Sixth Report of the Committee on Contact Catalysis - The Journal of

SIXTH REPORT OF T H E COMMITTEE ON CONTACT CATALYSIS‘

Downloaded by TEXAS A&M INTL UNIV on August 27, 2015 | http://pubs.acs.org Publication Date: January 1, 1927 | doi: 10.1021/j150293a001

BY ROBERT E. BURK*

As indicated by Chemical Abstracts, the volume of research in 1927 upon the theory of contact catalysis has not been enormous, many of the workers having been led into investigations of closely related fie theory of reaction velocity, adsorption, etc. At the newly-developed catalytic processes in industry, such as the synthesis of petroleum2 and of methan01,~progress in the study of enzyme^,^ etc., have rendered still more urgent the searching out of the mechanism of catalysis. The present report, in additiop to covering important new developments of the year in catalysis, summarizes the principal points of the five previous reports.6 An attempt is also made to summarize relevant work in other fields, which has come to our notice, with the hope that the contributing factors to catalytic phenomena will be thrown into somewhat sharper relief. These purposes are necessarily incompletely fulfilled. Two books of interest to catalytic investigators should be mentioned. A new edition of Rideal and Taylor’s “Catalysis in Theory and Practice” contains extensions and improvements and is well worth reading. The general plan of the book is good, and some of the chapters, for instance that on “promoters,” contain valuable information not readily found elsewhere. The theoretical treatment did not leave the writer with a feeling of very great satisfaction. The second book is a German translation of the second French edition of Sabatier’s “Catalysis in Organic Chemistry,” by B. Finkelstein. The translation proper is probably not so valuable to English-speaking readers as is the translation by Professor Reid; for the latter contains a chapter on theories of catalysis, many footnotes, and is arranged in better form. The new German volume, however, contains a summary by H. Hauber of catalytic literature between the years 1920 and 1926. *Sational Research Fellow. 1 Report of the Committee on Contact Catalysis of the Division of Chemistry and Chemical Technology of the Sational Research Council. Written by R. E. Burk with the assistance of the other membersof the committee: Messers. H.A. Adkins,E. F.Armstrong 0.W. Brown, C. G. Fink, J. C. R. Frazer, J. K. Pearce, E. E. Reid, H. S. Taylor, and D. Bancroft, Chairman. Killeffer: Ind. Eng. Chem., 19, 1077 (1927). 3 Jaeger: Abhandl. Kenntnis. Kohle., 7, 5 1 (1926); Chem. Abs., 21, 3531 (1927). Willstatter: “Problems and Methods of Enzyme Research,” Cornell University (1927). These five preceding reports are-First Report of the Committee on Contact Catalysis: Ind. Eng. Chem., 14, 326-31, 44 7, 545-8, 642-6 (1922); Nat. Res. Council Reprint No. 30, (1922). Second Report of t t e C o m m i t t e e on Contact Catalysis: J. Phvs. Chem., 27, 801-941 (1922); Nat. Res. Council Reprint No. j o (1923). Third Report ofihe Committee on Contact Catalysis: J. Phvs. Chem., 28, 897942 (1924); S a t . Res. Council Reprint KO. 59 (1924). Fourth Report d the Committee on Contact Catalvsis: J. Phys. Chem., 30, 145-71 (1926); Kat Res. Council Reprint No. 66, (1926). Fifth-Report of the Committee on Contact Catalysis: J. Phys. Chem., 31, 1121-49 (1927); S a t . Res. Council Reprint No. 78 (1927).

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1602

ROBERT E. BURK


Summary of Previous Reports

From the five previous reports of the Committee on Contact Catalysis, the following points remain unchallenged, so far as I know, and represent the principal features of general experience in contact catalysis. I . Contact catalysis does not take place unless one or more of the reactants is adsorbed on the surface of the catalyst.' 2. The complex thus formed on the surface may correspond to a definite orthodox chemical compound as in the case of the decomposition of hydrogen peroxide a t a mercury surface,2or it may not correspond to such a compound; for instance, no known oxide of carbon has the chemical properties of oxygen adsorbed on charcoaL3 3. The amount of gas adsorbed on the catalyst and the extent of catalysis do not show quantitative c~rrespondence.~ 4. Any substance strongly adsorbed on the catalyst will prevent reactants from reaching it, and will thus "poison" the ~ a t a l y s t provided ,~ the catalytic activity of the poison itself is small. The poison may be one of the reactants in case it is necessary for more than one reactant to be adsorbed upon the surface, as in the case of the poisoning of a platinum wire by carbon monoxide for the catalysis of the combination of this gas with oxygen.6 It may be one of the products of reaction, e.g., hydrogen poisons a platinum wire for the catalytic decomposition of ammonia by this metal.' Lastly, the poison may be a foreign substance of which there are many familiar examples such as the effect of sulphur compounds, stopcock grease, etc., upon the catalytic action of platinum. 5 . In some cases the amount of poison necessary to completely stop the activity of a catalyst is less than that necessary to cover it with a monomolecular layer.* 6. Some catalysts have been shown to have varying saturation capacities for various substances; and the ratio of the saturation capacities for adsorbed gases on a given substance has been faund to vary for different preparations, and for a given sample before and after heat treatment.1° 7. Poisons do not always have a proportionate effect upon adsorption and catalytic activity of the catalyst. 8. The results 5 , 6, 7 , and others, force the conclusion that some catalysts are not uniformly active." Bancroft: First Report of the Committee on Contact Catalysis. Bredig and von Antropoff: Z. Elektrochemie, 12, 581 (1906);von dntropoff: J. prakt. Chem., (2)77, 273 (1908). 3 Bancroft: First Report of the Committee on Contact Catalysis. 4 Taylor: Third Report of the Committee on Contact Catalysis. 6 Discussed in all of the reports of this committee. 'Langmuir: Trans. Faraday Soc., 17, 653 (1922). 7 Hinshelwood and Burk: J. Chem. Soc., 127, 1x05 (1925). 8 Armstrong and Hilditch: Tram. Faraday SOC., 17, 669 (1922). DPease: J. Am. Chem. Soc., 45, 2296 (1923);Hurst and Rideal: J. Chem. Soc., 125, 685,694 (1924). 'OMaxted: J. Chem. SOC.,127, 73 (1925);Pease: J. Am. Chem. SOC., 45, 2296 (1923); Taylor: Fourth Report of the Committee on Contact Catal sis. "Taylor: Fourth Report of the Committee on Contact %talysis.


SIXTH REPORT OF THE COMMITTEE ON CONTACT CATALYSIS

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9. The power of catalysts to activate molcules is often very specific.’ A particularly interesting case of the specificity of catalysts is the ability of different catalysts to accelerate the decomposition of a given molecule, e.g., formic acid2 in different ways. IO. It is sometimes found that the catalytic activity of a mixture of catalysts exceeds the additive. effect of the two. This is called promoter a~tion.~ 11. Continued reaction alters the activity of some catalysts, for instance, quartz in the decomposition of ammonia4; and causes visible roughening in other cases, for instance, platinum in the catalytic oxidation of sulphur dioxide.6 12. Certain substances called “negative catalysts” or inhibitors are found to decrease the rates of supposedly homogeneous reactions.e 13. Reactions involving a solid reactant and a solid product are often found to occur a t the interface between the two solid phases.’ 14. Processes such as heat treatment, which might well disturb special configurations of surface atoms are often found to decrease markedly the activity of catalysts.* Theories of catalysis must be consistent with these points.

Contact Catalysis and Reaction Velocity in General It may be thought that since catalytic phenomena have to do with altered reaction velocities, a solution of the problem must await a satisfactory solution of homogeneous reaction velocity. I n answer to this, it may be said, first, that the main paths are blazed in the theory of reaction velocity; secondly, as emphasized all along by BancroftJ8 that a great part of the problem of contact catalysis must consist in an analysis of the effect of adsorption upon the adsorbed molecule. Nevertheless, since the subject of contact catalysis deals with changes in reaction velocity, and since solid progress has been made in the study of the latter subject, it is thought desirable for the purposes of this report to summarize briefly the main points in the theory of homogeneous reaction velocity. Much of the work upon this subject has appeared since the first report of this committee was written. I This important point is discussed in all the previous reports of the Committee on Contact Catalysis. *Adkins and Krause: J. Am. Chem. SOC.,44, 385 (1922); Adkins and Xisson: 45, 809 (1923); 46, 130; Adkins and Lazier: 2291 (1924); Bischoff and Adkins: 47, 807 (1925); Hinshelwood, Hartley and Topley: Proc. Roy. SOC.,100.4, 575 (1922); Tingey and Hinshelwood: J. Chem. SOC.,121, 1668 (1922); Hinshelwood and Topley: 123, 1014; Hlnshelwood and Hartley: 1333 (1923). * Promoter action is ala0 discussed in all of the previous reports of the Committee on Contact Catalysis. ‘Hinshelwood and Burk: J. Chem. SOC.,127, 1105 (1925). See the photographs in Rideal and Taylor’s “Catalysis in Theory and Practice,” p. 177. See Reid: Fifth Report of the Committee on Contact Catalysis. Langmuir: J. Am. Chem. SOC.,38, 2263 (1916). Discussed in most of the previous reports of the Committee on Contact Catalysis. First Report of the Committee on Contact Catalysis, and elsewhere.

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ROBERT E. BURK

Attention was directed in the last report’ to three summarizing papers on this subject.2 I n addition to these, two recent books should be noticed. “Kinetics of Chemical Change in Gaseous Systems” by C. K. Hinshelwood (1926), although not exhaustive, is a stimulating review of the subject. It is written in such a clear fashion that the theory is within the reach of everyone. There are two chapters on heterogeneous reactions. The second book, “St’atistical Mechanics with Applications to Physics and Chemistry,” R. C. Tolman ( 1 9 2 7 ) is more mathematical than the book by Hinshelwood. If one is not frightened off by some of the topics in the earlier chapters such as “generalized space,” one finds much valuable material farther on in the book, especially with regard to photochemical reactions and certain physical questions which enter into work upon reaction velocity. Heterogeneous reactions are not discussed extensively. The modern theory of reaction velocity is founded on the very solid ground of the molecular kinetic theory of matter. To explain the reactions whose temperature coefficients exceed the rate a t which the total number of collisions increases with temperature, Arrhenius3 postulated that only certain “active” molecules enter into reaction, and that the number of these increases sufficiently rapidly with temperature to account for the observed reaction rates. It is now thought that these molecules are merely those with a sufficiently high energy content (an idea which accounts quantitatively for the experimental facts), rather than chemical isomers of the inactive molecules. Thus it would be difficult to imagine suitable isomers of hydrogen and iodine in accounting for the temperature coefficient of the combination of these two elements. On the assumption that active molecules are merely those of sufficiently high energy content, and that every collision between active molecules led to reaction, W. C. RlcC. Lewis4 was able to calculate theoretically the rates actually observed by Bodensteins for the homogeneous formation and decomposition of hydrogen iodide. Observed reaction rates a t different temperatures are found to conform to the empirical equation6 d In k / d T = A/RT2. “k” is the velocity constant, R the gas constant, T the absolute temperature, and A the heat of activation. A is constant when it is taken to be the excess energy’ the active molecules must have, over the average energy of all the molecules, in order for reaction to occur. From the experimental data and this equation.(a straight line is obtained with slope - h / R on plotting In k against I ~ T )an “observed” value of h is obtained. From the above assumptions the number of active 1 Reid: Fifth Report of the Committee on Contact Catalysis. 2Tolman: J. Am. Chem. Soc., 47, 1524 (192j); Lewis and Smith: 47, 1j08 (1925); Hinshelwood: Chem. Rev., 3, 2 2 7 (1926). 3 Arrhenius: Z. physik. Chem., 4, 226 (1889). 4 J. Chem. Soc., 113, 471 (1918). 5 2 . physik. Chem., 29, 2 9 j (1899). 6 Arrhenius: Loc. c i f . 7 Tolman: “Statistical Mechanics”, Zoc. cit.


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molecules of hydrogen is, according to the Maxwell distribution law, equal to the total number x e -AdRT (approx.), and the number of active molecules of iodine is equal to the total number X e-*A12/RT. The number of collisions between the active molecules, according to the formula given by kinetic theory, is, therefore, the total number of collisions between hydrogen and iodine molecules x e-hR2-A12/RT = the total number of collisions b%tween molecules of hydrogen and iodine x and by the assumptions of the theory should be equal t o the observed rate of disappearance of reactants. This gives a theoretical value of A. For hydrogen iodide formation, the observed value of A is 40,000 calories per z gram molecules, and the calculated value is 40,200 calories, a closer agreement than could be expected, granting the correctness of the interpretation. The individual values of AH^ and AI^, representing the extents t o which hydrogen and iodine must be separately activated, are not known. Indeed, it’ may be possible that it is immaterial how the energy of activation is distributed before collision, it merely being necessary that their joint transferable energy exceed the value A, the energy being pooled on collision. This possiblity was suggested by Hinshelwood,’ who has carried out calculations similar to the above for various reactions, which he and his students have investigated experimentally2. It is to be strongly emphasized that every known uncatalyzed bimolecular reaction is interpretable quantitatively on the above lines. Reaction, thermal decomposition 2 x 2 0 2 2

HI c120

A from collision calculation

Abs. T

A from temperature coefficient.

for equal rate 956

55,000

58,500

43,900

44,000

760

22,000

21,000

384

The last column gives the temperature a t which the various reactions attain the same speed. Hinshelwood3 in speaking of the parallelism between the figures of the last column and the values of A says: “This is the strongest and, a t the present time, the principal evidence for the reality of the energy of activation.” According to Hin~helwood,~ difficulties in accurately determining the heat of activation, inaccurate knowledge of the diameter of molecules, and the approximate nature of the distribution law used prevent one from deciding whether vibrational and electronic energies also contribute to the heat of activation in the above cases (only translational energy was considered in carying out the calculations). There is some evidence5 that the transference of these forms of energy depends upon the specific nature of the colliding molecules. Hinshelwood and Burk: Proc. Roy. SOC.,106A,284 (1924). and Hughes: J. Chem. Soc., 125, 1891 (1924). 3 Hinshelwood: “Kinetics of Chemical Change in Gaseous Systems,” p. 93. Chem. Rev., 3, 227 (1926). Stuart: 2. Physik, 32, 262 (1925).

* See, e.g., Hinshelwood


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ROBERT E. BURK

Of course, if upon introducing the reactants, subsidiary quickly-established equilibria are set up, so that the concentration of a given reactant will be a different fraction of that added for different temperatures, the heat of reaction for this subsidiary reaction will be included in the measured heat of activation.' Dey and Dhar2 propose the rule that the temperature coefficient is greater for smaller reaction orders. This cannot be an absolute rule for different reactions, for the temperature coefficients of the thermal decomposition of nitrogen pentoxide (monomolecular), and of nitrous oxide (bimolecular) contradict it. Termolecular reactions offer no particular additional difficulties. Their general characteristics are simply interpretable in terms of the molecular kinet,ic theory of matter. Vnimolecular reactions, on the other hand, continue to offer grounds for controversy. The question cannot be taken up in detail here. However, Tolman's calculation3 that collision cannot activate the molecules fast enough to account for the rates of known unimolecular reactions is thought by Lewis and Smith,4by Hinshelwood,j and by Fowler and Ridea16to offer no obstacle in case the molecules of reactant have sufficient degrees of freedom. However, Tolman, Yost, and Dickinson' think that there are still difficulties with activation by collision for this type of reaction, and favor an elaborated radiation theory as most satisfactory for unimolecular reactions. It has not been made clear why both collisions and radiation cannot play a part simultaneously in such reactions.s The confusion seems to arise over uncertainties regarding (I) the amount of radiation which the reactants can absorb, ( 2 ) the laws governing the transfer of internal energy a t the moment of collision. In recent experiments of Lewis and Mayer,Ya stream of pinene vapor was passed through a furnace under such conditions that there were no collisions, yet sufficient radiation for complete racemization. They found no racemization, and consider this to rule out radiation as the activating mechanism for this monomolecular reaction. I n the case of unimolecular reactions, the possibility was emphasized by LindemannlOthat molecules active on the energy score may have to await a suitable internal phase relationship before reaction is possible, and thus may have percentage reaction rates independent of pressure over certain ranges, even though the activation is brought about by collision. Giordani: Rend. Accad. Sci. Napoli, 32, 70 (1926); Chem. Abs., 21, 3797 (1927). Z.Elektrochemie, 32, 586 (1926). 3 J. Am. Chem. SOC.,47, 1524 (1925). 4 J. Am. Chem. SOC.,47, 1j08 (1925). 5 Proc. Roy. SOC.,1139, 230 (1926). 6Proc. Roy. SOC.,113.4, 570 (1927). 7Proc. S a t . &ad. Sei., 13, 188 (1927). 8 For further recent work on monomolecular reactions, see Ramsberger: J. Am. Chem. Soc., 49, 912, 1495 (1927); Proc. Kat. Acad. Sci., 13, 849 (1927); Christiansen: Proc. Cambridge Phil. Soc., 23, 438 (1926); Hibben: Proc. Kat. Acad. Sci., 13, 626 (1927); Hinshelwood and Askey: Proc. Roy. SOC.,115A,2 1 j (1927); 116A, 163 (1927). QProc.Nat. Acad. Sci., 13, 623 (1927); J. Am. Chem. Soc., 49, 3046 (1927). 'OTrans. Faraday SOC., 17, 598 (1922). 1

2


SIXTH REPORT OF THE COMMITTEE

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CONTACT CATALYSIS

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Thornson‘ now suggests a new type of internal phase relationship. He develops the view mathematically that if the intramolecular forces are intermittent, a Maxwellian distribution of energy amongst the molecules may be produced without the aid of collisions, or of the absorption and emission of radiation. He says, “The molecules with high energy will, however, not have acquired this energy by collision with other molecules but by drawing on the energyof the fields of force of the molecules.” decomposition of nitrogen pentoxide depends (on this basis) on the number which attain the critical energy in unit time, rather than upon the number which have this energy when things are steady.” From the point of view of catalysis, the important points .rvhich emerge from the work on homogeneous reaction velocity are that the velocities of reactions are governed ( I ) by the rates a t which the energyof activation can be supplied to the reactants, Le., by the number of activated molecules per unit time (2) by a factor corresponding to the probability of reaction after activation. ( 2 ) may result from the necessity of a suitable internal phase relationship. The second factor, in the case of bi- and multimolecular reactions may depend upon a possible necessary suitable orientation as the active reactants come into the sphere of collision. This effect will not change with the temperature because an increase in the number of revolutions of the molecule per unit time will not, ordinarily, change the chance that a given orientation will occur. According to Hinshelwood? (2) in no case can depend upon temperature. At least the variation with temperature cannot be very great-othernise reactions would not follow the hrrhenius equation. In his work upon homogeneous reactions, Hinshelwood3 also places preponderant weight upon the first factor for reactions other than bimolecular. This view is based largely upon the approximate constancy of A/RT for reactions of a given type at temperatures of equal reaction velocity. While this is evident enough for the simple bimolecular reactions, in view of the above calculation and others like it, it is a little surprising to find this relation true for unimolecular reactions since in their case so much weight has been placed on only a few of the activated molecules r e a ~ t i n g . ~ I n any event, the factor ( 2 ) , above, would seem to have more importance in catalysis than in uncatalyzed reactions for say its whole range of influence could lead but to a factor of ten in the reaction velocity. If a catalyst could change its value from one to ten, this would be of great importance for simultaneous reactions, and in other cases, but it might not be detected for the homogeneous reactions. According to Peacock,j substitutions in certain organic molecules affect their reaction velocities much more than the corresponding heats of activation Phil. Mag., 3, 241 (1927) Chemical Reviews, 3, 227 (1926). a Chemical Reviews, 3, 2 2 7 (1926). 4 Lindemann: LOC. at. J. Phys. Chem., 31, 535 (1927).


1608

ROBERT E . BURK

for these reactions. In some cases, the reaction with the higher heat of activation had the lower rate. This would indicate considerable importance for factors other than the heat of activation. However, the heats of activation which Peacock quotes are close together and accurate measurement of them is difficult. Although one cannot be sure that these two factors contain the whole story of homogeneous reaction velocity, yet they seem to account, at least in outliilZ, for the observed phenomena, and give very definite factors for the catalytic investigator to keep in mind in trying to ascertain how catalysts effect the acceleration of reactions. Thus, if a catalytic agent can facilitate the energy transfer to or away from the seat of reaction, if it can reduce the amount of energy necessary for reaction, or if it can increase the probability of reaction after activation, then the phenomenon of catalysis will ensue. As far as the writer is aware, little is known about the last factor. The possibility that the surface can cause the reactants to come into suitable juxtaposition1 more frequently comes under this head. Oriented adsorption may aid or hinder the transfer of energy to a given part of the adsorbed molecule, or may render a given part of the adsorbed molecule more, or less susceptible to collision from the gas phase. Facilitation of Energy Transfer Asszstance in the Dissipation of the Heat of Reactzon According to Polanyi and Herzfeld* and to Born and Franck3 certain B = C cannot take place unless there exothermic reactions of the type A is a third molecule present to carry off the heat of reaction, the system being unable to radiate it. This is one way in which a catalyst might assist certain reactions The reaction of dried ammonia with dried hydrochloric acid is visibly heterogene~us.~ The effect of the glass walls here might be due to their assistance in the dissipation of the heat of reaction. This mechanism might also explain the effect of metals in catalyzing the recombination of hydrogen atoms,5 which has been applied by Langmuir6 in an ingenious fashion to the welding of metals. (I)

+

I

Facilitation of Energy Transfer to the Reacting Molecules The Lewis-Perrin theory of catalysis? would come under this head. Although this theory offers no hope of explaining catalysis in general, no grounds exist for ruling it out completely. Catalysis by photosensitized mercury atoms, which is being actively investigated at the present time,* is Langmuir: Trans. Faraday SOC.,17,618 (1922). Z. I’hysik, 8, 132 (1922). “nn. Physik, (4)76,225 (1925). Burk: Dissertation, Oxford (1926). ‘Wood: Proc. Roy. SOC.,102A, I (1922). Bind. Eng. Chem., 19, 6 (1927). 1 See the First Report of the Committ,ee on Contact Catalysis for references. Also Penin: Trans. Faraday SOC.,17, 566 (1922). 8 Taylor and Bates (summarizing paper): Proc. Nat. Acad. Sci., 12, 714 (1926). 2


SIXTH REPORT O F THE COMMITTEE O S COSTACT CATALYSIS

1609

closely related to the mechanism proposed by Lewis and Perrin for the action of surfaces. The homogeneous thermal decomposition of nitrous oxide’ requires a heat of activation of about 58,000 calories per two gram molecules, and is bimolecular. The heat of activation of the heterogeneous decomposition on gold2is about the same per molecule, but the reaction is unimolecular, collision between active molecules being unnecessary. Here the increase in rate results from the kinetic simplification. Hinshelwood3 presents the following picture which is a possible factor in catalysis and comes under this head. “We may suppose a molecule to be composed of two parts, A and B, the separation of which constitutes the chemical decomposition of the molecule. Let B receive an impact from another molecule, which imparts to it momentum directed away from A. The small inertia of A enables it to follow B without the development of much strain between them. If, however, A were firmly enough held to a surface, its inertia might be so great that the accelerating force instead of drawing A after B would cause the disruption of the bond between them.’’ Lowering of the Heat of Activation The great importance of the energy of activation in homogeneous reaction velocity has already been emphasized. It is reasonable to expect the energy of activation to play a correspondingly important rBle in heterogeneous reactions, since the two types cannot be fundamentally different. I n some heterogeneous reactions, e.g., the catalytic thermal decomposition of ammonia upon a tungsten surface4 the other factors seem to drop out. At IOO mms. pressure, the reaction is experimentally of zero order, and therefore cannot depend upon collisions from the gas phase. Other experiments show that the reaction is really monomolecular and, therefore, cannot depend upon a suitable orientation of colliding molecules of ammonia upon the surface. The photochemical decomposition of ammonia requires ultraviolet light, and it is most unlikely that the tungsten can give rise to the necessary increase in density of this radiation under the conditions of thermal decomposition. For such a simple molecule, judging from the behavior of other simple molecules in the bimolecular reactions already discussed, the probability of reaction after activation must be high. The reaction is unretarded by products, and since the reaction is zero order (i.e., the active surface is practically completely covered with ammonia), the observed heat of activation, 38,700 calories per gram molecule, is the true heat of activation. The heat of activation for the homogeneous bi-molecular change must be greater than 80,000 calories per gram molecule, since this reaction is not perceptible at 1000degrees centigrade. The homogeneous unimolecular reHinshelwood and Burk: Proc. Roy. Soc., 106A,284 (1924). Hinshelwood and Prichard: Proc. Roy. Soc., IOSA,211 (1925). “Kinetics of Chemical Change in Gaseous Systems,” p. 197. Hinshelwood and Burk: J. Chem. Soc., 127, 1 1 0 5 (1925).


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ROBERT E. BURK

action would probably have a higher heat of activation still, and it is this reaction with which the unimolecular heterogeneous reaction should be compared. There seems to be little doubt that in this reaction the catalyst has effected a large reduction in the heat of activation. There is also little doubt that catalysts often do this. Euler and Olanderl find that the temperature coefficient for the enzymatic inversion of cane sugar i s smaller than without the enzyme. Many other examples could be quoted. How can the catalyst lower the heat of activation of the reactants by adsorbing them? That it must “distort,” “dislocate,” “strain,” or “profoundly modify” the adsorbed molecule in some way seems to be an opinion widely held.* But these terms are too indefinite to serve as a guide in prescribing catalysts for given purposes. Two rather definite ways have been suggested in which bonds in the adsorbed molecules can be loosened, or opened up, and the energy of activation thus lowered for the breaking of that bond. I. A t o m i c distortion theory. The older picture3 is, quoting Langmuir : “In view of the structure of atoms from positive and negative particles, it is clear that the atoms should have the properties of a dielectric. Thus, if we had a chain of atoms held together by duplets-as, for example, in the hydrocarbon chain of an organic compound -and we bring a positively charged body near one end of the chain, the electrons will be attracted and the nuclei repelled, so that a certain displacement of these particles with respect to one another will result. This effect is then transmitted with gradually decreasing intensity from atom to atom throughout the length of the chain. The chemical evidence indicates clearly that effects of this kind are transmitted relatively great distances . . . I n cases where atoms are not joined together by duplets, we should never expect the transmission of electric force to extend through more than about one atom.” On page 616 of the same paper we find, “when gas molecules condense on a solid surface in such a way that they are held on the surface by primary valence forces involving a rearrangement of their electrons, their chemical properties become completely modified. It is not surprising, therefore, that in some cases such adsorbed films should be extremely reactive, while in other cases they may be very inert to outside influences.” Presumably, some such idea is behind the speculations of Flur~cheirn,~ who postulates, if the writer understands him, that, if different atoms are attached to a foundation atom, the greater the strength of attachment of one atom the smaller is the remaining available total valence force of the foundation atom, for other atoms. lZ.anorg. allgem. Chem. 156, 143 (1926). ZBoeseken: Rec. Trav. chim., 29, 85 (1910);30, 381 (1911);,39,623 (1920);46, 458 7 (1926);Raachig: Z.angew. Chem. 19, 1748,2084,(1906);Bodenstein: Ann., 440,~ 7 (1924);

Zelinaki: Ber., 58, 2755 (1925);kinshelwood: Kinetics of Chemical Change in Gaseous Systems,” p. 188. ’Graham: Proc. Roy. Soc., 16, 422 (1868);Langmuir: Trans. Faraday SOC.,17, 610 (1922). Chem. and Ind., 3, 246 (1925)(summarizing paper).



1611

Possibly some such idea as that sketched above is back of the suggestion of Taylor’ that the action of oxides in accelerating both the dehydration and dehydrogenation of formic acid, though metals accelerate “only” the dehydrogenation,* is due to the presence of both postive and negative ions in the oxide. I n this connection, it is not obvious why the electrons in the metal cannot induce dehydration as well as the negative ions of the oxide, if the charge is the important thing. On the basis of Taylor’s idea, a stream of electrons, however slow, should be able to dehydrate formic acid, and should have general catalytic action. No experiments were avilable to the writer on this point. I n connection with this atomic distortion theory of surface action, i t is to be remarked that if the electrical forces of the catalyst actually do seriously distort the electronic orbits in the atom or atoms adjacent to the catalyst, in some cases it seems possible that this would have no effect upon the succeeding bond; in other cases i t might strengthen the succeeding bond, and in still other cases weaken it. These possibilities correspond to the three cases cited by F r a n ~ k in , ~ which bonds can be weakened, strengthened, or unaffected by electronic excitation in the constituent atoms. From the work of Birge and S p ~ n e rand , ~ Birge,s it seems that, according to the quantum interpretation of band spectra, the energy necessary for the dissociation of an excited molecule is sometimes about the same as that for the unexcited molecules. Further evidence that possible electronic disturbance does not greatly affect certain bonds is the persistence of the positions of certain infra-red spectra, and the characteristics of certain ultra-violet spectra which are traceable to certain bonds, for instance, C - H, C - C, 0 - H, and N - H, in various compoundss in which different substituents are attached to the atoms constituting the bond in question. On the other hand, Ephraim and Block’ found that in the reflection spectra of the series of ammonates of PrC13,there were shifts in the positions “Colloid Symposium Monographs,” 4, 19 (1926). does not exactly represent the accurate facts. There is some dehydration on silver, for example, a t 185”. (Hinshelwood and Tingey: J. Chem. SOC.,121, 1668 (1922)). At 355”, tinned iron gives mostly carbon dioxide and hydrogen; but a t 255”, the same catalyst gives also carbon monoxide. (Tro sch: Abhandl. Kenntnis Kohle, 7, I (I925).); Chem. Abs., 21, 3530 (1927). The remarkabye instance is given by Sabatier ("Catalysis ln Organic Chemistry”. p. 76) where fmely divided chromium oxide decomposes ethyl alcohol giving a gas containing 9 1 % ethylene (dehydrat,ion). The crystallized oxide is not so active, and at higher temperatures c a w s dehydrogenation almost exclusively. As Professor Taylor haa reminded me, films of oxide may bring these results on silver and tin into line; but films may also jeopardize the conclusions on the pure metals. ‘Discussion of the Faraday Society on Photochemistry, 536 (1925). Phys. Rev., (2) 28, 259 (1926). “Molecular Spectra in Gaaes,” Nat. Res. Council Bull., No. 57 (1926). e Coblenta: Investigations of Infra-Red Spectra, Camegie Inst. Pub. (1905-1908); Le Compte: Compt. rend., 178, 1530, 1698, 2073 (1924); Marton: Z. physik. Chem., 117 97 (1925); Ellis: J. Optical SOC.America, 8, I (1924); Phys. Rev. (2) 23, 48 (1924); 27, 298 (1926); 28, 2 5 (1926); Proc. Nat. Acad. Sci., 13,202 (1927); Bell: J. Am. Chem. SOC.,97, 2192. 3039 (1925); 48, 813 (1928); Bonino: Chem. Abs., 20, 709, 710,2950 (1926); Henri: “Structures des Mol6cdes” (1925); MortonandRiding: Proc. Roy. Soc., ll3A, 717 (1926). Ber., 59, 2692 (1926).

* This


1612

ROBERT E. BURK

of the spectra toward the ultra-violet, which they attribute to compression of the electron shells. However, according to Biltz,’ the molecular volume of ammonia in ammonates is very close to the value in the gas phase. I n this same connection, according to Schaefer and Schubert? the inner vibration frequencies of ions (e.g., nitrate, sulphate, and carbonate ions), are independent of the nature of the metal ion and therefore of large differences in electric field strength, (arising from differences in screening constants of the postive ions in the case of ions of the same apparent charge). According to Taylor and Rideal,3 the characteristic frequencies of vibration of sulphur are almost the same in all the phases. Bowen‘ has pointed out the constancy of heats of linkage in various compounds. The writers has recently carried out the thermal decomposition of ammonia upon fine electrically heated wires, in the presence of electric field strengths as high as 140,ooo volts per cm. This field is as strong as those used in investigations of the Stark effect and therefore produces some distortion of the electronic orbits. No effect was observed upon the rate of decomposition of the ammonia. The experiments, however, are not considered to be conclusive against the atomic distortion mechanism for lowering the energy of activation. The rate of decomposition of nitrogen pentoxide is about the same in a solution of chloroform as in the gas phase,* yet chloroform has a very considerable electric m ~ m e n t . This ~ is strong evidence that the electric fields of chloroform do not affect the bonds of nitrogen pentoxide. Similarly, Smith8 has found that the rate of the unimolecular racemization of pinene in solution in petroleum, acetophenone, and alpha methyl naphthalene is nearly the same as that in the pure liquid, and in the gas phases. On the theoretical side, Kembleg has pointed out that one of the characteristics of the quantum behavior of atoms, is the way in which electrons in orbits ignore forces which, on the basis of classical mechanics, would cause a profound change in their path. He points out that the electronic orbits must be essentially preserved upon collision, etc., if this is not too violent, and also points out the similarity of situation of atoms in collision and in compounds. Thus it may turn out that the quantum laws will rule out the atomic distortion mechanism for the lowering of the energy of activation by catalysts, at least in some cases. It is to be remembered in this connection that catalysts have to deal with molecules, in which loosely attached, or transferable electrons have had the opportunity of taking up positions of relatively greater stability than in the case of the free atoms. Z. anorg. allgem. Chem., 159,96 (1926). zAnn. Physik, (4) 50, 283 (1916); 55, 577 (1918). Proc. Roy. SOC.,115A, 598 (1927). * Discusaion of the Faraday Society on Photochemistry, 544 (1925). Proc. Nat. Acad. Sci., 13, 719 (1927). eLueck: J. Am. Chem. SOC.,44, 757 (1922). 7 Henri: “Structure des MolBcules,” Zoc. at. J. Am. Chem. SOC.,49, 43 (1927!. Kemble: In “Molecular Spectra in Gases,’’ Nat. Res. Council Bull., No. 57 (1926).

SIXTH REPORT O F THE COMMITTEE ON CONTACT CATALYSIS

1613


However, the work of Born’ and others seems to give considerable importance to the deformation of ions and atoms, in accounting for chemical bonds. One would expect the atomic distortion, or electronic orbit distortion effect, to influence allthe chemical bonds within the range of its action.2 This is not very well in line with the extreme specificity of some catalysts. One of the things catalytic chemists must explain some day is the ability of enzymes to accelerate exclusively one reaction with a large delicate organic molecule capable of many reactions. Atomic distortion, carried through the chain, does not seem to the writer to be a very promising approach to this problem. With regard to the chemical evidence for such effects, it is not obvious that the results of organic chemistry, such as the effect upon the strength of an acid of substituents in the chain, should point unequivocally to atomic distortions. Such substitutions could affect the degree of association of the organic molecules, could affect the spatial position of the atoms in the molecule with respect to each other,3 etc., without affecting the individual bonds. Finally, with regard t o the atomic distortion, or electronic orbit distortion theory for the lowering of the energy of activation by catalysts, we may conclude that it has not been shown definitely to be impossible in any case, although in some instances there is considerab!e evidence that i t is inoperative. Although the theory is old enough, it has never enabled one to prescribe a catalyst for a given purpose-which is not what one would expect of a correct theory of phenomena in which such wide interest has been shown as in catalysis. Nevertheless, the work of Born and others compels one to give the theory serious consideration. One knows that in some reactions there is a complete transference of one or more electrons. Presumably, however, such processes take place in accordance with quantum laws. The all-important point then is: If an external force is not sufficient to effect a complete quantum change, does it, nevertheless, distort the arrangement of the electronic orbits seriously or not? If it does distort them seriously, what part does this distortion play in catalysis? Molecular Distortion Theory There is another way in which adsorptive forces can lower the energy of activation, namely by actually partially separating the atoms forming the bond which is to be broken. A consideration of the extremely short range of chemical forces led Langmuir4 to the extremely fruitful, and now generally accepted, theory of monomolecular adsorbed gas films on solids. h consideration of the same thing, namely the short range of chemical forces, led See the discussion of “bonds” 1 Z. Elektrochemie, 30, 382 (1924)(summarizing paper). later in this report. Fajans and Joos: Z. Physik, 23, I (1924). cording to Langmuir (Zoc. cit.) the effect can he carried relatively long distances thr:$ a chain ’ ’ of atoms. 3By an effect which could be pictured as being similar to the opening or closing of a hinge by forces (due to the substituent in the case of the organic molecules). 4 ,J, Am. Chem. Soc., 38, 2252 (1916); Trans. Faraday Sac ., 17, 610 (1922).


1614

ROBERT E. BURR

the writer' to the theory that the forces of adsorption (apart from the atomic distortion effect discussed above) can separate the atoms forming a bond only if the molecule is multiply adsorbed on the surface, Le., adsorbed a t more than one point by more than one atom of the catalyst. The catalytic effect of the surface clearly would then depend upon the spacing of the adsorbing points in the surface, which spacing would have an optimum value for a given reactant. The catalytic effectivenessof the surface would also depend upon the specific attraction of the adsorbing atoms of the surface for the atoms of the adsorbed molecule by which it is, in the main, attached to the surface. That a gas could be attached to a surface by more than one point seemed the best explanation for observations of Hinshelwood and Burk2 that added hydrogen sometimes failed to poison a quartz surface for the thermal decomposition of ammonia. A similar situation exists in the case of the poisoning action of carbon monoxide for its oxidation a t a quartz ~ u r f a c e . ~

A rather striking result of Tropsch and von Philippovich4 should be mentioned in this connection. They found that the catalytic activity of nickel in decomposing carbon monoxide was almost completely destroyed by alloying the nickel with jifty percent tin. The possibility a t once suggests itself that isolated nickel atoms, and tin-nickel pairs cannot effect the decomposition, and that suitable nickel-nickel spacings in the alloyed catalyst are infrequent. The multiple adsorption theory is mechanically possible, and catalytic action must follow from it necessarily. If conditions (spacing, etc.) are right, it must exist, independent of the existence of the atomic distortion mechanism. It is therefore a mechanism which should be kept in mind when considering any case of catalysis. I n the case of symmetrical molecules such as H?,N?, 02, etc., why should one atom be attached to the surface in preference to the other? Similarly, why should not straight chain hydrocarbons, dibasic acids, and a great many other substances lie flat upon the surface? According to good evidence of Palmer and Constableb for the case of alcohols adsorbed on copper. the molecules are attached to the copper by the alcohol group. The entire CH20H group cannot possibly lie over a single copper atom, and if i t does not, then the localzzed forces of the other copper atoms cannot avoid stretching the alcohol group in the way sketched above. Sone of the points mentioned a t the beginning of this report is in contradiction to this theory. Promoter action, and the specificity of catalysis, even to the extreme found in enzyme action, become particularly clear on this basis. J. Phys. Chem., 30, 1134(1926).

J. Chem. SOC., 127, IIoj (1925). 3Bodenstein and Ohlmer: Z. physik. Chem., 53, 166 (19oj);Benton and Williams: J . Phys. Chem., 30, 1487 (1926). dbhandl. Ilenntnis Iiohle, 7, 44 (1925);Chem. bbs., 21, 3530 (1927). 6Proc. Roy. SOC., 107.A, 2jj (192j).


1615

The Nature of Chemical Bonds


If we are to come to a sound decision regarding the effect of the catalyst upon chemical bonds, i t seems necessary to discuss the question of the nature of such bonds. A short summary will therefore be given of recent efforts toward the solution of that problem.’ Starting with the universally accepted premise that matter is composed of positive and negative charges of electricity, that these are in motion, and that they have mass, it would seem that gravitational, magnetic, and electrostatic forces should be taken into consideration in an examination of the question of the nature of chemical bonds. There is no need to assume a special “chemical” force until the three ordinary types of physical force can be shown to be incompetent of explaining chemical combination. This problem presents, perhaps the most urgent immediate need for co-operation between chemists and physicists. According to Herzfeld* one can immediately neglect gravitational and magnetic forces. He says “concerning the nature of the fields (around atoms) only electrostatic actions need be considered. A11 other forces (electromagnetic or gravitational) are too weak: there are no grounds for assuming an unknown force.” Kemble says: “Calculations by Mensing seem to indicate that magnetic forces between magnetic atoms are not strong enough to play an important part in molecular formation.” Such statements come as something of a shock to chemists who have been under the influence of the Lewis-Langmuir theory of valence. Lewis says: “. . the pairing of electrons, which I have regarded as the most fundamental phenomenon in all chemistry is some sort of conjugation of two magnetons of such a character as t o eliminate their magnetic moment.” He also says, p. 147, “In previous chapters i t has occasionally been hinted that in place of the electric it is the magnetic properties of the atom and the molecule which determine their essential structure. I n the present chapter, we shall give free rein to this idea.’’ 1 Since completing the manuscript of this report, a copy of S . V. Sidgwick’s new book, “The Electronic Theory of Yalency” (I927), has come into the writer’s hands. He strikes an admirable note in the preface, where he says, “In developing the theory of valency there are two courses open to the chemist. He may use symbols with no definite connotation . . . . or he may adopt the concepts of atomic physics-electrons, nuclei, and orbitsand try to exulain the chemical facts in terms of these. But if he takes the latter course, as i s done in this book, he must accept the physical conrlusions in full, and must not assign to these entities properties which the physicist has found them not to possess. He must not, use the terminology of physics unless he is prepared to recognize its laws.” However, a rather superficial examination of the book (in the short time available) does not show that Hidgwick has emphasized the points about valence which, from the point of viex of catalysis, must be brought out. One must know something of the actual nature of the binding forces, the law according to which they vary with distance, and under what’conditions, and to what extent they can be affected by other forces. * Geiger and Scheel’s “Handbuch der Physik”, 2 2 , 392 (1926). In “LIolecular Spectra in Gases,” S a t . Res. Council Bull. KO.57 (1926). “Valence and the Structure of Atoms and llolecules,” p. 148 (1923).


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ROBERT E. BURK

Williams,’ Shaffer m d Taylor,* and T a y l ~ rfind , ~ molecules with an odd number of electrons to be paramagnetic and those with an even number diamagnetic, which they suggest to be evidence for magnetic valence forces. Shaffer and Taylor find that complex ion formation with paramagnetic ions reduces the susceptibility of these ions and say: “We must conclude from this that chemical bonds are partly or wholly magnetic in character, or that they are formed by the same electrons which give rise to the magnetic properties.,’ Although the latter conclusion is preferable to the former, neither seems inevitable. It seems that gravitational forces are more insignificant than magnetic forces, as compared to the electrostatic forces available around atoms. If then, gravitational and magnetic forces are really unimportant, it throws the whole burden of chemical valence upon electrostatic forces. The use of purely electrostatic forces in holding the simple extremely polar molecules (like sodium chloride) together needs no defense from the point of view of either the chemist or the physicist; but such bonds are the least interesting from the point of view of catalysis. While non-polar bonds have not been worked out to one’s complete satisfaction, their general features seem t o be in line with the possibilities of the purely electrostatic theory. The striking characteristics of such a bond are : (I) constituent atoms are themselves electrically neutral, (2) the nonpolar forces are extremely specific, (3) they fall off not in accordance with the inverse square of the distance, but with a power usually much higher than the second? If the charges in an atom are not symmetrically distributed (the negative charges around the nucleus,) the electrically neutral atom will have a resultant electrical moment; it will be a dipole. If the center of gravity of the positive charges in the atom coincides with the center of gravity of the negative charges, there can still exist a pole of highei order (in case all points in a spherical shell around the nucleus are not a t the same potential, which seems to be the case in atoms). Now dipoles and the poles of higher order attract one another. The following table5 gives the inverse power of the distance according to which the potential energy between two poles of the order indicated varies with distance. Dipole Quadrupole

Pole Dipole Quadrupole Cube pole

2

3

Tetrahedral Pole

Cube Pole

3 4

4

5

5

6

5

6

7 9

Phys. Rev., (2) 28, 167 (1926). J. Am. Chem. Soc., 48, 643 (1926). J. Am. Chem. Soc.. 48, 654 (1926). 4Langmuir: J. Am. Chem. Soc., 38, 16j2 (19x7). Herzfeld: Geiger and Scheel’s “Handbuch der Physik,” 22, 440 (1926).



1617

The values given in the table cover the range estimated from experiments with actual atoms and listed above as one of the characteristics of non-polar valence forces. The pole orders are all possible ones on the basis of the Bohr theory of the atom. If the electric charges are in motion, the pole orders are said to be increased’ somewhat. The possibility exists that there must be a synchronization of electronic periods in the two relevant atoms for firm binding.’ These considerations would lead one to expect forces with physical characteristics similar to the ones coming into play in non-polar unions around all atoms whose structure is according to the Bohr theory. For atoms of greatest symmetry (the rare gases) the pole order would be highest, the cohesive force would fall off with a very high power of the distance, and the f(r)dr would be energy necessary to overcome the cohesive forces small. f(r) is the force; r is the nuclear separation; ro is the equilibrium value of r without external forces. So far so good. But things are not so simple as this. The carbon atom has high symmetry, but the energy necessary to separate two carbon atoms is great. On the other hand, the energy necessary to separate two helium atoms is small, whereas the pole order cannot be very great (if the nucleus is considered as a point charge). A consideration of the preceding table gives a fair general picture of the way non-polar forces can be specific. If the synchronization factor, mentioned above, is real, it would of course involve specificity of the highest degree. Debye3 has sought to account for van der Waals forces through mutual induced polarization of the atoms. (The fields due to the permanent poles would cancel out a t large distances for rotating molecules). Born4 has calculated the magnitude of induced polarization in the case of the ions of sodium chloride and other molecules, and considers the effect of great importance in chemistry. He says, “The principal result of the investigations is that, energetically a t least, there is no place for directed valences, in the ordinary sense of the word, in the cases considered.” It would seem that induced polarization would be of greatest importance when an ion is in the immediate neighborhood of an easily deformed atom or ion; but that the very strong bonds between carbon atoms, for instance, should be due to any marked extent to the mutual deformation of the respective atoms seems rather unlikely. Palmers has continued his work on the critical voltages necessary to make a current flow through a coherer, when different gases are present. He appears to consider his results to be a verification of the dipole theory of adsorption. Whatever can be said for the theory, the experimental “test” of it led to

JE

Herzfeld: LOC.cit.

* Kemble: LOC.cit.

Physik. Z.,21, 178 (1920).

* Z. Elektrochemie, 30, 385 (1924)(summarizing paper). 5Proc. Roy. SOC.,llOA, 133 (1926).


1618

ROBERT E. BURK

results in contradiction to those obtained in other ways; for instance, the heat of adsorption of hydrogen on tungsten was found to be greater than that of hydrogen on platinum, and both much greater than the other results would lead one to expect. The possibility of an electron executing an orbit belonging to two nuclei is the next to be considered. This corresponds to the concept of the sharing of electrons, which has been found to be convenient in working out “octets”, etc. It is desirable to make this idea as definite as possible, for in so far as the nucleus of one atom exerts some attractive force upon all of the electrons of an adjacent atom, all of the electrons are shared.’ The idea of shared electrons, or binuclear orbits, does not seem to be in such great favor with those who have tried to work out the dynamic models in the simple cases.* Kemble makes the following statement. “Any explanation of the attractive force between the atoms in a non-polar molecule consistent with the possibility of an adiabatic dissociation into normal atoms must depend upon the dectric field due to the polarization of the atoms, whether that polarization be of the simple dipole type or of the quadrupole type.” (Such dissociation is apparently indicated by certain work on band spectra). This does not leave much room for shared electrons, if the statement is right. On the other hand, according to M ~ l l i k e n the , ~ multiplicity of the electronic levels revealed in the band spectra of certain molecules is of the same type as for the line spectra schemes for atoms with the same number of external electrons; for instance, CX is compared with sodium. This would indicate a rather aggravated case of the sharing of electrons. It would seem difficult to extend the idea, e.g., to cane sugar. It has been proposed that in certaininstances, groups containinganumber of electrons different from eight are stable. Thus, Stoner4proposes ten in some instances. Grimm and Sommerfelds suggest four in some cases; Niven6 allows six, etc.’ Because some molecules have been found to possess a considerable electric moment, there has been a tendency to consider the links as ionic, in disregard of chemical evidence. Thus, Hund* has considered the ammonia molecule to contain the ion N - - -and three H+ ions. Van Arkel and de Boerg consider methane to contain the ion C++++ and four H- ions. If these structures represent the facts, the chemist would like an explanation of their chemical behavior. 1

Sidgwick (“The Electronic Theory of Valency”) discusses such orbits.

* See Kemble: LOC.cit. 3

Phys. Rev.,

(2)

26,561 (19z5),and elsewhere.

‘ Phil. Mag., 3, 336 (1927).

2. Physik, 36, 36 (1926). Phil. Mag., 3, 1314(1927). Sidgwick, “Electronic Theory of Valency,” p. 62. * 2. Physik, 32, I (1925). Z. Physik, 41,27 (1927). 6 6


SIXTH REPORT OF THE COMMITTEE ON COXTACT CATALYSIS

1619

It is of great importance whether a link is ionic or not, and if only partially so, to what extent it is ionic. Thus Grimm and Sommerfeld’ present evidence that the ions in boron nitride are B-and N+and not Bfff and Y--as one might expect if the electrons transferred in such a way as to complete their octets. If catalysts act, in part, through disturbance of electrons by electrical forces, boron nitride might be expected to be a sensational catalyst, on the basis of B+++K---, but not as B-S+. Franck, Kuhn, and Rollefson,* and Franck and Kuhn3 seem to have developed a reliable method for distinguishing between ionic and non-ionic links, if their interpretation is correct. According to them, a purely polar diatomic molecule can be dissociated photochemically in one elementary act into either normal atoms, or a normal and an excited atom, whereas nonpolar or semi-polar molecules, on dissociation in a similar way, cannot give two normal atoms. Apparently, rules governing the formation of ions cannot be laid down from the ionization potential data. Noyes and Beckman4 say, “In general, energies of removal of the successive electrons seem to play a secondary part in determining the valencies that exist.” However, Rolla and Picardij trace certain relations, not altogether new, between chemical properties and ionization potentials. Ions like hl+++ would seem very unlikely on the grounds of the energy necessary to remove the electrons, yet the chemical evidence for their existence is hard to get around. One would like to know why Coulomb’s law should be brilliantly confirmed in Bohr’s theory of the hydrogen atom, and disregarded in favor of certain alleged stable arrangements of electrons when i t conie to the matter of the formation of ions6 Chemists apparently now recognize polar, non-polar, hemi-polar’ and coordinated links.* Sidgwick shows that the electron pair interpretation of the coordinate link results in an electric moment and is therefore inclined to identify the coordinate and hemi-polar links. There are, however, cases of molecules without coordinate or polar links which have marked electric moments (ammonia), in which case perhaps the idea of a hemi-polar link may apply. Investigations of band spectra are yielding some of the most direct and definite information regarding the nature of bonds, and may offer the best cit. Physik, 43, 155 (1927). Z. Physik, 43, 164 (1927). Proc. Sat. Acad. Sci., December (1927). Chimie et Industrie, 16, 531 (1926). 6 S o attempt will be made in this report to discuss valence or any of the problems connected with catalysis in terms of the new quantum mechanics. I t is too difficult to get p h sical concepts from the new mechanics, and there is no indication that the problem of v d n c e will receive an early satisfartory solution from this quarter. See, however, Heitler and London: Z. Physik, 44,455; Sugiura: 45, 484 (1927). For recent work see Perrin: Compt. rend., 185, 557 (1927); Xoyes: CohenFestschrift, Z. Physik. Chem., 130, 323 (1927); Lux: 2. physik. Chem., 121, 456 (1926). 8 Sidgwick: Chem. and Ind., 46, 799 (1927); “Electronic Theory of Valency,” (1927); Lessheim, Meyer, and Samuel: Z. anorg. allgem. Chem., 165, 2 j 3 (1927). 1 LOC.

* Z.

‘

1620

ROBERT E. BURR


way of finding out to what extent a bond is changed by outside influences, such as the fields of other molecules, or ions, substituents, etc. Some of this work has been previously discussed. Mecke' has attempted to correlate band spectra of various diatomic molecules and the periodic system of the elements. He calculated the binding forces in the various compounds using the potential law P = - [cl/rp - c2/rq], c l / P taking into account the attractive force and c2/rq the repulsive force. All the constants are spectroscopically accessible. Of the compounds examined, he concluded that the hydrides could be considered as polar since p = I (Coulomb's law). The remaining compounds he examined appeared not to be polar since p was found to be three or four, and q six to ten. I n addition to the work previously referred to regarding the influence upon bonds of substitution, etc., the following results have been obtained. Bovis* has found that the absorption spectra of bromine in water, carbon tetrachloride, chloroform, ethyl alcohol, and carbon disulphide consist in all cases of two absorption maxima, one of which lies in the violet a t approximate wave-length 0.410 p and the other in the ultra-violet a t approximate wavelength 0.260 p. The widths of the bands were unchanged from one solution to another. Scheibe, Felger, and Rossler3 show that the shifts in the wave-length of ultra-violet light absorption and changes in the amount of light of this wavelength absorbed when the same substance is dissolved in various solvents are two quite different phenomena. They attribute the first to solvation a t the active bond, and support this by cryoscopic measurements. The authors realize that in the spectral shifts one has a tool for getting a t the molecular deformation necessary to explain catalysis, although they did not determine just what the deformations were in the cases studied.

It is unfortunate that none of the work which the writer has been able to find upon band spectra enables one to decide whether substitution, solvent forces, etc., definitely alter bonds or not. I n some of the papers the position of a band attributed to a given link seems to persist practically unchanged from compound to compound, and from solvent to solvent. In other cases, there seems to be an appreciable, though usually small, shift, but it has not been traced definitely to a change in the binding force. The position and characteristics of a band, due to vibrations ( rotations) of a given link will, according to the quantum interpretation, depend not only upon the binding force and the law according to which it varies with distance, but upon the mass of the vibrating atoms. From the point of view of catalysis, what one would like to see brought out is the effect of substitution, etc., upon the strength of binding, distinct from any effect due to change in mass.

*

Z. Physik, 42, 390 (1927). Compt. rend., 185, 57 (1927) Ber., 60, 1406 (1927).

SIXTH REPORT OF THE COMMITTEE ON COSTACT

CATALYSIS

162 I


Condition of Adsorbed Gases It is thinkable that adsorbed gases can be essentially unchanged, can be stretched, can be resolved into atoms, can have their electronic orbits distorted, can be ionized, etc. Some of these possibilities have already been discussed . Resolution into Atoms.-H. S. Taylor' proposes resolution into atoms as the mechanism of activation for diatomic gases, e.g., hydrogen. It seems, however, to the writer that atomic hydrogen would probably be less active upon the surface than stretched molecules, because the bond catalyst-hydrogen atom would be stronger than the catalyst - H - H bond and would itself require activation in order that the atomic hydrogen become reactive. According to this view, there would be an optimum degree of stretchingwhich would give more active hydrogen than resolution into atoms. Kistiakowsky2 has continued the Princeton work on the investigation of the condition of adsorbed, gases on metal surfaces by the method of electronic impact. He felt confident that there is atomic nitrogen on the surfaces, but was not certain with regard to the hydrogen. If his conclusions are correct, it does not follow, however, that all the nitrogen on the surface is atomic, nor that the atomic nitrogen is the most reactive. Ions on the Surface.-The view is held in some quarters3 that the function of the catalyst is to ionize the adsorbed gases. The writer does not follow the argument that adsorption should necessarily lead to ionization, and can see no need for the assumption of ions as a preliminary state for chemical reactions in general. D h a f thinks reactions are hastened by the presence of ions. He considers these to combine with the reactants forming a more reactive complex. He proposes that heterogeneous catalysis is due to dissociation of the adsorbed gas into ions and electrons which then activate other molecules of adsorbed gas and induce the desired reaction. This is not impossible, as far as one can see, but neither does it seem likely as a general explanation of catalysis, for many catalytic reactions take place under conditions where ionization would seem most unlikely. In this connection, Hutchison and Hinshelwood5 find that the relative difficulty of decomposing ammonia and nitrous oxide by mblecular collision is maintained (roughly) when the collision is with ions. Pisarzhevskii6 thinks that in the catalytic decomposition of potassium chlorate by metals and oxides, the electrons of the catalyst ionize the adsorbed gases by their impacts. On leaving the catalyst, he thinks the gaseous 1

Proc. Roy. SOC., 1138, 77 (1927).

* J. Phys. Chem., 30, 1356 (1926).

3Bone: Proc. Roy. Soc., 112A, 477 (1926); Brewer: Phys. Rev. Proc. Kat. Acad. Sci., 12, j60 (1926). Z. anorg. allgem. Chem., 159, 103 (1926). 6 Proc. Roy. Soc., 117A, 131 (1927). 8 Chem. Abs., 21, 2415 (1927).

(2) 26,

633 (1925);

ROBERT E. BURK


1622

ions act on the molecules of the chlorate and bring about a series of electronic transferences resulting in the decomposition of the chlorate molecule. Gray’ suggests that clouds of electrons in the immediate vicinity of the surface are the active influence in promoting vigorous reactions a t surfaces. Bone2 examined the influence of drying upon the catalytic combustion of carbon monoxide over oxides and metals. The immediate effect of drying was always to increase the apparent catalytic power of the surface, and the ultimate effect to diminish or completely stop the catalytic combustion. I t is difficult to see how he concludes that these results would be best explained if the prime function of the catalyst surface is to ionize the reacting gases. These authors have not been very definite as to how they considered ionization of the reactants to explain catalysis. Also, their experiments, so far as the writer can see, have not demanded the ionization explanation. Severtheless, there is considerable evidence for electrical effects accompanying adsorption in some cases. Finch and Stimson3 have examined the charging of gold and silver surfaces when heated in the presence of various gases, and in vacuo. They conclude, (I) that the charge on hot gold or silver surfaces in contact with gases is due to an activation of the gas whereby its molecules become electrically charged, ( 2 ) that the activation of a compound molecule such as water-vapor or carbon dioxide involves a t least its dissociation. These conclusions can be true if the definition of activated molecules is different from that used for purposes of reaction velocity, and there is nothing in the experiments to show that all molecules chemically active have been activated in such a way as to involve electrical charging or complete dissociation. B r e w e ~has found that the heterogeneous oxidation of xylene, toluene, benzene, ethyl alcohol, etc., produced currents of I X IO+ amperes per square centimeter of reaction surface a t temperatures of 580’ - 750’ Abs. He found no current with the gases separately. He postulates nevertheless that the adsorbed gases are ionized by the “combined image and contact forces.” A current of I x 10-15 amperes per square centimeter is an extremely minute one. Becker5 has demonstrated the existence of both caesium atoms and ions upon the surface of a tungsten filament in the presence of vapor of that metal, and is able to estimate the proportion of each. He found that the wire attained a maximum in electron emission when covered with a monomolecular layer of caesium atoms. Below a certain critical temperature, a second adsorption layer started to build up. Kunsmad finds that the active Fc-X1203-Alkaline oxide catalysts (for ammonia) are excellent emitters of positive (alkali) ions. These catalysts

* Chem. and Ind., 46, 218 (1927,. 2

LOC. cat. Proc. Roy. Soc., 116.1, 379 (1927).

4 LOC.

nt

Phys. Rev., ( 2 ) 28, 3 4 1 (1926). J. Franklin Inst., 203, 635 (19271.


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give large and comparatively steady emissions. When A1203 is absent the emission is larger but less steady. He says, “It would seem quite reasonable, therefore, that one should look for a partial explanation, a t least, of the increased catalytic activity of an Fe - A1203 - Alkaline oxide mixture to the presence of the alkali on the surface, either in terms of an interface phenomenon, a reduction of the electron work function of the surface, the emission of postive ions with a potential or neutral vapor without a potential.” However, the presence of alkali decreases the activity of iron alone. He suggests that the alkali diffusing to the surface might act by keeping the surface in a state of constant eruption. If so, this effect would be present for iron alone. The effect of the alkali is very specific. It is altogether possible that the presence of an ion would affect the reactivity of some molecules either by atomic distortion (already discussed), or by separating the atoms in case they differin electrical polarity. The widespread catalytic effect of hydrogen ions may well be due to this cause. Alany of the foregoing experiments were conducted a t elevated temperatures, where the thermal emission of electrons and ions is well enough established. These phenomena disappear a t room temperature, except when enormous fields are applied to the surface. The experiments, therefore, cannot be used to demonstrate the existence of ions on the surface a t room temperature. Furthermore, it is quite possible, in some of the experiments, that the substances do not exist upon the surfaces as ions, but that the ions are created by the very collision which ejects them from the surface. The writer’s feeling is that to assume that substances are always adsorbed as ions, or that the preliminary formation of ions, is a prerequisite in catalytic processes is, t o say the least, a very extreme view. At room temperature one might expect a molecule to remain unionized upon being adsorbed, a t least in many cases, yet catalytic effects at room temperature are numerous.

Differential Heats of Adsorption Some very interesting results have been obtained during the year by measuring the heat of adsorption for successive small portions of added gas. Beebe’ found that these “differential” heats of adsorption for carbon monoxide on copper were distinctly higher for the initial portions than for later portions. After accidental poisoning, resulting in decreased adsorptive capacity, the integral heat of adsorption was greater than on the unpoisoned catalyst. Beebe concluded that the most active points on the catalyst are not the most unsaturated ones. Magnus and Kalberer2 found that the heat of adsorption of carbon dioxide on wood charcoal was about 7450 calories at pressures greater than 5 mms., but rose to 12,460calories a t 0.076 mms. pressure. They interpreted this as being due to points of high adsorption potential. With carbon dioxide on silica, the heat of adsorption decreased linearly with increase in equilibrium pressure, which they attributed to the Joule-Thomson effect. J. Fhys. Chem., 30, 1538 (1926).

* 2. anorg. allgem. Chem., 164, 345 (1927)


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Garner’ has found a maximum in the differential heat of adsorption curve for oxygen on charcoal. Xistianowsky,* Taylor and Kistiakowsky? and Kistiakowsky, E’losdorf and Taylor; have found that the differential heat of adsorption of hydrogen on active copper and nickel catalysts is small for the first portions of gas adsorbed. It rises sharply to a maximum and then settles to a steady intermediate value as more gas is added. Heat treatment shifted the maximum toward smaller amounts of added gas, and made it less pronounced, and finally eliminated the maximum altogether. When the catalyst was treated with oxygen, the initial value of the heat of adsorption was high, and there was no maximum. They interpret the results (‘as indicating in the active preparations a greater fraction of the surface capable of activating the adsorbed species, the activating process being endothermic, and possibly a dissociation into atoms. The heat treatment destroys these active areas on the surface preferentially.” This seems probable enough. However, they go on to say, “A difficulty of theoretical importance in connection with these results on the heat of adsorption on active catalysts may now be stated in the hope that some suggestions may be forthcoming as to its solution. It has been pointed out to us by Professor Herzfeld of Johns Hopkins University that, with the active preparations, the process first occurring is one in which the decrease in internal energy is very much less than that occurring subsequently a t the maximum point and beyond. There is no evidence that the entropy changes involved would cause the free energy decrease a t low adsorptions to be greater than those obtained a t the maximum. The problem therefore arises as to why the process with a smaller free energy decrease should occur preferentially to that in which the large free energy decrease can take place.” The writer does not see the thermodynamic difficulty, and feels that multiple adsorption might account, a t least in outline, for the results. There is a certain kinetic complexity involved in desorbing the multiply adsorbed molecule, which must be taken into account as well as the energy factor. This is not an unknown factor in thermodynamics for, according to the kinetic nature of the process, the terms involving the sum of the chemical constants can outweigh the term involving the heat of reaction in calculations of equilibrium. Let us consider the molecules attached to the surface by two atoms. Judging from the distribution of the velocities of thermally emitted electrons, one can conclude that there is a Maxwellian distribution of energy amongst the surface atoms. There are two adsorption bonds to be broken before the adsorbed molecule can leave the surface. This can be done in two ways. First, one atom can receive a thrust hard enough to endow it with sufficient energy to break both bonds. However, if E is the energy necessary to break Garner and McKie: J. Chem. SOC., 130, 2451 (1927). Sci., 13, I (1927). SZ. physik. Chem., 125, 341 (1927). 4 J. Am. Chem. SOC., 49, 2200 (1927). 1

* Proc. ”et. Acad.


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each bond, the energy with which one atom must be endowed in order to break both bonds may be considerably greater than 2E, for energy will be used up in stretching the adsorbed molecule. The doubly adsorbed molecule can also leave the surface if each adsorbed atom receives a thrust endowing it with energy a t least equal to E. However, in this case, unless both atoms receive this energy within a very short time, the first atom will be adsorbed again. It is a little difficult to calculate the exact magnitude of this effect. However, if the time interval within which both atoms must receive E is short enough, it is clear that the molecule will be desorbed less often than a singly adsorbed molecule whose energy of desorption is zE.

Amount of Adsorption and Catalytic Activity Sabatschka and Moses' verify the point set forth a t the beginning of this report that there is no quantitative correspondence between amount of total adsorption and catalytic activity. Experiments of Griffin2 also support this point. However, Taylor and Kistiakowsky find that two active methanol catalysts show quite extraordinary adsorptive properties for hydrogen and carbon monoxide. Intermediate Compounds Sabatier; whose personal experience in catalysis is very great, appears t o think that the formation of intermediate compounds is a satisfactory explanation of catalysis. However, his point of view does not remain unchallenged5 and one cannot restrain a desire to know what compounds are formed in specific instances (particularly when it seems unlikely that any orthodox compounds should be formed), and what there is about them which makes them so reactive. Perhaps a pursuit of such questions will lead to the fusion of the adsorption and intermediate compound points of view. There are cases where the hypothetical intermediate compounds cannot exist except on the surfaces. I n these cases, perhaps, the postulate is unfruitful. ZhukowG has obtained some interesting results along this line. He heated a number of metals in nitrogen. Magnesium began t o absorb nitrogen a t 700°, Mg&;, being formed; calcium reacted a t the same temperature to form Ca3RT2; lithium formed I&S. These three compounds show no measurable dissociation pressure below I 2 50'. With manganese and chromium, isotherms were obtained representing equilibria with solid phases of continuously varying composition (910'- 1 2 0 0 ~ ) . Aluminum began to form AlN a t between 850' and 875'. Impure titanium behaves like manganese and chromium. Ber., 60B,786 (1927). J. Xm. Chem. Soc., 49, 2136 (1927). J. Am. Chem. Soc., 49, 2468 (1927). Ind. Eng. Chem., 18, 1005 (1926). Boeseken: LOC.cil. Ann. Inst. anal. physik. Chem., 3, 14 (1926).


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Molybdenum absorbed only 2 0 cc. of nitrogen per five grams of metal, and uranium 2 7 cc. of nitrogen per four grams of metal a t 1000’. No reaction below 1200’ was observed with tungsten, zinc, copper, and iron. The electrical conduction of azotized chromium, manganese, and titanium was of the same order as that of the pure metals; that of AlN, CasKz, and Mg,X;, was less than 5 X IO-?. There seems to be a distinct difference between the nitrides and the azotized metals. Sabatier and Fernandezl have found that various difficultly reducible oxides are able to effect the hydrogenation and dehydration of organic compounds. Intermediate compounds seem rather a useless conception here. According to BalareffJ 2 alcohols merely add to phosphorus pentoxide below 130’ forming normal phosphates and acid phosphates. No catalytic decomposition takes place. At 130’-15o’, the addition still takes place, but the acid phosphate thus formed decomposes into ethylene and a volatile ester of phosphoric acid. At from 16oo-26o0,only part of the alcohol is combined with phosphorus pentoxide and catalysis takes place partly by formation of intermediate compounds and partly by adsorption and resulting thermal decomposition. At 3g0°, the alcohol is only adsorbed by the phosphorus pentoxide and the catalysis is accomplished not through the formation of intermediate compounds but by thermal dissociation of the adsorbed molecule. At this temperature no volatile ester of phosphoric acid is found in the reaction products. He considers that no sharp dividing line can be drawn between the two mechanisms.

Active Centers and the Structure of Surfaces At the beginning of this report, it was set forth as an established fact that all catalytic surfaces are not uniformly active. T a y l ~ rin , ~the Fourth Report of the Committee on Contact Catalysis, developed the view that the most active centers are peaksof “extra lattice” atoms of high degree of unsaturation and that the edges and surfaces of crystals are less active according to their degree of unsaturation. This is a theory of catalytic surfaces, not of catalytic action, and Taylor did not claim otherwise for it. It does not qualify as a theory of catalysis because catalytic action at the peaks would be no more lucid than on a catalyst of uniform activity. I n the writer’s opinion, the work on active centers (much of which is due to Taylor and his collaborators) by no means forces the conclusion that active centers are always necessary, nor even that when they exist they are peaks of “extra lattice” atoms, etc.4 The interfaces between minute crystals in a catalyst which without assumption must exist, would serve just as well as active centers. I n other cases, the results apparently indicating active centers can be explained on plane surface^.^ I n still other casesJ6there is evidence Compt. rend., 185, 241 (1927). zZ. anorg. allgem. Chem., 158, IOj (1926). * J. Phys. Chem., 30, 145 (1926). J. Phys. Chem.; 30, 1134 (1926). 6 This view is also held by Constable: Proc. Cambridge Phil. SOC.,23, 832 (1927). 6 Miss Wright: Proc. Cambridge Phil. Soc., 23, 187 (1927).



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that the active centers consist of pairs of different atoms (predictable on the basis of multiple adsorption nevertheless, it has not been shown that Taylor’s “extra lattice peaks,” etc., never exist. Constable1 has decomposed copper salts of monobasic organic acids and investigated the catalytic action of the newly formed copper, measuring the rate of reaction by the rate of evolution of hydrogen. He concluded that the catalytic activity was independent of the nature of the compound from which the copper was derived.* He considers these results to agree with the hypothesis that the centers of activity of the surface are “frozen groups of atoms with strong specific external fields.” This idea of an active center is not up to the standard of Constable’s experimental work. If Constable objects to pursuing the question a little further and saying that on the grounds of the short range of surface forces, etc., multiple adsorption is a likely interpretation of the specificity, etc., of some active centers, he has not stated what his objections are. On the basis of the multiple adsorption mechanism, one would have some idea of how t o set about creating more active centers. I n an investigation of the nature of the sintering of copper catalysts, Constable3 concluded that the coarse structures collapse even more easily than the fine ones, which he says shows that the energy present in the centers of activity is not very greatly in excess of that possessed by the regular arrangement of the surface atoms. Miss Wright4 considers her data on the auto-oxidation of certain charcoals to show that the oddaton proceeds along chains of carbon atoms. The reaction is zero order, andonly 0.38% is auto-oxidizable according to her measurements, Rideal and Wright; in an investigation of the oxidation of oxalic acid by charcoal, consider that blood charcoal contains two types of active centers, which they think are Fe - C, and C - C complexes, and that sugar charcoal6 contains three types of active centers, attributed to Fe - C, C - C, and Fe - C - N complexes. They estimate the Fe - C complex to be 57 times as active as the C - C complex (on blood charcoal). The more active area on this charcoal was estimated, by the method of selective poisoning with KCN; to occupy 1.2% of the total surface area (estimated by Paneth’s method, based upon the assumption that a surface is completely covered, one molecule deep, with methylene blue, when in the presence of that dye). There are many points to be watched in the method of selective poisoning, e.g., steric effects not dependent upon the active centers; the quantity of poison must not be sufficient t o cover the catalyst one molecule deep; one must be sure that the poison itself does not have some catalytic activity. l

Proc. Cambridge Phil. SOC.,23, 432 (1926).

* Compare the work of Adkins: Second Report of the Committee J. Chem. SOC.,130, 1578 (1927).

‘L O C . cit.

J. Chem. SOC.,129, 3182 (1926). J. Chem. SOC.,127, 1347 (1925);129, 1813(1926).

on Contact Catalysis.


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Rideal and Hoover’ conclude that different patches upon thoria are active in the dehydrogenation and dehydration of ethyl alcohol. According to Constable* reduced copper has an activity 8,000 times greater than that of electrolytic copper. Bredig and Allolio3 conclude from X-ray studies that, except for particle size, there is no difference between the finely-divided, highly-active platinum, palladium, and nickel catalysts and inactive catalysts of these metals. Madenwald, Henke, and Brown4 find that different preparations of lead, after use, assume the same activity for the hydrogenation of nitrobenzene. This may be purification of the surface, chemical change, or structural change. According to Adam,5 it is unnecessary to suppose that liquefaction occurs a t a surface upon polishing? Von Weimarn and Hagiwara’ claim that all rigid substances are crystalline. If amorphous patches in metallic catalysts are in all cases impossible it would be well worth knowing.s

Promoter Action This is the name given to the phenomenon, sometimes observed, that the activity of a mixture of catalysts exceeds the additive effects of both. Promoter action, while of the greatest importance to catalysis in general, may be conveniently discussed in connection with the problem of the nature of active centers. Constableg has found that in the dehydration of alcohol over thoria, water acts as a promoter in small concentrations but becomes a poison a t higher concentrations. This would be simply explained if the reaction occurred a t the border between the water film and the underlying catalyst, a possibility which multiple adsorption suggests. Water would then become a poison as the film becomes complete in case the water itself has but small catalytic power for this reaction. Constable may not accept this explanation. KunsmanlO in an investigation of the thermal decomposition of ammonia on promoted iron catalysts found the heats of activation (using the time of half-life as a measure of the reaction rate) to be between 38,000 and 42,000 calories for all the catalysts examined, but that the catalytic activity varied as much as eighteen-fold. He concludes that the primary action of promoters is to increase the number of atoms upon which decomposition takes place and that heat treatment decreased this number; and that the quality of the centers upon which reaction takes place is not altered sufficiently to cause J. Am. Chem. SOC., 49, 104 (1927). *Nature, 119, 349 (1927). Z. physik. Chem., 126, 41 (1927). J. Phys. Chem., 31, 862 (1927). 6 Xature, 119, 162 (1927). e Beilby: “Aggregation and Flow of Solids,” (1921). Kolloidchem. Beihefte, 23, 400 (1927). a See, however, Langmuir: J. Am. Chem. SOC.,2252 (1916). uproc. Cambridge Phil. SOC., 23, 593 (1927). Wcience, 65, 528 (1927).

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an appreciable change in the heat of activation. He considers that the heat of activation may be a function only of the reacting gas and not of the catalyst. The variation in the heat of activation which he records is sufficient to account for an eleven-fold change in rate, and it depends entirely upon the circumstances of the action (unpublished) whether the measured heat of activation can be taken as a measure of the true heat of activation or not. His last conclusion would seem most unlikely if the different catalysts lower the heat of activation of ammonia, and there is extremely good evidence1 that they do. In promoter action as ordinarily defined, the quality of the active centers is changed, for the striking feature about promoter action is that the reaction velocity with the mixture is not that which would be expected by adding the velocities for the separate ingredients, nor by producing one new active center, of a kind effective on one of the separate ingredients, for each new particle of the new ingredient. The thermal decomposition of ammonia upon platinum, a t I O O mms. pressure, is retarded by hydrogen, but not by nitrogen. That on tungsten, is not retarded by hydrogen.2 The decomposition upon molybdenum3 is retarded by nitrogen and not by hydrogen. It would be a good guess, therefore, that tungsten adsorbs the nitrogen atom of ammonia more strongly than the hydrogen atoms and vice-versa on platinum (there is considerable other evidence in support of this). On the basis of multiple adsorption, therefore, the decomposition of ammonia upon a mottled surface of tungsten and platinum should be more rapid than upon either alone. The write^ has recently evaporated platinum from a spiral on to a tungsten filament in high vacuum, and found that the rate of decomposition on the mixed surface was unmistakably greater than on the underlying tungsten; and when tungsten was evaporated on to platinum, the activity of the mixed surface reached a value unmistakably greater than upon the underlying platinum. It is doubtful that the surface after the evaporation could accommodate many more molecules of ammonia than before. Although the experiments have not been completed, this seems a clear case of promoter action which was predicted theoretically. The magnitude of the promoter action was not as great as is often found (the rates on the mixed surfaces were about twice as great as upon the unalloyed surfaces). Almquist,j also, thinks that the main function of promoters is to increase the number of catalytically active atoms. Again, he must define promoter action differently from the writer. QuartaroW has found that Mg(OH),, Cd(OH), and Iii(OH)e greatly diminish the rate of decomposition of alkaline hydrogen peroxide and greatly diminish the power of Pb(OH), to accelerate the decomposition of the subDiscussed under “Lowering of the Heat Activation.” Hinshelwood and Burk: J. Chem. SOC.,127, 1105 (1925). a Burk: Proc. Nat. Acad. Sci., 13, 67 (1927). Experiments soon to be published. J. Am. Chem. SOC.,48, 2820 (1926). 6 Gam., 57, 234 (1927).


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stance. But these same compounds greatly increase the accelerating action of Ag20on the decomposition of hydrogen peroxide. Quartaroli’s explanation is that the added substances protect or destroy the real catalyst (not named). Yant and Hawk’ have investigated the catalytic oxidation of methane by metals and oxides and mixtures of these. Promoter action was not the rule here, although mixtures of C O Z Oand ~ nickel were almost as efficient as Co203alone, nickel alone being quite inefficient. Taylor and Kistiakowsky2 say, “The mode of action of promoters can be traced to the following causes, ( I ) protection of the catalyst from sintering, and in this way the adsorbing surface, especially the number of unsaturated atoms; ( 2 ) the place of contact between different chemical substances may be the seat of special actions upon the adsorbed molecules and thereby give rise to selective effects; ( 3 ) the possibility is not excluded that the added substance has chemical action, e.g., in breaking up the intermediate compound formed Rith the catalyst.” In point ( I ) these authors are considering phenomena different from true promoter action, as defined above for one particular phenomenon, e.g., the effect of “supports” may be in part that given in ( I ) . It seems to the writer that the most likely interpretation of point ( 2 ) is multiple adsorption? Surely the promoter effect would not be attributed to the force at a mathematical dividing line between two atoms. The individual forces of the two atoms will be the more intense the nearer one approaches the respective atoms. The effective specific force would then become a pair of forces with definite strength and spacing dependent upon the nature of the individual atoms. Promoter action is a most natural consequence of multiple adsorption.4 It is a little difficult to say what the profitable attitude5 toward point (3) should be. When there is more than one molecule reacting, the promoter may act simply by placing the second molecule is a suitable juxtaposition more frequentlyP Supported Catalysts Tropsch’ has found that a t 355’ the dissociation of formic acid {waterfree) to carbon dioxide and hydrogen occurs most readily over tinned iron. There is a possibility regarding supported catalysts which does not seem to have been recognized. The possibility is that in addition to the prevention of sintering, in the case of thin films, as in the foregoing instance, the support can enforce its spacing upon the catalyst film, which may be more suitable than the normal spacing of the catalyst. Morns and Reyersons have found J. Am. Chem. Soc., 49, 1454 (1927). Z. physik. Chem., 341,125 (1927). See the experiments upon the thermal decomposition of ammonia described above. Burk: J. Phys. Chem., 30, 1134 (1926). See the discussion of intermediate compounds in this report. Rideal and Taylor: “Catalysis in Theory and Practice ” p. 31 (1919); Langmuir (Trans. Faraday SOC.,17,607 (1922))discusses this type of s p a h g effect. See also Hinshelwood: “Kinetics of Chemical Change in Gaseous Systems,” p. 195 (1926). Abhandl. Kenntnis Kohle, 7, I (1925);Chem. Abs., 21, 3530 (1927). J. Phys. Chem., 31, 1220 (1927);32, 113 (1928). 3

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that thin films of metals upon silica gel have marked catalytic activity; whereas, the experiments of Gauger1 will be recalled who found thin films of nickel or platinum on Pyrex to be inactive (there is no apparent reason for “extra lattice” atoms to be precluded on the Pyrex-metal films.)


Films on Catalysts The rather disheartening point arises that the real catalysts may not be what we think they are. The writer has observed that if a tungsten wire, which is in a perfectly steady state of activity for decomposing ammonia, be heated, at the same temperature used for decomposition, in a high vacuum, its activity falls off markedly, but is regained upon decomposing more ammonia. The steadiness which the wire attains-which is easily reproduced --would make an actual change in structure with decomposition unlikely (though not impossible) in this case. The probable explanation is that the tungsten becomes covered with a film of nitrogen, possibly atomic. This conclusion is supported by the work of Kenty and Turner2 who find that active nitrogen produces a film on tungsten, when this metal is maintained a t dull red heat. The writer has also found evidence for the existence of a similar film upon molybdenum3 when used for the thermal decomposition of ammonia. Bone and Forshawl find that the catalytic action of fireclay upon the combustion of carbon monoxide and hydrogen is stimulated by previous exposure to the combustible gases. Spitalsk3.5 found that electrochemical pretreatment of platinum with oxygen decreased its activity in decomposing hydrogen peroxide, while cathodic polarization increased it. Thus a film of hydrogen on platinum has a beneficial effect when used for this reaction, and a detrimental effect when used for the decomposition of ammonia. There is much evidence to show that many catalysts behave quite differently when the films of moisture are removed from them.5 Langmuir has shown how tenaciously films of oxygen are attached to tungsten.’ Clark and Toplef encounter the film question in experiments on the catalytic decomposition of formic acid. The point has been subjected to special investigation for the case of water synthesis over nickel by Hughes an& B e ~ a n . ~ J. Amer. Chern. SOC.,47, 2278 (192j). Nature, 120,332 (1927) Proc. S a t . Acad. Sci., 13, 67 (1927). Portions of this paper have been cailed into question, and the research has not been fully repeated. Proc. Roy. Soc., 1144, 169 (1927). Ber., 59,2900 (1926). Bone: Proc. Roy. SOC.,112.4, 477 (1926). ‘Trans. Faraday SOC.,17, 618 (1922). J. Phys. Chem., 32, IZI (1928). Proc. Roy. SOC.,117A, I O I (1927). 2


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It is therefore well to consider whether a given catalyst is ordinarily bare (except for loosely adsorbed substances), or whether the reaction is probably taking place on a stable film, which may vary in nature according to the reaction catalyzed. Poisons Constable’ has investigated the effect of inert diluents on the rate of decomposition of alcohol over copper; he found their behavior to fit his previously deduced formula.2 If the heat of adsorption of the reactant is greater than that of the diluent, the temperature coefficient of the reaction is less than the true one, and vice versa. He says, “this is a hitherto unsuspected cause of error in temperature coefficient measurement.” The point is well taken, but is not new.3 Selective poisoning has been discussed under “active centers.” There is nothing very mysterious about the action of poisons. Dust Particles According to Rice4 many reactions would take on a different aspect if extreme precautions were taken to remove dust particles which act as catalysts. Rice and Kilpatrickb find the decomposition of hydrogen peroxide to be such a reaction. Rice and Getz6 find that the thermal decomposition of nitrogen pentoxide is unaffected by carefully removing the dust particles. The reactions which have been found to be homogeneous and kinetically accountable would not be expected to depend upon the presence of dust particles. It would not seem that dust particles offer a very great menace to the sound investigation of reaction velocity; surely dust particles cannot be universally effective as catalysts, and their effectiveness would not often be expected to compare in magnitude with that of the walls, which are probably similarly constituted, and of larger area, and with that of added specific catalysts. Moisture The catalytic effect of water vapor stands in a similar, though probably more important position. Garner and Johnson’ find that the presence of moisture greatly depresses the emission of infra-red radiation from exploding carbon monoxide-oxygen mixtures. Their interpretation is that the water increases the rate a t which thermal equilibrium is reached, energy being conserved which otherwise would be radiated away. They call this “energo-thermic” catalysis. If the energy is used merely to raise the temperature of the system, it is questionable whether it should be considered a special type of catalysis, any more Proc. Cambridge Phil. SOC.,23, 593 (1927). Phil. SOC.,23, 172 (1926). 3 Hinshelwood: “Kinetics of Chemical Change in Gaseous Systems,” 178 (1926). See Fifth Report of the Committee on Contact Catalysis. 5 J. Phys. Chem., 31, 1507 (1927). 6 J. Phys. Chem., 31, 1572 (1927). ’Phil. Mag., 3, 97 (1927).

* Proc. Cambridge


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than highly reflecting walls should be. The authors did not state whether or not they considered water molecules to be necessary in the sphere of reaction. Armstrong’ denies that water in this reaction reacts directly with the carbon monoxide. His objection is discussed later. New experiments of Smits2 indicate that the density of dried ammonium chloride is greater, and the vapor pressure smaller than for the moist substance below 3oo0C. He considers that his experiments show that intensive drying causes a shift in an inner equilibrium. Now the amount of moisture which will cause ammonium chloride to behave normally is less than one part in a million, as the writer knows from personal experience with this problem. This amount of moisture cannot cause decided shift in an equilibrium by reacting with one of the species. If the change in the equilibrium is caused in some other way, it would seem that a negligible free energy of reaction would have to be assumed, or that an enormous amount of energy would be involved in removing the last traces of water. The explanation in general does not strike one as plausible. In an earlier paper, in discussing the possibility of the displacement of an equilibrium Smits3 says, “It is tacitly assumed that a trace of moisture can displace the inner equilibrium to a very considerable extent, the thermodynamic significance of which is, that a very large amount of work is necessary to withdraw the last trace of water from the substance, and this is not improbable.” He does not say why it is not improbable, and it certainly seems improbable on the face of it. Smits4 has also found that intensive drying increases the vapor pressure of nitrogen tetroxide by as much as 3.5 cms., the change being accompanied by a deepening of the brown color. He found intensive drying to decrease the vapor pressure of hexane, benzene, carbon tetrachloride, and carbon disulphide. Again he suggests shifts in the inner equilibrium. Smith5 also found that intensive drying increases the vapor pressure of nitrogen tetroxide and agrees with Smits as to interpretation. Balareff6found that after standing in contact with phosphorus pentoxide for three and one-half years, methyl alcohol, ethyl alcohol, benzene, hexane, and ether had added phosphorus pentoxide forming phosphates, which remained as a residue after evaporating the liquid. He thinks these addition compounds, or phosphates, may account for the results on the change of vapor pressure of liquids with intensive drying. However, Baker’ says his experiments do not admit this possibility. xature, 120, 659 (1927).

* Rec. Trav. chim., 46, 445

‘ 6

’

(1927).

J. Chem. SOC.,125, 1069 (1924). J. Chem. SOC., 129, 2655 (1926). J. Chem. Soc., 130, 867 (1927). J. prakt. Chem., 116, 57 (1927). J. Chem. SOC.,130, 2902 (1927).


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Briner’ has performed experiments which make him doubt that water plays an essential part in the oxidation of nitric oxide, a result previously reported with not too great certainty by Baker*. This is not the first time Baker’s results have been called into question,3 and he has made good his claims on the other occasion^.^ Nevertheless, the oxidation of nitric oxide is a homogeneous reaction satisfactorily accounted for kineti~ally,~ and there is, therefore, no reason for expecting it to depend upon the presence of watervapor. Briner points out that the velocity of this homogeneous reaction is greatest a t liquid air temperatures, where the concentration of water-vapor is extremely small. According to experiments of Cohen and Heymer; the combination of hydrogen and chlorine is not a wall effect, and the inhibition observed upon drying, therefore, cannot depend upon removing the water from the walls. Negative Catalysts or Inhibitors Backstrom’ has carried out some interesting experiments on negative catalysis in the oxidation of aldehydes and of sodium sulphite. He found negative catalysis both for the light and dark reactions, and found quantum o per efficiencies for the light reactions of the order of ~ o , o o o - ~ o , o omolecules quantum. His results led him to come out for Christiansen’s theory. of negative catalysis. Christiansen’s t h e o g ie an extension of his “hot molecule” mechanism,9 which he used to account for reactions with otherwise necessary high rates of activation. The idea has been shown theoretically by TolmanIo to be of doubtful value for some types of reactions, and has been shown experimentally by Hinshelwood and Hughes” to be most improbable in the favorable case of the thermal decomposition of chlorine monoxide (which easily becomes explosive). Christiansen believes that because the molecules of reaction products, when first formed, possess high energy content, they can activate new molecules of reactant a t the first encounter, the process being repeated, thus giving rise to long reaction chains. He then proposes that negative catalysts act by breaking up these chains, either by taking up the energy from the “hot” molecules of reaction products, or by reacting with them in some way or other. ~

J. Chim. phya., 23, 848 (1926). * J. Chem. SOC.,65, 613 (1894). 3 Gutmann: Ann., 299, 267 (1898). ‘Baker: J. Chem. SOC.,73, 422 (1898). Bodenstein: 2. physik. Chem., 100, 68 (1922). Ber., SQB, 1794 (1926). J. Am. Chem. SOC., 49, 1460 (1927). 8 J. Phys. Chem., 28, 145 (1924). 9Christiansen and Krsmera: 2. physik. Chem., 104, 451 (1923). ’OJ. Am. Chem. SOC., 47, 1524 (192j). “J. Chem. SOC., 125, 1841 (1924). See also Hinshelwood: Chem. Rev., 3, 227 (1926).



1635

The writer feels that there are certain objections to such a mechanism for negative catalysis, in addition to those referred to above, and others‘ for the general scheme of “hot” molecule chains in slow thermal reactions, which should be disposed of before accepting the chain mechanism in any specific case. First, the amounts of negative catalyst necessary are often minute, which in the ordinary case would seem to rule out reaction with the “hot” molecules, because the supply of negative catalyst would soon be used up.2 Secondly, there are present, after the reaction has proceeded a small distance, molecules of products in enormously greater numbers than those of the negative catalyst. These would also rob the newly formed molecules of products of their energy, and while it may be argued that they would then be as effective as the original “hot” molecules of products in activating new reactant molecules, nevertheless, if the energy passes from product molecule to product molecule in this way before being transferred to a reactant molecule, the chance of breaking up the chain through distribution of the energy would be very great. The transfer in this way from product molecule t o product molecule cannot be ruled out on the groundsof specificity in the transfer of energy on collision, since they are molecules of the same kind. Thirdly, if the molecules of negative catalyst really do rob the “hot” product molecules of their energy, they would, as far as the writer can see, be on precisely the same basis as molecules of products, unless additional assumptions are made, and could communicate the “stolen” energy to reactants as well as the “hot” molecules whose energy they have just taken up.3 It is not true in general that the molecules of inhibitor are more complicated than the molecules of product, which might otherwise cast doubt upon this possibility. I n other words, if the “cold” products do not themselves act as negative catalysts, why should the negative catalysts act as such? The old and simple idea that a negative catalyst acts by destroying the activity of a positive catalyst is difficult to dispose4 of, and its incompetence must be shown in each individual case, when negative catalysis is restricted to the case of a given reaction actually being slowed up (not replaced by faster side reactions, etc.). F. Perrin5 considers the action of antioxygens to consist in the selective deactivation of the active molecules. MujamotoB has found that the velocities of oxidation of mixtures of stannous chloride and sodium sulphite in sodium hydroxide solution are

+

+

E.g., with a reaction of the type A = A’ = R’, A = A‘ R, A’ and R‘ are probably indistinguishable for the purposes of reaction velocity so that the whole theory loses point in this case. * (Sote added to the proof.) Prof. H. S. Taylor informs me that since the completion of this manuscript, a great many results have been obtained in his laboratory which can best, be interpreted in terms of the chain mechanism. The inhibitor, or negative catalydt, in some of these experiments was itself destroyed by oxidation, but a t an extremely slow rate. (One molecule for some fifty thousand molecules of sulphite normally oxidized). The very slow endothermic destruction of the inhibitor would also account for a point raised a t the end of this paragraph, See, however, footnote 2 . ‘See the section on moisture in this Report. Compt. rend., 184, I I Z I(1927). 6Bull. Chem. Soc., Japan, 2, 191 (1927).

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ROBERT E. BURH


less than those of either alone, except when the concentration of stannous chloride is low. Cases of negative catalysis found by Tropsch and von Philippovichl and by Quartarol9 have been discussed under “promoter action.”

Homogeneous Catalysis Since contact catalysis is the subject of this report, homogeneous catalysis will not be elaborately considered. Nevertheless, an account of the status of contact catalysis would suffer if homogeneous catalysis were completely neglected. Dawson2 and his eo-workers have carried out an extensive series-of investigations on catalysis in solution. Some of their conclusions are the following. The “protion” theory of chemical reactivity is inconsistent with the facts. Catalytic effects must be attributed to the hydrogen ion, the undissociated acid, the acid ion, and the hydroxyl ion. Catalytic activity is proportional to the volume concentration of molecular or ionic catalyst, and has no apparent connection with the thermodynamic activity. Skrabal? also, takes into consideration the catalytic activity of water molecules, of hydrogen ion;, and of hydroxyl ions. He finds that the catalytic effect of hydrogen ions is relatively constant (for the saponification of different esters) but that the catalytic effect of the hydroxyl ions varies over extremely wide limits. It is interesting in this connection, that an association related to multiple adsorption is possible for the hydroxyl ion but not for the hydrogen ion. Bronsted and Guggenheim4 do not ascribe the catalytic activity of acids and bases to the hydrogen and hydroxyl ions, but to the acid and base molecules. The catalytic efficiencies of acids and bases were found to be proportional to their strengths (for the mutaroration of glucose). Bergstein and Kilpatrick5 and Bergsteins find that the catalytic minimum point (the point of minimum velocity), for the reaction between iodine and acetone, was not displaced by neutral salts, nor by temperatures within the limits 25°-4S0. They find the reaction to be complex, and were unable to identify all the catalysts with certainty. According to Baudisch and Davidson,? the ferrate and ferrite ions have identical catalytic activity for the oxidation of certain organic compounds. 2 I- = 2S03 iIP, Kiss and Z o m b o g think that the reaction, &Oscatalyzed by iron ions, takes place in accordance with the mechanism:

+

1 LOC.

cit.

Dawson and Hoskins: Proc. Leeds Phil. LBt. Sci. Sect., 1, 108 (1926);Dawson and Dean: J. Chem. Soc., 129, 2872 (1926);Dawson and Carter: 129,2282 (1926);Dawson: 130,458, 1290 (1927);J. Phys. Chem., 31, 1400 (1927). a 2.Elektrochemie, 33, 322 (1927). J. Am. Chem. SOC.,49,2554 (1927). 6 J. Phys. Chem., 30, 1616 (1926). 6 J. Phys. Chem., 31, 178 (1927). 7 J. Biol. Chem., 71, 501 (1927). *Rec. Trav. chim., 46, 225 (1927). 2


+ +

+

2 Fe+++ 21- = 2 Fe++ IZ Feft SzOs- = 253042 Fe* Akerlofl and others have attempted to correlate reaction velocities in solution with the :ctivities of the catalysts. For the decomposition of dincetone alcohol, Akerlof finds that the reaction velocity is proportional to the activity of the catalyst when the latter is potassium hydroxide, or sodium hydroxide. When lithium hydroxide is the catalyst, however, the velocity is accurately proportional to the conceqtration and not to the activity. He was unable to obtain a definite relationship between activity and reaction velocity for the hydrolysis of ethyl acetate and of cyanamide. A great many straight lines are shown in the paper with only two points on them. Thermodynamics is not concerned with the rates a t which systems approach equilibrium. If activities are the numerical coe5cients which one must substitute for concentrations for thermodynamical purposes, i t would seem that they should be irrelevant to the subject of reaction velocity, and that the volume concentration, a priori, should be the important thing, for it is the volume concentration which comes into the expression for the number of collisions. If the non-ideal behavior of one species dissolved in another is due entirely t o the real concentration not being what we think it is, and if the activity coefficient is a measure of this divergence and of nothing else, then one should use activities in expressions for reaction velocity--otherwise, not. If a single molecule can exhibit different activities in different thermodynamic environments, as presumably it can, then the activity coe5cients cannot be an index to the true concentration of the dissolved species and to nothing else since this value (which is real in contradistinction to the temperature of a single molecule), is constant, or zero, in the various thermodynamic environments.* 2


1637

+

Heat of Reaction and Catalytic Mechanisms There is another point of confusion in the use of thermodynamics in reaction velocity work, namely, one finds statements that certain reactions are probable or impossible on the basis of the value of the corresponding heat of reaction. The question is, of the two reactions, (I) AR = AR1 Rzand Ra, (2) AR = AR, can we predict from a knowledge of thermodynamic quantities which will take place in preference to the other? That ( I ) will take place in preference to (2) simply means that under the given conditions, the rate of ( I ) is large compared to that of (2). The very existence of endothermic reactions is proof enough that the value of the heat of reaction is not an index to reaction velocity. The very existence, for instance, of hydrogen-oxygen gas mixtures a t room temperature is proof enough that the free energy of reaction is not the determining factor in reaction rate.

+

+

9

J. Am. Chem. SOC., 49, 2956 (1927). Compare Dawson and Dean: LOG.cit.


1638

ROBERT E. BURK

It can easily be shown that in accordance with the theory of reaction velocity as given in this report, the heat of reaction is equal to the difference in the heats of activation of the direct and reverse reactions. Therefore, the only condition under which the heat of reaction has any significance for reaction velocity is when the heat of activation of the direct or reverse reaction is known. Thus, if it can be shown that the heat of activation in one direction is zero (as may be possible, e.g., in the combination of some atoms), then the heat of reaction can be substituted directly for the heat of activation in the velocity expression for the reverse reaction. The heat of reaction must, however, be at least as small as the largest heat of activation which can be allowed in the case of a given endothermic reaction proceeding at a given rate a t a given temperature. Therefore, without special knowledge, the value of the heat of reaction is no argument that a given mechanism is probable. Nor can such an argument be used against any mechanism. Thus, originally advanced Armstrong’ says the mechanism CO OHz = C 0 2 Hz, by Dixon to explain the rBle of water in accelerating the combustion of carbon monoxide, “is precluded by the fact that the heat of combustion of hydrogen to liquid water is greater than that of carbonic oxide.” Kapannaz has found that in the conversion of NH,CNS to S C ( S H Z ) ~ , the direct and reverse reactions are of the first order, and are uncatalyzed by glass wool and platinum strips. He found the difference in heats of activation to be equal to the heat of reaction.

+

+

Simultaneous Reactions These reactions are of great theoretical and practical interest. It is often found that, of two possible modes of decomposition of a molecule, some catalysts accelerate one almost exclusively, different ones accelerate the other almost exclusively, and still others accelerate bothS3 Clark and Topley have decomposed formic acid on several new catalysts. The various researches with this substance have not enabled the investigators to correlate the velocity of its modes of decomposition on different catalysts with the respective heats of activation. Differences in adsorption relations can, however, account for the discrepanciesP Senderen$ has found that the decomposition of formic acid in the wet way by sulphuric acid, potassium sulphate, and ortho-phosphoric acid is purely catalytic; that is, the action is not merely the taking up of the water formed. The sole products are carbon monoxide and water. I n the dry way, Senderens says alumina gives carbon monoxide and water at z0oO-z soo,but thoria gives carbon dioxide and formaldehyde. Particular attention should ~~

xature, 120,659 (1927). * Quar. J. Indian Chem. Soc., 4,217 (1927). See the reference in the introduction to this report. J. Phys. Chem., 32, 121 (1928). Hinshelwood: “Kinetics of Chemical Change in Gaseous Systems,” 181 (1926). 6 Compt. rend., 184,856 (1927).

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163~

be called to the fact that both of the latter catalysts are compounds of the same type, namely, difficultly reducible oxides. According to Tropsch' the decomposition of anhydrous formic acid to carbon dioxide and hydrogen, a t 3 5 5 O , occurred most readily over iron, and tinned iron. At 2 5 5 ' with tinned iron, formic acid decomposed to carbon monoxide plus hydrogen and small quantities of methane. Glass wool, asbestos, and pumice, a t 3 5 5 O , formed both carbon monoxide and dioxide, as did calcium carbonate a t 400°, and lithium carbonate a t z55'-410O, and T h o z a t 355'. Alumina a t 305' and 43 j oformed almost exclusively carbon monoxide and water, whether the formic acid was dry, or contained watervapor. I n particular, the results with tinned iron should be noticed. According to hliiller and Hentschel,*the decomposition of formic acid by ultra-violet light in the presence of highly dispersed platinum falls off greatly when the latter becomes aggregated. I n boiling aqueous solutions, the reaction yields carbon dioxide and hydrogen, and also water and carbon monoxide. According to Tropsch and R ~ e h l e n ,the ~ catalytic decomposition of formaldehyde was very rapid a t 3o0°-5000 over sodium carbonate, calcium carbonate, barium carbonate, zinc oxide, aluminz, thoria, chromium oxide, uranium oxide, alkalyzed iron, and lead, but not over quartz or antimony. Lead was the best catalyst. Considerable methyl alcohol was formed over thoria, alumina, uranium oxide, sodium carbonate and activated carbon. Formation of methane occurred especially with chromium oxide and with activated carbon, while the formation of unsaturated hydrocarbons occurred with alumina. A five percent yield of formic acid was obtained with uranium oxide. Zinc oxide gave few reaction products of high molecular weight; it decomposed dry formaldehyde smoothly to carbon monoxide and hydrogen, and moist formaldehyde smoothly to carbon dioxide and (hydrogen?). The latter is a very striking result. Tinned iron, whose properties in the decomposition of formic acid were so remarkable, had but slight effect on the decomposition of methyl alcohol.' The decomposition products with this catalyst, and with iron, tin, and aluminum a t 520' contained considerable quantities of carbon dioxide, carbon monoxide, hydrogen, and methane. According to Sabatier,5 thoria, alumina, and the blue oxide of tungsten, active in the catalytic decomposition of methanol, accelerate almost exclusively the reaction 2CH30H = CzH4 zH~O,

+

while 1lnO gives dehydrogenation, like the metals. Other oxides produce both effects simultaneously. The action of zinc oxide upon ethyl alcohol is Abhandl. Iienntnis Kohle, 7, I (1925); Chem. Abs., 21, 3530 (1927). Ber., 59, 1854 (1926). 3Abhandl. Kenntnis Iiohle, 7, 15 (1925); Chem. Abs., 21, 3530 (1927). 'Tropsch atid Schellenberg: Abhandl. Kenntnis Kshle, 7, 13 (IgZj); chem. abs., 21, 3530 (1927). Compt. rend., 185, 17 (1927).

ROBERT E. BURK

1640

principally dehydrogenation, while that of chromium oxide is principally dehydration. Both reactions are almost equal with zirconia and glucina. According to Kurtenacker and Werner,l in the decomposition of hydroxylamine, the ratio of the two reactions


+ + + +

3 N H 2 0 H = NH3 N2 3H20, and 4NH20H = 2”3 N 2 0 3H20, was controlled by the alkalinity. Platinum black favors (2), while platinum sponge, or platinum gauze does not affect the proportions of the end products. Mutiple adsorption seems particularly apt in explaining the effect of catalysts in selectively catalyzing one of two or more possible reactions of a compound, since on this mechanism a particular bond can be separately stretched. It is not so obvious that the atomic distortion theory can account for the results, especially when very specific and different effects are produced by compounds ionically of the same kind. (I)

(2)

The Thermal Decomposition of Ammonia Certain reactions, for one reason or another, have received a particular amount of attention. One of these is the thermal decomposition of ammonia. Reasons have already been given that the lowering of the heat of activation is the only effect of tungsten in catalyzing this reaction, and that the real catalyst in this case is a film of nitrogen upon the tungsten. The results of Kunsman have been described under “promoter action,” as have some unpublished results of the writer. The writer2 has found that the thermal decomposition of ammonia upon molybdenum is somewhat retarded by nitrogen, but not by hydrogen. It developed in this investigation that the cases of a poison blocking certain active centers completely, and leaving others free, and of a poison covering the entire catalyst, but itself having some catalytic activity, lead to equations of the same form expressing the reaction velocity. Schwab3 has decomposed ammonia upon heated strips of platinum and tungsten, and compared the behavior with that found by Hinshelwood and Burk.4 Schwab’s work was a t pressures of the order of 0.01 mm. while the latter was at pressures some 10,000times as great. At low pressures, the decomposition on both platinum and tungsten was found to be monomolecular; a t the higher pressure, the decomposition on tungsten is zero order. Nitrogen was found to retard both reactions a t the low pressure; a t the higher pressure i t retards neither the reaction on platinum, nor that on tungsten. Hydrogen was found to retard both reactions a t the low pressures, but a t the higher pressure, retards only the reaction on platinum. The temperature coefficients did not agree in the two researches. This is probably to be explained by the different poisoning phenomena. 1

Z. anorg. allgem. Chern., 160, 333 (1927).

* Proc. Nat. Acad. Sci., 13, 67 (1927). a 4

Z.physik. Chem., 128, 161 (1927). J. Chem. SOC., 127, 1105 (1925).



1641

These results are not surprising when one takes into consideration the knowledge of catalysts, of variable surface films, etc. Thus the nitrogen film on tungsten may be completely formed at IOO mms. but not a t 0.01 mms. The film a t the low pressures may be of a different structure, and have different properties from that a t the higher pressures. The gases may poison certain centers at the low pressures in proportion to their pressure. At the higher pressures, these centers may be blocked out almost from the start, and poisoning upon others may not yet have set in. Schwab considers similar explanations, but also considers multimolecular layers possible. He appears to consider the reaction to take place by collision from the gas phase. This could not be the mechanism, or the reaction on tungsten could not become zero order a t the higher pressure, as it does.

The Decomposition of Hydrogen Peroxide Kiss and Lederei-' h d that metals of invariable valence-calcium, cadmium, magnesium, strontium, and zinc-were without action upon the decomposition of hydrogen peroxide. However, the action of cobalt, manganese, and nickel, also, was so small that it was attributed to impurities (copper or iron). The action of copper and iron was very marked. They suggest a mechanism involving the rather strange assumption that molecular addition is the slow reaction. Their mechanism involves change in valence for both the copper and iron. The change in valence of the copper was thought to be shown by the color change of the solution. According to Spitalsky and Funck? the behavior of sodium molybdate as a catalyst for the decomposition of hydrogen peroxide depends upon whether it is in a stable state or not. This compound is thought to go over to the catalytically stable state by a preliminary reaction with hydrogen peroxide, and is rendered unstable again after a definite time, which is conditioned by the slow formation, and very slow decomposition, of an intermediaG compound. According to Pisarzheviskii and Roiter? the decomposition of hydrogen peroxide in solution is a unimolecular process. They consider the nature of the mechanism to be the same whether platinum, or manganese dioxide is used as a catalyst. Spitalsky and Koboseff4 think it necessary to assume more than one intermediate compound in order to explain all the known results for the catalytic decomposition of hydrogen peroxide by chromic acid. Elder and Ridea15 decomposed pure hydrogen peroxide vapor on quartz, platinum, and mercury. On quartz a t 85') the reaction was found to be zero order, inhibited by oxygen. This reaction stopped 80% short of completion. On platinum wire, the apparent order of the decomposition was the first, 'Rec. Trav. chim., 46, 453 (1927). * Z.physik. Chem., 126, I (1927). * Sei. Mag.Chem. Cath. Katerinoslav, 93 (1926);Chem. Aba., 21, 2414 (1927). Z.phyaik. Chem., 127, 129 (1927). 6 Trans. Faraday Soc., 23, 545 (1927).

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ROBERT E. BURK

which they consider to be determined probably by the rate 9f diffusion through an adsorbed or dissolved layer of oxygen. The reaction on a mercury surface was thought to consist of a preliminary direct oxidation to HgzO, followed by a coupled oxidation to HgO. According to Robertson,’ the catalytic decomposition of hydrogen peroxide by K2Cr2O7takes place in the following steps: HZOZ= zKCr04 HzO (I) K2Cr20, 2 KCr04 H~OZ = K2Cr20, HzO 02. (2) HVOI is thought to act as a negative catalyst because the rapid reaction ( 2 ) is superseded by HVO4 HzOz = HV03 HzO 02, HV04 being produced by oxidation of HVOI by H 2 0 2 ;it is not clear why a rapid reaction should be superseded by a slow one. The catalysis of this reaction by K2Cr207is promoted by MnC12.2 Robertson considers the shape of the curves for the promoted reaction t o show the presence of two superimposed first order reactions, one related to the concentration of intermediate compound due to the promoter. (The nature of this intermediate compound could not be determined by absorption spectral measurements). The other monomolecular reaction he considers related to the velocity constant of the promoted reaction. For other results on hydrogen peroxide, see the discussions of “promoter action,’’ and “dust particles.”


+ +

+

+ +

+

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Water Synthesis Benton and Elgin3 have found that in the synthesis of water over gold, the rate of reaction varied as the square of the hydrogen pressure, as the first power of the oxygen pressure, and inversely as the water vapor pressure. Independent adsorption measurements showed hydrogen not to be adsorbed appreciably, while oxygen is strongly adsorbed. They say: “A mechanism based on the interaction of adsorbed oxygen with two hydrogens adsorbed on adjacent spaces on the catalyst surface will not account for the observed kinetics, since it may readily be shown that with the relative adsorptions found for the two gases, this assumption requires the velocity to be proportional to the square of the hydrogen pressure and to the inverse first power of the oxygen pressure.” One must be very careful in drawing such conclusions from independent measurements of adsorption? In a previous work, these authors6 considered the same reaction over silver to involve collisions of gaseous hydrogen with adsorbed oxygen on the fraction of the surface not covered by water. They could have settled this point definitely had they used a selectively heated catalyst, and varied the bulb temperature. In case the reaction does not take place by collision from Proc. Kat. Acad Sci , 13, 192 (1927). J. Am. Chem. SOC.,49, 1630 (1927). * J. Am. Chem. SOC.,49, 2426 (1927). ‘Compare Hinahelwood end Prichard: J. Chem. SOC., 127, 806 (1925). t, J. Am. Chem. SOC., 48, 3027 (1927).


SIXTH REPORT O F THE COMMITTEE ON CONTACT CATALYSIS

1643

the gas phase, this variation in bulb temperature would, ordinarily, be without great effect upon the reaction velocity. Remy’ has investigated this reaction extensively, using metals of group eight as catalysts. He assumes, with Hoffmann,* that both hydrogen and oxygen, in the water synthesis, are activated by the catalyst. He thought it might be possible to connect catalytic activities and the quantitatively measurable chemical affinities betweeen the catalyst and hydrogen and oxygen. Attempts to correlate catalytic activity and the ability of the metal to dissolve hydrogen did not materialize experimentally. He found, however, without exception, that if a metal has a greater affinity for oxygen than for hydrogen, that its catalytic activity was increased by pretreatment with hydrogen and vice-ver~a.~He gives a series of elements arranged in the order of hydrogen affinities and oxygen affinities. The oxygen series was based on the heat of formation of the simplest oxide while the hydrogen series was based on the amount of hydrogen adsorbed by the metal. With alloys containing one rare metal, and one iron group metal, the catalytic efficiency, after pretreatment, was that predictable from the foregoing principle, with the exception of ruthenium-iron. Palladium-platinum was also an exception to the rule. From the point of view of the more important question of the absolute effectiveness of the catalyst, Remy distinguished the further factor of the relative ease of stable oxide formation as compared to the combination with hydrogen, those catalysts favoring the reaction to stable oxide being relatively inefficient. He makes the wrong assumption that this is necessarily proportional to the heat of formation of the oxide.

Special Points According to Sabatier; nickel cannot be used for the hydrogenation of the oxides of carbon to alcohols because i t is too good a hydrogenation catalyst, taking the substances on through to methane. Bogandy and Polanyis find that hydrogen atoms are emitted when a jet of sodium vapor is directed against a film of sulphuric acid, kept a t liquid air temperature in a high vacuum. Only hydrogen molecules come off when a filmof ice is Eubstituted for the sulphuric acid. They think the explanation sufficient that the heat of reaction is negative in the latter case, and positive in the former one. According to ClarkI6radiation of the platinum catalyst, (used for the oxidation of sulphur dioxide) with X-rays, produced erratic changes in its activity. Belenski’ has found that for the cata!ytic decomposition of potassium chlorate, the poorest catalysts are those which are easily oxidized. Z. anorg. allgem. Chem., 157, 329 (1926). For experimental method see, Remy and Gonningen: Z. anorg. allgem. Chem., 149,283 (1925). * Ber., 49, 2369 (1916). a Compare the discussion of “films on catalysts.” Compt. rend., 185, 17 (1927). Naturaissenschaften, 14, 1205 (1926). Brit. J. Rahology, 23, 112 (1927). Chem. Abs., 21, 2415 (1927).

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ROBERT E. BURK

Jaboaynski and Rytel' have found a novel type of auto-catalysis. H2S203 decomposed monomolecularly and reversibly with the formation of monatomic sulphur. After the reaction had proceeded a certain distance, aggregates of sulphur started to form, stopping the reverse reaction, with a net increase in reaction velocity. According to Schlesinger and Malkina-Okun,2 the halogen acids shift the equilibrium, in the formation and saponification of ethyl acetate, in proportion to their concentration, and simultaneously catalyze the reaction. The shift in equilibrium is allowable, but is not connected with the catalytic action. Experiments of Richards and Loom$ seem to indicate that high-frequency sound waves catalyze the iodine "clock" reaction, and others. Their method of investigation was a little unorthodox. Thoren4 found a stepwise increase in the activity of nickel as the temperature was raised. The reactions investigated were: C2H4 = C2H6; C6H6 3Hz = CBHIZ; 0 2 Hz = HzO. The actual percentage increase in activity a t the various activation temperatures was not reproducible. Adsorbed substances can move around on the surface15as well as evaporate without having moved. Constable,6 accordingly, points out that collisions between more or less freely moving molecules in the film can take place. He also points out that this effect is probably of no great importance in catalysis. H. A. Taylor and Pickett' find, in decomposing hydrogen sulphide on platinum by a streaming method, that there is a region where the rate of decomposition is independent of the rate of flow. The temperature coefficient at this stage was 11,750 calories. They conclude that this represents the heat of evaporation of sulphur from the surface. This does not seem a necessary conclusion. Miscellaneous Results

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Pearce and Ott8 decomposed various esters upon pumice-supported nickel. Added hydrogen was found not to change the reaction products. Ottensooserg has found a zinc oxide catalyst effective for the oxidation of alcohols to aldehydes or ketones. Uchidallo has studied the catalytic decomposition of nitric oxide on platinum and on FezO3,and says that this reaction affects the yield in the oxidation of ammonia by these catalysts. Zelinski and Balandin" studied the dehydrogenation of decahydronaphthalene over various catalysts, and found the order of activity to be platinumasbestos, platinum-charcoal] palladium-asbestos, nickel-asbestos. 'Chem. Abs., 21, 2415 (1927). * Ber., 60, 1479 (1927). J. Am. Chem. SOC.,49, 3086 (1927). 4 Z . anorg. allgem. Chem., 163, 367 (1927); 165, 5 E.g. Becker: Phys. Rev., (2)28, 341 (1926). 6 Proc. Cambridge Phil. SOC.,23, 593 (1927). J. Phys. Chem., 31, 1 2 1 2 (1927). 8 J. Phys. Chem., 31, 102 (1927). 'Bull., 41 324 (1927). 'OJ. Sac. dhem. Ind., (Japan), 30, 171 (1927). "Z. physik. Chem., 126, 267 (1927). J

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SIXTH REPORT O F THE COMMITTEE O S CONTACT CATALYSIS

1645

hladenwald, Henke, and Brown‘ studied the rate of hydrogenation of nitrobenzene on lead catalysts. The formation of azobenzene is a side reaction, the extent of which depends upon, ( I ) the metal constituting the catalyst, ( 2 ) the activity of the particular catalyst, (3) the temperature. Lang2 finds silver to catalyze the change of manganic to manganous chloride, and the evolution of chlorine from dilute hydrochloric solution by cerium nitrate. Walkel3 has conduofed an extensive investigationof the reactions of ethylene over various catalysts. Various products of decomposition were identified. A polymerized oil was also found. He was unable to identify the active constituent of Jena glass by trying them out separately. He concluded that, “The activation of the ethylene by the walls of the reaction chamber is thought to be due to a combination of two or more of the oxide constituents, which come into play a t the elevated temperature.” Silica gel, silica gel and borax, silica gel and calcium hydroxide, calcium hydroxide. borax, calcium silicate, zinc oxide, ferric oxide, the oxides of lead, sodium, nickel, cobalt, iron, alkalized iron, chromium oxide-iron, and pumice-iron were investigated a t various temperatures. According to Palmer: the rate of reduction of copper oxide on clay by carbon monoxide is not governed by the rate of diffusion through the already reduced copper film, and the reduction is not hindered by the carbon dioxide formed. When hydrogen was the reducing agent, after a rapid initial reduction, the reaction came almost to a standstill. lt7ater apparently poisons the reaction. Only such results have been placed in this section as do not fit very well under the previous heads. No attempt has beeen made to exhaust the literature of the year (for instance, that of organic chemistry) with respect to miscellaneous results, desirable as that may have been for some purposes. The spirit of this report has been rather to systematize the results, and to try to enumerate and evaluate, as far as possible, the various factors which could be expected t o operate in catalysis, which, after all, cannot be a new trick of nature entirely unrelated to the rest. It is not universally realized that none of the work on catalysis has been, in a strict sense, quantitative. In no case has the number of reacting molecules upon the surface a t a given time been known. I n no case has the number of active catalyst atoms, or centers of a given type been accurately known. We have not been sure of the identity of the actual catalyst in many cases, nor of the reactants on the catalyst. Some of these unknowns are starting-points for investigations of homogeneous reaction velocity. It is no wonder, then, that catalytic results are so often ambiguous and difficult to interpret. January 25, 1928

J. Phys. Chem , 31, 862 (1927). Ber , 60, 1389 (1927). J. Phys. Chem , 31, 961 (1927). Trans. Am. Electrochem. SOC.,51, 4 (1927).

Sixth Report of the Committee on Contact Catalysis - The Journal of

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