Irreversible Reduction and Catalytic Hydrogenation - The Journal of

Publication Date: January 1923. ACS Legacy Archive. Cite this:J. Phys. Chem. 1924, 28, 10, 1096-1107. Note: In lieu of an abstract, this is the articl...
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IRREVERSIBLE REDUCTION AND CATALYTIC HYDROGENATION' BY J A M E S B. CONANT AND HAROLD E. CUTTER

Organic compounds which are rapidly reduced in solution by soluble reducing agents may be divided into two classes. In the one class are those substances of which the oxidized and reduced forms are in mobile equilibrium a t room temperature with reducing or oxidizing agents, and in the other are compounds whose behavior under ordinary conditions does not correspond to the existence of such equilibria. As examples of the first class may be mentioned the systems, quinone-hydroquinone, indigosulfonate-leuco-indigosulfonate, nitrosobenzene-phenylhydroxylamine in which the relationships between the oxidized and reduced substances are essentially those so common in inorganic chemistry and represented by the system ferric chloride-ferrous chloride. Numerous investigations2 of the last few years have shown that the oxidation-reduction potentials of such organic systems may be measured and these potentials condition the behavior of the compounds in question towards oxidizing or reducing agents. In almost all cases of such reversible oxidation or reduction, the speed at which equilibrium is attained is very rapid even in dilute solutions a t room temperature. The reduction of compounds of the second class such as azo dyes, nitro compounds, and certain unsaturated ketones also proceeds rapidly, but, unlike the reduction of quinone does not reach a final equilibrium. This is evident from the fact that it is impossible to find a reducing agent with which the reduction will proceed to such a measurable equilibrium (reducing agents either are without effect or cause complete reduction) and also because the reduction can not be reversed by treating the reduced compound with an oxidizing agent of potential higher than that of the weakest reducing agent necessary for the reduction. Thus while I , 4 naphthoquinone is reduced to the hydroquinone by stannous chloride (potential= +o . 4 ) and the hydroquinone is oxidized to the quinone by ferric chloride ( + o . 8 ) ) dihenxoylethane (C6H5COCH&H2COC6HS) is not affected by ferric chloride although the oxidized form, dibenzoylethylene, (C,H,COCH = CHCOCsHb) is not reduced by stannous chloride but only by the more powerful titanous chloride ( + o . 0). We shall designate the reduction of a substance of this second class by the term irreversible, in order to distinguish it from the reversible reduction of a compound of the first class. Contribution from the Chemical Laboratory of Harvard University. Granger and Nelson: J. Am. Chem. SOC.43, 1401 (1921); W. M. Clark: J. Wash. Acad. Sci. 10, 255 ( I 20); Public Health Reports 38, 443, 666, 933, 1669 (1923); Biilmann: Ann. chim. 16, 321 q1921); 19 137 (1923); La Mer and Baker: J. Am. Chem. SOC.44, 1954 (1922); Conant, Kahn, Fieser and Lute: Ibid, 44, 1382 (1922); Conant and Fieser:Ibid, 44, 2480 (1922); 45, 2194 (1923). 2

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No oxidation-reduction potential can be measured in the case of a compound which is irreversibly reduced since the oxidized and reduced forms are not in equilibrium with the reducing agents. Thus, while an inert electrode immersed in mixtures of a quinone and the corresponding hydroquinone rapidly attains a definite potential which is a function of the concentrations of the two compounds and the hydrogen ion, no reproduceable potentials can be obtained with mixtures of dibenzoyl-ethylene and dibenzoylethane. It would seem at first sight, therefore, as though it would be impossible to describe quantitatively the conditions limiting the reduction of all those organic compounds which fall in the second class noted above. However, while no real oxidation-reduction potential of such irreversible systems can be formulated or measured, it is possible to characterize them with considerable precision by determining the potential of the weakest reducing agent (Le. the reducing agent of highest oxidation-reduction potential) which will cause reduction. Thus if a series of reducing agents of known potential are available, it is a relatively simple matter to determine that dibenzoylethylene, for example, is not reduced by reagent A of potential +o . 3 I but is reduced by B of potential + o . 24; some value within this range of + o . 3 1 to + 0 . 2 4 can be considered as the potential of the critical reducing agent which would just cause appreciable reduction ; the potential of such a hypothetical reducing agent we have called the “Apparent Reduction Potential” of the substance in question; in the case of dibenzoylethylene, the value of the apparent reduction potential is + o . 2 7 ko . 0 3 volts. In previous papers* from this Laboratory we have determined the apparent reduction potential of a variety of organic substances supplementing the usual inorganic reducing agents with certain hydroquinones and similar substances whose potential could be readily measured. The accuracy of the method obviously depends primarily on having a series of reducing agents such that the potential of each member is only slightly less than that of the preceding member; with the reducing agents now available this can be satisfactorily realized within certain ranges and the apparent reduction potential determined within o 03 of a volt. For compounds which are not reduced by reducing agents of potential $ 0 150, however, there are relatively few reagents available and the apparent reduction potential is correspondingly uncertain. The actual experimental procedure is much simplified by using an electrochemical device for determining whether or not a certain reducing agent reacts with the substance in question. This has been done by introducing the substance into an equimolecular mixture of the reducing agent and its oxidized form (e.g. a quinone and the hydroquinone) and noting whether or not the potential of an inert electrode immersed in the mixture changes during a definite time (30 minutes, for example). A slight reaction between the reducing agent and the substance introduced will obviously cause a considerable change in the ratio of the concentrations of the reducing agent and its oxidized form and a corresponding change in potential. With a great Conant and Lutz: J. Am. Chem. SOC.45, 1047 (1923); 46, 1295 (1924)

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majority of the substances investigated there was a sharp distinction between the reagent which was without action and the next in the series which caused reduction, since with the latter the reduction process proceeded a t such a rate that it was practically complete in I O minutes or less. In a few instances where it was possible to use a closely graded series of reducing agents, “border line cases” were met which very nearly corresponded to the “critical reducing agent” whose potential would represent the actuaI apparent reduction potential of the substance under investigation. In such border line cases the reduction proceeded a t a relatively slow rate, while a reducing agent of slightly higher potential was without effect, and a reagent of slightly lower potential caused rapid reduction. Before discussing the probable mechanism of irreversible reduction and the relation of this problem to catalytic hydrogenation, it will be well to consider certain obvious limitations to the quantitative formulation just outlined. Since we are dealing with a reaction which runs to completion, the problem is one of rates of reaction and not of equilibria. The question may, therefore, well be raised as to how much significance can ke attached to these apparent reduction potentials. Are these irreversible reductions governed primarily by the potential of the reducing agent? Does it necessarily follow that all reducing agents of pctential less than a certain value will reduce the substance in question? May not the whole process be determined by specific properties of the reducing agents and not by their potential? The answer to these questions is a matter of experiment and while the final verdict can not be rendered until all the conceivable cases have been examined, the evidence a t present seems very clear. If one specifies the solvent, (in particular the hydrogen ion concentration), and the temperature and defines the term “appreciable reduction” as a few per cent reduction in 30 minutes, the results of experiments in this Laboratory with some 7 5 conipounds of a half dozen types can be summarized in the statement, under definite experimental conditions the potential qf the reducing agent employed determines whether or not appreciable reduction wzll take place in homogeneous solution. We have found no exception to this although the reducing agents employed have been as different as titanous chloride, leuco-indigosulfonate and sulfonated anthrahydroquinone. Fxperiments have been carried out in both aqueous solutions and solutions containing large amounts of acetone or in some cases of alcchol, the use of these organic solvents being necessary with substances insoluble in water. The few compounds which could be investigated in all three solvents (all about 0 . I N in HC1) were found to have essentially the same apparent reduction potentials in each solution. The effect of changes in hydrogen ion concentration (investigated in aqueous buffer solutions) need not be considered in detail in this paper but the results can be summarized by the statement that the concentration of the hydrogen ion affects the process of irreversible reduction in much the same way as it affects the reversible reduction of quinones and similar compounds.

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The following table (taken from a previous paper) is of interest as showing the influence of structural changes on the apparent reduction potent,ial; the 0 0 II

II

series of conipounds given below all contain the system -C -C = C-C which is also present in many substances which are reversibly reduced, such as the quinones and indigo. These results were all obtained a t 24'*z0, and are, therefore, comparable with each other but not with the action of a reducing agent a t some higher or lower temperature. For example, undoubtedly a t 100' a substance might be rapidly reduced by a reducing agent of potential higher than the apparent reduction potential determined a t 25'. The exact magnitude of this effect of increased temperature has not yet been ascertained.

TABLE I Apparent Reduction Potentials Substance

Solvent

App. Red. P o t . (normal hydrogen electrode = G )

Dibenzoylethylene (cis and trans) A +o 2 7 (*.03) (C~H~COCH=CHCOCEH~) Benzoylacrylic acid (CeH5CCCH=CHCOOH) A + o . c ~( * . ~ 6 ) Benzoylacrylic acid B $0.08 ( * . 0 6 ) Benzoylacrylic ester(C6H5COCH= CHCOOCZH,) A + o . c ~( * . ~ 6 ) Maleic ester (C,H,OOCCH = CHCOOCZH,) A - . 2 5 (*.IO) Maleic acid A - . 2 5 (*.IO) Maleic acid B -.25 (*.IO) A = 75y0acetone, 2 5 % aqueous hydrochloric acid; total acidity o . 2 N . B =aqueous 0 .zN HCl.

Low Temperature Catalytic Hydrogenation It will be noted from Table I that maleic acid has an apparent reduction potential of - 0.z 5 * 0 . I O ; it is instantaneously reduced by chromous chloride ( - 0.40) but is not affected by vanadous chloride ( - 0 . I 5) even after many hours; it is, however, reduced in dilute hydrochloric acid by hydrogen and colloidal platinum. The potential of the hydrogen-platinum combination is, obviously, that of the hydrogen electrode and in o.zN hydrochloric acid is about - .05, a value considerably above the potential of vanadous chloride which is without effect. How is this to be explained? A similar problem was raised in an earlier paper' in connection with the reduction of dimethylacrylic acid. This compound is not reduced by chromous chloride but is readily hydrogenated by hydrogen and a catalyst in 0 . zN hydrochloric acid at room temperature. We concluded from these previous experiments with dimethylacrylic acid that the process of catalytic hydrogenation, unlike reductions brought about by soluble reducing agents, could not be formulated electrochemically and that J. Am. Chem. SOC.,44, 265

(1922)

JAMES B. CONANT AND HAROLD B. CUTTER

I IO0

probably a different mechanism was involved. We suggested that in the case of a reduction process caused by chromous chloride the addition of electrons and hydrogen ions was involved and in the case of catalytic hydrogenation the direct addition of activated hydrogen atoms. Bancroftl in the Second Report of the Committee on Contact Catalysis has reviewed our earlier paper and has stated that the experiments as they stand are not convincing. He has suggested that it is important to determine whether or not colloidal platinum mill catalyze a process of reduction involving a soluble reducing agent. The results of such an experiment with chromous chloride would, however, not be very significant since, as lie points out, the catalyst will cause the evolution of hydrogen and reduction must take place unless some catalytic poison is present, as the system gaseous hydrogen catalyst dimethyl acrylic acid is known to give isovalerianic acid. A crucial experiment can be carried out, however, with a reducing agent which is not decomposed by colloidal platinum. An equimolecular mixture of anthrahydroquinone 2 , 7 disulfonate and anthraquinone z , 7 disulfonnte has a potential in 0 . I N hydrochloric acid of + o . 1 7 ; a platinum catalyst is without action on such a mixture. If now maleic acid is introduced into the mixture together with the platinum catalyst it is found that no reduction takes place. That the catalyst is not poisoned can be shown by carrying out the usual process of catalytic hydrogenation with the same solution at the conclusion of the experiment. Similar results were obtained with dimethyl acrylic acid. The experimental details are given below. In a glass cell of 350 cc. capacity equipped with a stirrer, electrodes, gas inlet and outlet tubes, and connected to a saturated calomel electrode with a saturated KC1 bridge, there were placed I O O cc. of vanadous chloride solution, I O O cc. of 0 . I N hydrochloric acid, and I O cc. of a solution of colloidal platinum. The apparatus was swept out with nitrogen and a solution of anthraquinone 2 , 7 disulfonate in O . I N hydrochloric acid was added in such amounts ( 2 0 cc. of a 0.05 1 4 solution) that an approximately equimolecular mixture of the anthraquinone and anthrahydroquinone sulfonates resulted; the amount of vanadous chloride being sufficient to reduce only half the added anthraquinone sulfonate. The potential of the bright and platinized electrodes immersed in this mixture were constant for half an hour a t +o. 184 (on the hydrogen scale); 2 0 cc. of a o . r M solution of dimethyl acrylic acid in 0.1i3'lhydrochloric acid was now introduced and the potential of the cell noted over a period of half an hour. There was a change of 2 millivolts in the direction opposite to that which would have been caused by any reduction. If any reduction had occurred there would have been a change of a t least 20 millivolts as den?onstrated by many experiments in this Laboratory in connection with the determination of apparent reduction potentials. T o test whether the catalyst was still active, a quarter of the mixture was shaken with hydrogen in the usual manner after the completion of the above experiment. In the course of

+

J. Phys. Chem. 27, 801 (1923).

+

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half an hour I O cc. of hydrogen was absorbed and a decided odor of isovalerianic acid resulted. T o another portion 0.25 g. of dimethyl acrylic acid was added and the mixture shaken with hydrogen; after an hour 40 cc. of hydrogen had been absorbed (theor. am’t 58 cc.) and a strong odor of isovalerianic acid was apparent. The catalyst used in this experiment was prepared by the reduction of 3 cc. of a 10% solution of chlorplatinic acid mixed with 7 cc. of 1% gum arabic solution by shaking with hydrogen with one drop of ‘(seeding solution” prepared in the usual manner.’ The experiments with maleic acid were performed exactly as above except that in place of the colloidal platinum solution, a suspension of o I g. of a platinum catalyst prepared according to Adams’ directions? was employed, the platinum oxide being reduced to the active state by shaking with hydrogen before the catalyst was used. After the electrode potential had been constant for 1 5 minutes a t +o 130, 2 0 cc. of a o I M maleic acid solution was added. A change of only one millivolt was noted during the next 30 minutes. TO test the catalyst, 0.5 g. of maleic acid was added to 50 cc. of the solution and the mixture shaken with hydrogen. After 90 minutes, 160 cc. of hydrogen had been absorbed (theor. = 1 2 0 cc.), the solution was filtered and extracted three times with ether, and the ether layer on evaporation yielded 0.35 g. of succinic acid. If the process of catalytic hydrogenation is similar to reduction by soluble reducing agents then the catalyst must function not only to make gaseous hydrogen an effective electrochemical reducing agent a t room temperature, but also to alter in some way the maleic acid or dimethyl acrylic acid so that it is reduced a t a much higher potential than otherwise is the case. In other words, if an electrochemical account of catalytic hydrogenation is to be formulated, we must assume that the apparent reduction potential of maleic acid is much higher (above - 0 05) in the presence of a platinum catalyst than in its absence. That some catalyst having such an effect might be found is entirely to be expected; in fact it is readily conceivable that in the presence of the proper catalyst the syatem maleic acid-succinic acid might be as mobile as the quinone-hydroquinone system. If this is the effect of the platinum catalyst, it should function equally well with other reducing agents (providing it is not “poisoned” which can be determined experimentally). However, the experiments given above show conclusively that platinum does not cause thereduction of either maleic or dimethyl acrylic acid by a reducing agent having a potential only slightly above that of hydrogen itself. It might be objected that the difference in potential between the reducing agent chosen (+0.17) and hydrogen ( - 0 o s ) was sufficient to vitiate any conclusions; it is possible that the energy relationships in the system maleic acid, hydrogen and succinic acid correspond to a true oxidation-reduction potential between these values, for example +o os. If this were the case, however, and an electrochemical explanation of catalytic hydrogenation Skita: Ber. 45, 3589 (1912) Adams and Vorhees: J . Am Chem. SOC. 44, 1397

(1922)

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J A M E S B. CONANT AND HAROLD B. CUTTER

correct, an oxidizing agent such as ferric chloride (+o 8) should oxidize succinic acid to maleic acid in the presence of platinum and isovalerianic acid to dimethyl acrylic acid. Such reactions do not proceed as the following experiments show. The apparatus used was identical with that described above. A o 01 solution of ferric chloride in 0 . I M hydrochloric acid ( 2 5 0 cc.) was introduced into the cell together with 7 5 cc. of a colloidal platinum solution prepared in the usual manner. After the potential had been constant for 30 minutes a t i - 0 . 8 3 8 , 7 5 cc. of a o I Msuccinic acid solution was added and a potential change of only 3 millivolts occurred during the half hour. In another experiment with Adams’ catalyst the same results were obtained. Siniilarly when a solution of isovalerianic acid was introduced in another experiment, the potential was constant within a millivolt for 30 minutes. To illustrate the effect on the potential of any action between the ferric chloride and the substance introduced, the following experiment may be mentioned. To 2 5 0 cc. of O.IMferric chloride and catalyst as above, 7.5 cc. of O.IMsolution of hydroquinone was added. The potential fell from +0.833 to $0.666 within ten seconds after the addition of the hydroquinone. That the ferric chloride had not poisoned the catalyst in the above experiment was demonstrated in each case by taking 30 cc. of the solution after the experiment, adding 0.5 g. of dimethylacrylic acid and shaking with hydrogen; in each case 1 2 0 to 160 cc. of hydrogen was absorbed in 90 minutes. (Theor. = 1 2 0 cc.) In the experiment with Adams’ catalyst, ferric chloride and succinic acid, the efficiency of the catalyst was proved by the addition of 0.5 g. of maleic acid; after 60 minutes, 1 2 5 cc. of hydrogen was absorbed and 0 . 2 5 g. of succinic acid was obtained by extraction with ether. To recapitulate, if the process of catalytic hydrogenation (in solution a t room temperature) is to be interpreted in the same terms as the process of reduction by soluble reducing agents, it must be assumed that the catalyst niarkedly changes the apparent reduction potential of the system. Experiments with a reducing agent almost as powerful as hydrogen show that this is not the case; the possibility of poisoning of the catalyst being excluded by experiment. The fact that the catalyst does not promote the oxidation of the reduced compound by an oxidizing agent of potential 0.8 volt above hydrogen clearly shows that the platinum is not a catalyst which has converted an irreversible system into a mobile reversible one. Our previous statement that “the process of catalytic hydrogenation can not be successfully formulated in terms of oxidation-reduction potentials” seems to represent accurately the facts and to place the process of catalytic hydrogenation in strong contrast to processes of irreversible reduction brought about by soluble reducing agents. It is hardly necessary to add the qualification that we are discussing the catalytic hydrogenation of substances whose reduction plcducts are I ot in mobile equilibrium with the oxidized form and a reducirq agent. The action of hydrogen and platinum on substances of the first class enumerated at the beginning of this article, such as quinone, is, of course, a strictly reversible

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process and will proceed to an equilibrium point depending on the reduction potential of the organic system and the potential of the hydrogen electrode. The final position of equilibrium in such a mobile system must be in accord with an electrochetnical formulation irrespective of whether the hydrogenplatinum combination acts in the same way as titanous chloride, for example, or in some entirely different manner. The mechanism of such relatively rare examples of catalytic hydrogenation can not be deduced, therefore, from electrochemical considerations. The Mechanism of Irreversible Reduction and Catalytic . Hydrogenation The sharp contrast between the action of hydrogen and a platinum catalyst on maleic acid and the action of vanadous and chromous salts on the same substance points to a totally different mechanism for the two ways of irreversibly adding hydrogen atoms to the double linkage. Let us first consider, in some detail, the irreversible reduction of maleic acid and see if any reasonable mechanism can be suggested which would account for the observed facts. I n the following discussion we shall make a number of more or less probable assumptions in regard to the numerical value of certain equilibrium constants and rates of reaction in order to outline a more definite picture than could be formulated in merely qualitative terms; we do not, of course, pretend that such an arbitrary procedure leads to any conclusions of quantitative significance but hope that it may suggest at least a possible explanation of the difference between reversible and irreversible systems. Lacking information in regard to the free energy change involved in the formation of succinic acid from maleic acid and hydrogen, we shall use the value 38 Cal. which is the total energy change from calorimetric measurements; this assumption that the total and free energy change are equal is probably not greatly in error. It may be supposed that the action oE a soluble reducing agent on maleic acid is similar to the reduction of quinone and that the first step is the addition of two hydrogen atoms to the oxygen atoms a t the ends of the doubly conjugated system. This di-enol form of succinic acid (B in the equation below) would be very unstable and would rearrange to give succinic acid, and the mechanism of the reduction would be as shown below:

//o /OH CH-C -OH + ~ H + + ~ E ~ C H-OH = C

l

//o

CHzC - OH

//o

CH-C-OH CH=C-OH CH,C -OH (A) (1) (B) (2) (C) fast ----t relatively slow + If it be assumed that the addition or removal of hydrogen atoms to com-

0

0

ii

pounds containing the system -C-C

II =

C-C - by means of soluble reducing

'

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JAMES B. CONANT AND HAROLD B. CUTTER

or oxidizing agents proceeds only through some intermediate compound such as the di-enol form (or its ions), the difference between quinone and maleic acid is readily explained. This differencc! is to be found in the fact that with hydroquinone the di-enol form is stable while wit8hsuccinic acid it is highly unstable. I n other words, the free energy of the first step of the scheme outlined above (AJB) is prac,tically equal to the free energy of the entire process in the case of quinone, while with maleic acid it is very much less than that of the total free energy change and may be of the opposite sign. For example, if the enolization constant' of a single enol group were IO-^^, the free energy of the change of the di-enol form of maleic acid to succinic acid ( z enol = 60 Cal. a t 2 5 ' ; in this case the free groups involved) would be I .364 Xlog energy change of the system maleic acid, hydrogen, di-enol form of succinic acid, would be 38-60= - 2 2 Cal. which corresponds t'o a potential (on the hvdrogen scale) of -0.5 volt. Under such circumstances the action of a reducing agent of potential -0.30 in a solution normal with respect to hydrogen ions would cause the formation of only IO-' moles of di-enol from I mol of maleic acid since the difference in potential between the reducing agent and the potential of the system ( -0.5) is equal to .ozg5 log [di-enol] when the amount of di-enol formed is very small. The action of a more powerful reducing agent of potential - 0 . 3 8 would f o r m IO-* mols. This equilibrium condition first attained would be continually upset by the rearrangement of the di-enol form into succinic acid, the concentration of the di-enol, however, would remain practically constant, more being very rapidly formed by the action of the reducing agent. The rate of this rearrangement' and the amount of di-enol formed by the reducing agent in question would determine whether or not "appreciable reduction" occurred. Let us see how rapidly such ketonization would have to proceed in order to have 1% of the maleic acid reduced in 16 to 1 7 minutes (103 seconds) by a reducing agent of potential -0.30 volts on the one hand and a reducing agent of -0.38 volts potential on the other. If we start with one liter of molar maleic acid solution, this reduction corresponds to the formation of IO-^ moles in 1000seconds or 10-j moles per second; the di-enol form must rearrange at, t,his rate which will be practically uniform since this is a rather unusual instance of a monomolecular reaction fed by a much more rapid reversible reaction. With t'he reducing agent of potential - 0 . 3 0 , the amount of di-enol in equilibrium is only IO-'" moles and in order for 10-j moles of this to rearrange each second it is obvious that t,he rate of this monomolecular isomerization must he such that one-half of the equilibrium amount (0.5 x IO-^) will rearrange in 5 X IO-^ seconds. On the other hand, with a reagent of potential -0.38 since there are IO-^ moles in equilibrium the rate of rearrangement would only have to be such that half the material ketonized in 5 seconds. The actual rate of ketonization of the The value IO-^^ for the enolization has been assumed quite arbitrarily in order to demonstrate the influence of this factor. Some such very small value probably is somewhere near the truth since the 01 hydrogen atom in acids is much less active than in ketones where the value is too small to be measured.

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enol form of acetoacetic ester as measured by Meyerl in aqueous solution corresponds to half isomerization in 6 seconds. However, as this rate would undoubtedly be faster in the presence of acids, not too much significance can be attached t o a comparison between it and the rates just calculated, but it indicates the order of magnitude of such speeds of ketonization as have been observed. Since the ratio of the speeds of the enolization and ketonization are equal to the equilibrium constant of the change from succinic acid t o the di-enol form which we have assumed to be very small, one can postulate either a very high speed of ketonization or a very low speed of enolization. We are inclined to believe that the magnitude of the rate of ketonization is not very different from that observed with such substances as acetoacetic ester and that the rate of enolization is correspondingly small. This would account for the fact that oxidizing agents which readily oxidize hydroquinones are without appreciable effect on succinic acid, for if these oxidizing agents are only effective in removing hydrogen atoms from enolic and phenolic substances the process could proceed no faster than this very slow enolization. If kl/kz= IO-^* and kz corresponded to the ketonization rat,es given above, it is evident that any process that must proceed through the enol form would not be appreciable in a life time. Thus, if one merely assumes that soluble oxidizing and reducing agents of the type of ferric chloride and chromous chloride can not directly remove or add hydrogen atoms to the carbon linkage but only to the oxygen atoms a t the ends of the doubly conjugated system, one arrives a t a satisfactory explanation of the reversible reduction of quinones on the one hand and the irreversible reduction of maleic acid on the other. The fact that no case appears to be on record of the reduction of an ethylene hydrocarbon hy a soluble reducing agent is in harmony with the assumption underlying the mechanism just proposed. It is true that acetylene is reduced to ethylene by chromous chloride but this may be due to peculiarities inherent in the triple linkage. Ethylene hydrocarbons are readily hydrogenated catalytically and there is every reason to believe that the hydrogenation of the double linkage in maleic acid proceeds in a similar manner. From our point of view, the catalyst activates the gaseous hydrogen and adds hydrogen atoms directly to the ethylene linkage in contrast to chromous chloride which adds hydrogen only to the oxygen atoms at the end of the conjugated system. From the difference in the two mechanisms arise the differences between irreversible reduction and catalytic hydrogenation which we have considered in this and our earlier paper. If one now seeks for the reason why soluble reducing agents and oxidizing agents are only effective in their action on organic substances containing certain linkages, the answer must be found in the mechanism underlying the oxidation and reduction of these reagents themselves. While the electrochemical equations for reversible oxidation and reduction can be derived ~~

~

Ann. 380, 235 (1911).

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J A M E S B. CONANT AND HAROLD B. CUTTER

either from the concept of hydrogen in equilibrium with the reagents or in terms of electron transfer, the most satisfactory picture of these reversible changes seem to be arrived a t from the latter point of view. Thc tremendously small hydrogen pressures which must be post,ulated and the fact that satisfactory oxidation-reduction potentials can be measured with electrodes‘ which will not function as hydrogen electrodes, are arguments against the concept of an oxidation-reduction electrode as a low pressure hydrogen electrode. The fact that satisfactory potentials of mixtures of ferrous and ferric salts can be measured in anhydrous pyridinZ seems convincing proof that a direct interchange of electrons between the ions is the essential feature of many cases of oxidation and reduction in inorganic chemistry. If the action of chromous chloride3 on ferric chloride, for example, is thus merely the transfer of an electron from the chromous to the ferric ion, there is every reason to believe that a similar mechanism takes place when chromous chloride reduces an organic compound. In the cases which we are considering, however,the resulting negatively-charged organic molecule can not exist by itself in aqueous acid solution but combines with hydrogen ions and becomes an electrically neutral molecule. It is a n interesting problem whether this addition of two electrons and two hydrogen ions takes place in the order just mentioned or in the reverse order or simultaneously; no evidence appears now available which would enable one to distinguish between these three possibilities. The change may be represented as: A 2~ 2HS+AH2

+ +

From our point of view, the reduction of an organic compound in homogeneous solution, whether reversible or irreversible, is primarily a process of electron transfer usually accompanied by the capture of hydrogen ions by the organic molecule.4 It is very tempting to spkculate as to why this electron transfer takes place with certain types of compounds and not with others but until we have completed our study of the reversible and irreversible reduction of a wide variety of substances, we prefer not to hazard an explanation of this fundamental question. There has been a tendency in recent years to represent all processes of oxidation and reduction as essentially electronic but this seems to us a rash and unwise generalization which greatly endangers the fruitfulness of the modern electrochemical conception of oxidation and reduction. As we have attempted to bring out in this discussion of irreversible reduction and catalytic hydrogenation, there are certain processes involving the gain or loss of hydrogen atoms (or their equivalents) which can be successfully formulated in terms of electron transfer and a t least one other which L. P. Hammett: J. Am. Chem. SOC.46, 7 (1924). Elektrochem. 15, 264 (~gcg). 8 F. Foerster: “Elektrochemie wasseriger Losungen,” pp. 21 1-21?. 4 As pointed out in another paper from this Laboratory (J. Am. Chem. SOC.44,2480 (1922)in certain cases of reversible reduction where the reduced compound forms a disodium salt, the actual reduction product formed in alkaline solution is a negatively charged ion; no hydrogen is apparantly involved in the reduction. However, since this represents a final equilibrium condition, no conclusions can be drawn from such facts in regard to the mechanism by which this equilibrium is reached. 2 2 .

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can not. Whether any given process is to be formulated in electrochemical terms is a problem to be settled by experiment and not by purely theoretical argument. In this paper we have confined out attention to the irreversible reduction of one particular type of substance in order to examine the difference between this process and catalytic hydrogenation. An extension of this general point of view to other types of irreversible reduction and also to cases of irreversible oxidation will be reserved for a later paper. The number of examples of such irreversible processes of reduction and oxidation is very great not only in organic but also in inorganic chemistry. The familiar fact, that sulfurous acid or sulfites may be oxidized to sulfuric acid or sulfates but that these can not be reduced in aqueous solution is a striking example of an irreversible oxidation. It should be possible to characterize such processes by determining their “apparent oxidation potential” much as we have measured the apparent reduction potential of a variety of organic compounds, and we are now examining the oxidation of aminophenols, and benzoin from this point of view. If our electrochemical interpretation of such irrevresible reductions and oxidations is correct, there must, in every instance, be a relatively fast reaction involving electron transfer followed by an essentially irreversible process leading to the final product. I n reversible processes the oxidized form can rapidly give up electrons and the reduced form rapidly acquire them whereas in the cases of irreversible oxidation or reduction only one of these steps takes place a t an appreciable rate. A more precise formulation of the mechanism of each type of .irreversible oxidation or reduction must await further experimentation.