CATALYSIS OF IONIC OXIDATION-REDUCTION REACTIONS BY DYES AND ITS PROBABLE MECHANISM’ P. A. SHAFFER Department of BioEogieal Chemistry, Washington University School of Medicine, St. Louis, M ’ i S S O U T i Received June 11, 1966
In order to make clear the significance of the main results to be reported briefly in this paper, and indicated in its title, some introductory explanation is called for. The basic idea underlying nearly all of present-day theory of chemical kinetics is that only activated molecules react, and that the absorption of a critical increment of energy is the process of activation. From this point of view reaction velocity is thought to depend merely upon the proportion of the total molecules having a certain energy content or distribution; little or nothing is said about other properties the reacting molecules must possess. The r81e of catalysts, according to this view, is either to provide a different reaction path for which a lower activation-energy level suffices, or to furnish energy-rich molecules t o supply the activation energy. Without venturing to estimate, in general, how successful the theoretical treatment based on these concepts is in predicting specific reaction velocity, it seems to the writer that there is reason to doubt whether the energy concept alone is adequate to account for the velocity of many reactions in solution. Additional concepts of a different sort appear to be necessary. Because many, perhaps most, ionic reactions in solution are immeasurably rapid it has seemed to some logical to infer that ions in general are in an activated state and require for reaction little or no added energy of activation. In making this generalization it appears to have been overlooked that there are many cases where demonstrably ionic-and demonstrably “active”-substances, thermodynamically capable of substantially complete reaction, fail to react or react very slowly. And so far as I know no basis has been suggested for predicting whether a given ionic reaction will be immeasurably rapid or very slow. That there are cases of very slow reaction between demonstrably “active” ions-and also cases of very rapid reaction involving uncharged molecules-would seem to disprove the idea that the ionic state, without further qualification, represents a 1 Presented before the Thirteenth Colloid Symposium, held a t St. Louis, Missouri, June 11-13, 1936. 1021
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kinetically activated state. If this argument is valid, Lhe question arises as to what other factors determine reaction velocity. Some of the striking cases of slow ionic reactions are found among inorganic oxidation-reduction systems. Survey of a number of these (and comparison with rapid reactions of similar type in which one of the “slow” reactants participates with a different partner) led to the hypothesis that one important factor is €he equality or inequality of valence change between oxidant and reductant. One of the most remarkable cases is the failure of reaction even on boiling of ceric sulfate and thallous sulfate, both components of reversible, ionic, electromotively active couples. When separated as half-cells of a battery or with a number of catalysts the reaction is fairly rapid and complete a t room temperature (10). Another example: titanous chloride (a powerful reducing agent) and iodine in potassium iodide (a fairly strong oxidant) are toward many other substances very rapidly reacting reagents. Both are also electromotively active toward inert electrodes. It is therefore surprising to find that on mixing these two solutions (both in 1 M hydrochloric acid) reaction is quite slow, the half-period with 0.01 N concentrations being nearly an hour at room temperature. (A number of similar cases, some much more striking, have been examined.) According to the equi-valence change principle a probable explanation of the slowness of reaction is that Ti3+ can give up one electron but not two, while I,- (or 12) can accept two but not one. This explanation is supported by the fact that in this and other similar slow ionic oxidation reactions various substances have marked catalytic effect which are able to be reduced in two steps by the reductant and thereafter to yield two electrons t o the oxidant, or vice versa. Among the substances which exhibit marked catalytic effect on the Ti*+and I- reaction are various dyes, notably indigosulfonates, pyocyanine, and other phenazines, and to less extent rosindulin, flavins, indophenols, and anthraquinone- and naphthoquinone-sulfonates. Can it be shown that these dyes owe their catalytic effect in this reaction to an ability to accept or yield one as well as two electrons? If so, the hypothesis suggests itself that the same property may explain the marked catalytic effect of dyes in other reactions, notably in oxidations by molecular oxygen, including biological respiration. As a matter of fact it was this last problem, the “activation” or “transport” of 0 2 in biological oxidations, which directed our interest to a study of the simpler ionic reactions.2 It was observed some years ago that while a number of inorganic ionic 2 The points referred to were discussed in some detail in a paper before the Federation of American Societies for Experimental Biology a t Cincinnati, April 12, 1933, and in later addresses before several local sections of the American Chemical Society, but have not so far been published.
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oxidation reactions comply fairly well with the equi-valence change rule (that is, that reauctants requiring the loss of one electron to pass to a known state of higher oxidation are rapidly oxidized by sufficiently powerful oxidants capable of accepting one electron, but generally not rapidly by oxidants requiring two or more electrons), the components of organic reversible oxidation-reduction systems appeared to be exceptions to this rule. I n virtually all of the many cases of dyes which had been examined at that time, the oxidant and reductant differed by a pair of electrons, and the point w2s emphasized, notably by W. M. Clark, that the two electrons appeared to be gained or lost together, not separately. Yet the oxidized forms were found to be reduced about equally rapidly by, for example, titanous chloride (yielding one electron) and by stannous chloride (yielding two). Similarly the reduced forms are as rapidly oxidized by ferricyanide or ferric ions as by 1 2 (or la-), Hg++, or TP+, the last three being regarded as two-electron oxidants. This exceptional behavior of dyes was obviously a serious objection to the equi-valence change idea, and left unexplained the very remarkable catalytic action of the dyes in, for example, the reaction of Ti3+and IS-, where the addition of a very small concentration of indigocarmine enormously accelerates the oxidation. A possible way out of the difficulty was indicated by the discovery independently by Elema and by Michaelis of the two-step oxidation-reduction of the natural pigment pyocyanine, and identification of the intermediate between the fully oxidized and fully reduced forms as a unimolecular free radical, a semiquinone. The property thus established for this pigment (and soon thereafter also for rosindulin by Michaelis) of losing one electron in each of two separate steps (as well as two electrons in one step) is precisely the property needed to account for its reactions with both one- and two-electron oxidants and reductants and for compliance with our hypothesis as to the r6le of pyocyanine in catalysis. I n the meantime the theory of the two-step reaction and semiquinone formation was formulated by Michaelis and by Elema (6). Could it be that the same property is possessed by the other organic reversible systems above named, even by those for which it has been looked for and not found? This has proved to be the case with a number of these systems, of such diverse structure that it now seems permissible to predict that the two-step oxidation-reduction with intermediate semiquinone formation may be a rather general property, in varying degrees, with many organic systems. The first case we studied was that of the indigosulfonates. These had been systematically investigated in 1923 by Sullivan, Cohen, and Clark (ll), who found no evidence for the existence of an intermediate form. That an intermediate does nevertheless exist seemed probable from the fact that on cautious or slow reduction of indigocarmine at about pH 12 a cherry-color appears between the blue and the yellow of the fully oxi-
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dized and fully reduced forms. A careful reinvestigation of the four indigosulfonates, in collaboration with Dr. P. TV. Preisler and later with Dr. E. S. Hill and Mr. Robert Loeffel, gave data the interpretation of which, according to the principles of TV. 11.Clark and of Michaelis, proves the existence of well-characterized semiquinones. Only preliminary (5, 7, 9) reports have so far been published. I n the meantime the catalytically active natural alloxazine pigments, the flavins, were shown by Stern, Kuhn, Hastings, and Stare to exhibit the two-step property on reduction in acid solutions. (The flavins were found to be catalytic in the Ti3and 13- reaction.) We next turned t o the anthraquinonesulfonates, which have the advantage of fair stability in highly alkaline solution, where Michaelis predicted that anionic semiquinones should appear. The anthraquinonesulfonates had been studied by Conant and coworkers (3), who observed abnormal slopes of potentiometric titration curves, which they interpreted as due to meriquinone formation. Dr. E. S. Hill and I (4) found with several of these quinones clear evidence that semiquinones are formed on reduction. On cautious reduction (or reoxidation) an intermediate reveals itself by a characteristic color, the intensity of which is independent of dilution; the slope of the potential-titration curves increases systematically with rise of pH, the E: values being independent of concentration. Very recently Hill and I have reexamined also several hydroxynaphthaquinones, including phthiocol, a pure sample of which was kindly furnished by Professor R. J. Anderson. The latter substance had been studied by Ball (l),who concluded that there was no evidence for semiquinone formation or two-step oxidation. We find, however, what seems to be conclusive evidence with phthiocol of two-step oxidation-reduction at high pH, with semiquinone formation which reaches a maximum of about one-half of the total. The evidence is based on the criteria developed in Michaelis’ theoretical analysis and mentioned above in connection with the anthraquinonesulfonates. Statement of the data and their interpretation will be presented in papers now in preparation. Finally Michaelis (8) reports data showing that P-naphthaquinonesulfonate exhibits semiquinone formation. We had also observed the abnormal slopes of titration curves and the existence of a colored intermediate with this quinone. It is significant for our argument that the naphthaquinones and anthraquinones, as well as indigo, pyocyanine, flavin, and rosindulin are all markedly catalytic in the Ti3+and 13- reaction and also in the oxidation of glucose in sodium hydroxide solution by oxygen. There are of course many dyes capable of reversible oxidation-reduction and having catalytic properties in oxidation reactions, for which the twostep process has not so far been demonstrated. Whether in these cases
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also the existence of semiquinones can be proved remains to be seen. Rut in view of the variety of types of structure in which semiquinones have already been found, it seems not unlikely that this property is rather general among reversible organic oxidation-reduction systems, and the writer is disposed, tentatively, to attribute the catalytic activity of such systems in oxidation reactions to the possibly widespread existence of this property. It should be noted in this connection that there are doubtless other ways in which organic systems may exert catalytic activity,-such as complex formation. Also that substances capable of semiquinone formation are not by any means all equally catalytic in any one reaction. The relative levels of the potentials-of the catalyst with respect to that of the primary reductant and of the primary oxidant-appear to be very important. This point obviously concerns the old unsolved problem of the possible relation between dynamics and kinetics (2). I n conclusion I desire to mention a point concerning the behavior and potential of molecular oxygen, and a suggestion as to the possible rale of catalysts which promote aqueous oxidations by dissolved oxygen. The calculated potential usually cited for the oxygen electrode is 1.23 volts, a value which corresponds to a very intense oxidant, an intensity which if applied unchecked to biological systems should destroy them. The fact is somehow overlooked that this potential is for the reduction of 0 2 t o 2H20, Le., for the acceptance by 02 of four electrons (and 4Hf ions). But if the reductant yields to 0 2 only two electrons (and 2H+ ions are added) Hz02 is formed, and for this reaction the normal potential is only 0.68 volt, which at p H 7 is not very far above the intensity level of methemoglobin, an active natural oxidant. Since hydrogen peroxide is known to be formed in many autooxidations by air or oxygen, it seems probable that oxygen in aqueous solutions (including biological respiration) acts a t this lower intensity level. If, however, it could be shown that reversible catalysts exist which are capable of yielding simultaneously four electrons to 02,(and subsequently accepting electrons in steps of one or two) the full oxidizing intensity of 1.23 volts might perhaps become available. By anchoring such a catalyst a t a strategic point in the cell the large energy yield might conceivably be advantageously disposed of. Such a picture is pure speculation, but it provides a possible mechanistic reason for the known fact that the “oxygen-activating” type of respiratory catalyst is essential for the breathing of cells. I know of no other suggestion as to why this type of catalyst is essential. REFERENCES (1) BALL:J. Biol. Chem. 106, 515 (1934). (2) See, for example, BARRON:J. Biol. Chem. 97, 287 (1932). (3) CONANT, KAHN,FIESER,AND KURTZ:J. Am. Chem. SOC.44, 1382 (1922).
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(4) HILL AND SHAFFER:J. Biol. Chem. 114, li (1936). (5) LOEFFEL:J. Biol. Chem. 109, lvi (1935). (6) See MICHAELIS:Chem. Rev. 16, 243 (1935), for a summary of the subject. (7) MICHAELIS:Chem. Rev. 16, 282 (1935). (8) MICHAELIS:J. Am. Chem. Sac. 68,873 (1936). AND SHAFFER: Abstracts of the Eighty-sixth Meeting of the American (9) PREISLER Chemical Society, held in Chicago, September, 1933. (10) SHAFFER:J. Am. Chem. Sac. 66, 2169 (1933); Science 81, 464 (1935). COHEN,AND CLARK:U. S. Pub. Health Repts. 38,1669 (1923). (11) SULLIVAN,
