Frederick R. Duke Institute for Atomic Research Iowa State University, Ames
I
Oxidation Reduction Mechanisms
I
During the past twenty years, and particularly since 1945, the field of oxidation mechanism studies has been a busy one, and enough is now known concerning such mechanisms that the information may well be incorporated into oxidation-reduction theory as it is taught to undergraduates. Inclusion of mechanistic along with thermodynamic theory in this area adds spice and interest to the subject for most students. When one is considering electron transfer from one particle to another, it is fair to ask how the electron makes its move; since the electron may be credited with no senses, it is immediately clear that an easy path must be provided for the transfer. In most of the cases which have been studied, this path is provided through close association of the oxidant and rednctant, the bond between them being of the Lewis acid-base variety. In other cases, the electron apparently is able to penetrate a thin-skinned insulation in moving from reductant to oxidant upon collision. These situations will now be considered in turn. Reactions Involving a Cation and an Anion
The simplest mechanism of electron transfer involving the acid-base properties of the reactants is that in which a cation reacts with an anion. Cations, particularly those of variable valence, are generally Lewis acids and anions are Lewis bases. After the bond between them is formed, intramolecular electron transfer is provided for by the bond. Some examples of this type of reaction are shown in Table 1. In the case of the Mn(Il1)-oxalato reaction (S), three complex ions are formed: the mono-oxalato, the dioxalato, and the trioxalato manganate ions. All of these ions acts a intermediates; however, the electron transfer rate is much reduced as the number of oxalate ions in the complex is increased. Consequently the reaction slows down as the oxalate ion concentration increases, either by the addition of the ion or by increasing the pH to convert oxalic acid to oxalate ion. A similar situation occurs in the cerinm(1V) oxidation of glycols (5). Again, the complexes with the greatest number of glycol molecules allow slowest electron transfer; pH has no effect here, because the glycol is complexed as the molecule rather than as glycolate ion. The intermediate in either of these situations following electron transfer is a free radical since only one electron is transferred. The probability is that this free radical is an oxalate ion minus an electron, G O , in the one Prrv ntrd ns part of thr Symposium orr Rvcent Advnncce i n Inmganil: C h m h r y , spunaorrd jointly by the Diriiiona crl Inorm n i r Chrmistrv and Chrmiral Eduwttion. nt the 137th Meeting of the ACS, levela and, Ohio, April, 1960.
-~~ ~~
~
case, and the radical RCHOH plus RCHO in the other. Such radicals may be studied through reactions which they induce in normally unreactive substances placed in the solution (17). [Such reactive intermediates have been called leoctates (18). The chemistry of reactates is not well studied and deserves much more work; their importance is emphasized by the realization that probably most oxidations would not proceed if reactates of moderate energies did not exist.] In the case of the Fe(II1)-I--reaction, the acid base intermediate may be Fe12+or FeI(OH) +, but not FeI++ (I). Apparently the radical ions 4- or IOH- have much lower free energies than the bare atom, I. In general, when halide ions are oxidised by one-electron oxidants, the reaction involves two halide ions leading to the electron-transfer intermediate, Xz-, where X is a halogen. Reactions of this sort, then, tend to he higher than first order in halogen. Another example of this effect is in the cerium(1V) oxidation of chloride (19), where CeC13+is an important intermediate. The cation may he either oxidant or reductant. For example in the case of the perchlorate oxidation of titanium(III), it is believed that weak complexes of cation and anion are formed prior to electron transfer
(4). Reactions Between Two Cations
Cations cannot bond directly, being acidic, hut two cations may combine with a single base provided the base has two or more electron pairs available for such bonding. Complexes of this type have been shown to be responsible for some electron transfer reactions and are believed to be quite generally involved in reactions between two cations. Examples are shown in Table 1. The first proof of existence of a basic bridge between two cations was accomplished by Taube and his coworkers (8). They chose as oxidant, Co(NH3)sClf+ and as reductant Cr+2. Co(II1) and Cr(II1) both have the property of exchanging ligands very slowly with substances in solution (20); thus, one could he sure of the position of the chloride ion before and after the reaction. Before reaction, it was certainly attached to cobalt; after reaction, it was shown to be attached to chromium. The fact that both cations were attached at the same time to the chloride was proven by placing radiochloride in the solution and finding no radiochloride attached to Cr(II1) at the end of the reaction (8); it is hardly conceivable that the transfer of chloride from Co(II1) to Cr(I1) could have occurred without exchange unless both cations were attached to the same chloride ion. Incidentally, ammonia does not act as a bridging group because it has available only one pair of electrons and is indeed a very poor electron-transfer bridge. Volume 38, Number 4, April 196 1
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Bases other than chloride ion were tested for their efficiency as bridges in the Co(III), Cr(I1) reaction. In order of decreasing effectiveness, the bases were found to arrange as follows: CN- > Br- > C1- > S04-2g HzO > NH8 (9). Recently, HnOwas shown to be a trne hridge, apparently through use of all four electron pairs on oxygen (25). In further work, Taube and his students showed that conjugated double-bonded systems acted very effectively as electron transfer bases. For example, fumaric acid,
occurred on the carbon-hydrogen bonds (24). In some respects, this is most easily explained by assuming that the electron spends sufficient time in the bridge so that the organic molecule has the properties of a free radical and may add and lose hydrogen or deuterium. Thus, it appears that the electron may be traced to a certain extent as it progresses through the bridge. Other reactions which are believed to include bridges of this sort are electron exchanges between Fe(II1) and Fe(I1) (12, IS), Tl(1) and TI(II1) ($6,26), Eu(111 and Eu(II1) (11), and a number of other reactions of this type which must be traced radiochemically. Some reactions using bridges which involve two different elements are the Fe(II1)-Sn(I1) reaction (7) and the above mentioned Co(II1)-Cr(I1) reactions.
allowed much more rapid electron transfer than did succinic acid, the latter being no better than acetic acid (21). The fact that the addition of hydrogen ion increased the rate a t which fumarate ion allows electron transfer was attributed to the fact that the use of the acid molecule, rather than fumarate ion as bridge, forces the cation to be hound to the double-bonded oxygen providing better electronic conduction between cations along the chain. I t was further shown that p-phthalic acid is an excellent electron bridge (22); in all cases the bridge in the beginning is attached to cobalt, and a t the end of the reaction is attached to chromium. To prove that Co(II1) and Cr(I1) had hold of opposite carboxyl groups rather than the same one during electron transfer, one of the carboxyl groups was replaced by an aldehyde group. This molecule, when still containing the conjugated system, is a good electron transfer agent; however, the aldehyde group does not stay with the chromium(III), being subject to ready replacement by a water molecule. When the aldehyde is used, no organic complex of Cr+aresults. In more recent experiments, using fumaric acid as a bridge in the Co(II1)-Cr(I1) reaction and using Dz0 as solvent, it was found that some exchange of D for H
Reactions Involving Anions Only
Table 1.
oxidant
Intermediates in Selected Examples of Oxidation Reduction Reactions
Reductant
Intermediate
Reference
HONSO~-~
(14)
\SO. BrOsMnOl-
HSOI Mn0,c2
(HO)~NSF~;01Br-OSOa M~OA-M-M~OA-~ where M is alkali metal ion +
162
/
HSOI-
+ HONO [H+l ONSOsH + HzO; ONSOsH + H 8 0 d
- gg
NSOsH.
PN-
Hobcacoo~ HSO.-
NO*-
where the sulfur accommodates five electron pairs and the sulfite acts as a base; this is proved by using hromate containing 018and finding that the resulting sulfate contains one oxygen in four of those originally part of the bromide ion. Again, it is not considered possible that the bromate oxygen could be transferred to the sulfur without exchange with solvent unless both Brand S had hold of the oxygen simultaneously. Another example is that of the nitrite-sulfite reaction (14). In this case, the sulfite ion acts as a hase and in the presence of hydrogen ion, displaces hydroxyl ions from the nitrite:
This is followed by further displacement of hydroxyl by sulfite until either nitrilodisulfonic acid or nitrilotrisulfonic acid is formed, and upon hydrolysis, sulfate and either hydroxylamine or ammonia are formed. The halate-halide reactions have been studied and are believed to involve HOX09and X - as follows:
Cu(CN),P Fe--CISnCb+ FeOHSnOHt' Co-X-Cr where X i s OH-, so,-, c1-, RI-
Although anions are primarily basic in the Lewis sense, acidic properties are also exhibited by many anions, either by accommodation of more than four electron pairs by the central element of the anion, or by displacing other bases normally part of the anion. For example, in the reaction between bromate and sulfite (15), the intermediate appears to be
Journal of Chemical Education
(16) (16)
The fufther steps depend upon the nature of X-, hut it is evident that the base, X-, has displaced the hase, HzO, from the oxidant in the process of the reaction (27, 28). One recently studied reaction of interest is the electron exchange between Mn04-2 and Mn04- (16). It has been found that alkali cations apparently intervene, possibly as a bridge, between the two anions. For instance, the reaction is faster in the presence of Cs+ than in the presence of Na+, and as Cs+ is substituted
for Na+, the rate increases linearly with the Cs+ concentration. There may he an intermediate or activated M + - Mn0,c2. complex of the type Mn0,-
-
(7) DUKE,F. R., AND PINKERTON, R. C., J. Am. Chem. Soe., 73, 3045 (1951). DUKE,F. R., AND PETERSON, N. C., Iowa Slate Coll. J. Sn'., 32, 89 (1957). (8) TAUBE,HENRY,A N D MYERS,H., J. Am. Chem. Soc., 76, 2102 ... (19.54). (9) TAUBE,HENRY,MYERS,H., AND RICK,R. L., J. Am. Chem. Soc., 75, 4118 (1953). (10) HORNIG, H. C., A N D LIBBY,W. F., J . Phw. Chem., 56, 896 (1952). C. S.. J. Phvs. Chem.,. 56.. 853 (11) . . MEIER.D. J.. AND GARNER. (1952). (12) SILVERMAN, J., AND DODSON, R. W., .I.Phgs. Chem., 56, 846 (19.521. D ~ K E , ' R., ~ . A N D WOLF, E. D., Iowa state COU. J . s&., 34, 157 (1959). RUTENBERG, A. C., HALPERIN,1.. AND TAUBE,HENRY, J. Am. Chem. Soc., 73, 4487 (1951). HALPERIN, J., A N D TAUBE,HENRY,J . Am. Chem. Soc., 74, 375, 382 (1952). . . GJERTSEN.L.. AND WAHL.A. C.. J. A n . Chem. Soc... 81.. 1572 (1959): (17) I. M.. AND STENGER. V. A,. "Volumetric And~- , KOLTHOFF. ysis," 2nd ed., Interscience Publishers, Ine., New York, 1942, Vol. I, pp. 16880. (18) DUKE,F. R., "Mechanisms of Oxidation-Reduotion Reactions," in "Treatise on Analytical Chemistry," edited by Kolthoff and Elving, Interscience fincyclopedia, Inc., New York, Val. 1, pp. 63&59. (19) . . DUKE.F. R.. AND BORCEERS. . C... J.. Am. Chem. Sac... 75.. 5186 (1953j. (20) TAUBE, HENRY, Chem. Revs., 50, 69 (1952). (21) TAURE,HENRY,J. Am. Chem. Soc., 77, 4481 (1955). (22) SEBERA,H., AND TAUBE,HENRY,ta be published. See Can. J . Chem., 47, 129 (1959). (23) KRUSE,W., AND TAUBE,HENRY,J . Am. Chem. Soc., 82, 526 (1960). (24) FRASER,R. T. M., AND TAUBE,H., J . Am. Chem. Soe., 81, 5514 (1059). 125) W. C. E.. J . Chem. . B A S ~ R S TK.. G.. AND HIGGENSON. Soc., 3 0 k (1953). B., J . Phys. Chem., 60, 1015 (26) DUKE, F. R., AND BORNONG, (1956). (27) Srcamn, J., J. Chem. Phys., 55, 758 (1958). (28) EDWARDS, J. D., Chem. Revs., 50, 455 (1952). (29) THOMPSON, R., J . Am. Chem. Sac., 70, 1045 (1948). (30) EICHLEE,E., AND WAHL,A. C., J . Am. Chem. Soe., 80, 4145 (1958). (31) ZWICKEL, A. M., A N D TAUBE,H., J. Am. Chem. Soc., 81, 2915 (19591. (32) ZW&IN&I, B. J., MARCUS, R. J., AND EYFSNG, H., C h . Revs., 55, 157 (1955). \----,-
Reactions Nol Involving Acid-Bare Properties
There are a number of reactions which proceed readily and yet do not fall into any of the classes discussed above. For example, the electron exchange between ferrocyanide and ferricyanide proceeds at a greater rate than the cyanide exchange; thus, there can be no removal of cyanide ligands in the reaction (39). Similarly, the rapid reactions of the 1-10 phenanthroline complexes of Fe(I1) and Fe(II1) and similar complexes cannot be explained on the basis of acid-base properties of the ions (50, 51). It appears that here we have complex ions containing unsaturated and conjugated systems such that the electron can get very close to the outside boundary of the ion. Further, such ions are weakly solvated so that collisions usually involve the ion directly with no intervening water molecules. In such cases, the insulation between particles becomes very thin and the electron may make its way through from one particle to another without difficulty. This type of reaction at present is best explained on the basis of quantum mechanical tunneling, the latter phenomenon being responsible for electron transfer (Sf?). In general, it may be stated that if the electron can be involved close enough to the outer boundary of the particle, transfer is probably due to the thin electrical insulation between particles. Literature Cited (1) HERSHEY, A. V., A N D BRAY,W. C., J. Am. Chem. Sac., 58, 1760 (1936). (2) TAUBE,HENRY,J. A n . Chem. Soc., 70, 1216 (1948). (3) D u n , F. R., J . Am. Chem. Soc., 69, 2885 (1947). (4) DUKE,F. R., AND QUINNEY, P., J . Am. Chem. Soc., 76, 3800 (1954). . . (5) DUKE,F. R., A N D BREMER,R. F.,J. Am. Chem. Sor., 73, 5169 (1951). . . (6) DUKE,F. R., AND COURTNEY, W. G., J . Phy8 Chem., 56, 19 (1952).
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