Oxidation-Reduction: A Re-evaluation CALVIN A. VANDERWERF, ARTHUR W. DAVIDSON, and HARRY H. SISLER University of Kansas, Lawrence, Kansas
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agent, were still in OME years ago, Dr. Malcolm Dole, in discussing was provided by the oxidizing the pH concept, said, in effect: "When Sfirensen general use. This practice gradually fell into disfavor in the fivefirst proposed the use of the pH scale, he supposed i t to be a measure of hydrogen-ion concentration. After it vear oeriod from 1925 to 1930. Durine., these years. had been established that electrode potential is a func- several of the more enthusiastic idvocates of the ion tion of activity rather than of concentration, pH was electron method of balancing oxidation-reduction equasupposed to measure hydrogen-ion activity. Still later, tions performed a notable service by bringing into sharp when i t became clear that neither single electrode focus, particularly through their contributions to THIS potentials nor single ion activities can be determined JOURNAL; the question of the true nature of oxidation experimentally, that definition had also to be aban- and reduction. So convincingly did these advocates doned. Now we cannot say with certainty what it is present their case that today the definition of oxidation that pH measures, except that it' is something very as a chemical change involving the loss of electrons is almost universally accepted. Reduction is defined as impo&ant." The history of the oxidation-reduction concept pre- a chemical change involving a gain of electrons. From sents a rather strikine ~arallelto that of DH. We do these definitions, the inference has been drawn that the not, indeed, wish to suggest that oxidation-reduction basis of all systematic methods for balancing oxidationcan no longer be defined except as "something very reduction equations is the obvious fact that in any comimportant." Yet i t is obvious that the accepted use plete electron transfer the total number of electrons of this term has wandered far from its original meaning, gained by one atom or group of atoms must be equal and we propose to maintain that its present significance to the total number of electrons lost by a second atom is quite different even from that which has come to be or group of atoms. The present authors have always been sympathetic attributed to it during the past four decades. toward the electron-transfer approach to the problem of HISTORICAL DEVELOPMENT oxidation and reduction; a t least one of the three The earliest application of the term "oxidation" was, learned the ion-electron method of balancing redox as is implied in the word itself, to reactions consisting equations before he was introduced to any other of the combination of other substances with oxygen; method. After years of presenting the electron-transfer "reduction," on the other hand, meant the removal of idea of oxidation and reduction to their students, howoxygen from its compounds. Thus, a typical oxidation ever, the authors have come to view their own teaching with skepticism, and to believe that, in the interests was such a reaction as of logic and consistency, a re-investigation of the mean2Cu + on 2 c u o ing and application of the terms "oxidation" and "reand a typical reduction, duction" is necessary. Hence we are attempting in this paper critically to examine the accepted definitions in CuO + Hl Cu H.0 the light of present knowledge of atomic and molecular Soon, however, the scope of "oxidation" was broad- structure, and perhaps to open a frank discussion conened so that combination with sulfur, chlorine, or other cerning the best method for presenting the concept, nonmetallic elements was often included in the same from a pedagogical standpoint. category as combination with oxygen, and the removal In the early years of the present century, largely as a of such elements came to be regarded as a form of reduc- result of the pioneer work of J. J. Thomson1in correlattion. From this point, i t was but a short step to the ing chemical properties, and especially valence, with definition of oxidation as any reaction in which the electron structure, the concept of oxidation-reduction valence of an element toward oxygen or any other non- came to be associated with the electrical state of an metallic element is increased, and of reduction as any element. Stieglitq2 particularly, emphasized the fact reaction in which such valence is diminished. that the changes which take place a t the anode and In spite of these broadened definitions, however, the cathode of an electrochemical cell are oxidation and idea that oxidation necessarily involved the element reduction, respectively. The nature of half-cell reacoxygen persisted, and even as recently as 20 years ago, ' THOMSON, J. J., Phil. Mq.,7, 2 3 7 4 5 (1904). antiquated and cumbersome methods for balancing "6% for example, STIBDLITZ, J., "The Elements of Qualitative oxidation-reduction equations, based upon improbable Chemical Analysis," The Century Company, New York, 1911, DO. 251-5. partial reactions in which oxygen (sometimes"nascent") . 450
- +
tions, together with the assumption of electron transfer in the formation of valence bonds, served as the basis for the ion-electron method of balancing equations. NEED FOR RE-INVESTIGATION By the time that Lewis8 first presented his idea of the covalent bond, consisting of a shared electron pair, a number of prominent chemists had already fully accepted the electronic definitions of oxidation and reduction and were teaching the balancing of equations by the ion-electron method. With the advent of the Lewis theory, i t should have become clear, as Remick4 has pointed out, that i t was time for chemists to re-examine their definitions of oxidation and reduction. Yet such a re-evaluation was considerably delayed, for the true nature of the covalent bond was not a t first fully appreciated. Until 1925, or thereabouts, perhaps because of the continued influence of the ideas of Tbomson, who had assumed that all valence forces resulted from the complete transfer of electrons, every interatomic bond, even in organic compounds, was still generally thought of as being a t least vaguely "polar" in nature; hence valence was regarded as being either definitely positive or definitely negative. Thus there arose the concept of valence number or oxidation state as applied to an element, which will play an important part in our subsequent discussion. We now recognize that all chemical reactions may be considered to involve either the complete transfer or the sharing (or resharing) of electrons. Confusion has arisen from the fact that these two types of reactions are not clearly distinguishable, for one merges gradually into the other. In so far as the division is clear cut, however, are we imparting the total intended meaning t o the terms "oxidation" and "reduction" if we classify as redox reactions only those chemical changes in which electrons are transferred outright? That question bas probably perplexed all who have written on the subject. Some writers have answered it entirely in the affirmative, no doubt with personal misgivings in varying degrees. Many have dodged it, while a few have suggested that the answer should be in the negative. Meanwhile the..idea of complete electron transfer persists as the foundation for the accepted definitions of oxidation and reduction.
ELECTRON TRANSFER?
simple ions or ionic compounds. This relationship may be illustrated by means of such reactions as
In each case, the first reactant has lost electrons and bas been oxidized, whereas the second reactant has gained electrons and has been reduced. All clear-cut cases of complete electron transfer in a chemical reaction are admittedly oxidation-reduction; i t is obviously a fallacy to conclude, however, that all oxidation-reduction is electron transfer. Consider, for example, the complete combustion of carbon in air: C
+ 02-cog
By universal agreement, the carbon is oxidized by four units. But the assumption that carbon has lost four electrons in its oxidation to carbon dioxide requires that the carbon exist in the compound as C+4 ion, a postulate which is, of course, entirely inadmissible. It might be argued that the four carbon electrons are shifted partially from the carbon atom toward the oxygen atoms; but this claim can hardly be made in the case of the oxidation of carbon to carbon monoxide, the very small or zero dipole moments of which indicates that no such electron displacement has taken place in its f ~ r m a t i o n . ~ Evidence favoring the complete equivalence of reduction and electron gain is not much more convincing. Consider the familiar inorganic redox reaction :
- + 5S08- + 6H+' 2MnOi
-
2Mnf+
+ 5SO,= + 3H.O
6 PAWNG, L., "Nature of the Chemical Band," Cornell University P r w , Ithaca, New York, 1940, p. 135. a The case for complete electron loss in oxidation is even more difficult to uphold in such reactions as the oxidation of trinitrobenzene topicric acid, or of acetaldehyde to acetate ion in alkaline solution:
CHIC^
+ 2KMnOl + ROH + ~cIL
