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J. DUDLEY HERRON Pvrdue University West Lofoyene, lndiono 47907
What is Oxidation? There is a maxim in education that "teachers teach as they were taught." This probably explains why many teaching practices persist without apparent logical justification. How else can we explain the fact that most teachers continue to define oxidation and reduction in terms of a gain and loss of electrons rather than in terms of a change in oxidation numher? Virtually all general chemistry texts in use today present a discussion of oxidation number and when redox reactions are discussed, oxidation numbers are assigned to those atoms which undergo some change in oxidation numher. The definition of oxidation as a gain in oxidation number and reduction as a decrease in oxidation numher is straight forward, chemically correct, and easy for students to follow. Still, it appears that most teachers continue to define oxidation as a loss of electrons and introduce a complicating factor that is unnecessary and sometimes confusing. In introducing oxidation and reduction, the teacher normally begins with a simple example of an element combining with oxygen; for example, 2Ca (s)
+ 0, (g)= XaO (s)
There may he some discussion, of the historical definition of oxidation as the combination of an element with oxygen with reduction as the reverse process, reducing the oxide to the elements. The discussion quickly proceeds to the observation that in this reaction the calcium apparently gives up two electrons to form Caz+ and the oxygen gains two electrons to form OZ-. Oxidation is then defined as a loss of electrons while reduction represents the opposite process, the gain of electrons. A further example of an oxidation reaction which does not involve oxygen is given and the generality of the definition to any reaction which involves an exchange of electrons is then established. No real harm is done at this point. The definition seems to be related to the chemical process that is taking place and it is general enough to apply to many reactions that will he encountered later. However, by focusing on the gain and loss of electrons rather than the change in oxidation numher, an unnecessary complication is introduced. The definition implies that oxidation always involves a loss of electrons and students are expected to rememher a definition that no longer contains any logical clues as to its meaning-a minor hut important point when students must rememher a lot of new terms. Unlike the definition of oxidation as a loss of electrons, the historical definition "made sense." If an element combined with oxygen, it was called oxidation; if the compound was reduced to its elemental form, it was called reduction. These are "logical" names for the processes. We can see the same kind of logic in "ionization energy" as the term that we assign to the energy required to form a positive ion from a neutral atom or in "electronegativity," "atomic weight," "bond length" and many other terms
that are part of the chemist's special jargon. But what association can a student make between "oxidation" and a "loss of electrons?" Even worse. the associations that are obviously made with "reduction" are opposite to the correct definition, a aain of electrons. Of course chemists understand the hist'orical origin of the terms but this does not help the beginning student who is struggling to remember the hundreds of new terms that he is bombarded with in the introductory course. It seems rather fortuitous that a new word association follows if we s i m.~.l vdefine oxidation as a gain in oxidation numher and reduction as a decrease in oxidation numher. Althoueh no clue is found in the definition of oxidation, it is quite iogical to associate "decrease" with "reduction" and the student has a strong. association to aid in remembering the new terms. Defining oxidation and reduction in terms of a gain and loss in oxidation number requires very little change in the normal discussion of redox reactions. In the calcium example given previously, the discussion simply focuses on the fact that we assign an oxidation numher of zero to calcium metal and a plus two to the calcium that is combined with oxygen. In going from a zero to a plus two, the calcium has increased in oxidation numher and, by definition, we say that the calcium has been oxidized. The instmctor may even wish to point out that in this case the change in oxidation numher was due to the loss of two electrons by each calcium atom to produce a calcium ion with a plus two charge. However, emphasis should be placed on the change in oxidation number as an indication of oxidation and reduction rather than on the transfer of electrons. Teachers who define oxidation and reduction in terms of loss and gain of electrons normally assign oxidation numbers to elements but continue to focus on the electron transfer. This is unnecessary and leads to additional confusion. A discussion of the typical redox reaction
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usually proceeds something like this: (After assigning oxidation numbers to the appropriate elements) "You see that the charge (or oxidation numher) on zinc has changed from zero to plus 2. Has zinc lost electrons or gained electrons?" "Then was the zinc oxidized or was it reduced?" There are alwavs some beainnina students who respond, "Zinc has gained electrons,%ndihe instructor then explains, "No. Remember that electrons are negatively charged and since the charge on the zinc is increasing, we must have lost some of those negatively charged electrons." Eventually, the students appear to understand and the lesson proceeds. The test later reveals that comorehension was not as comolete as we had hoped. The apparent contradiction of ''l&ing2' something Ghich results in a larger numerical value is confusing to students who are struggling with a large number of newly acquired concepts which are not firmly fixed in memory. This confusion can certainly he overcome and the majority of our students eventually get it sorted out but the point is that it can easily he avoided entirely. If oxidation is defined as a gain in oxidation numher, there is no need to ask whether electrons were gained or lost a t all. Since the oxidation numher of zinc increased in this reaction, the zinc must have undergone oxidation. I t follows immediately from the definition. The arguments for defining oxidation and reduction in terms of the change in oxidation number rather than in terms of electron transfer may appear trivial to many readers and, standing alone, perhaps they carry little weight. However, there is another, more weighty argument. The simple fact is that by defining oxidation and reduction in terms of electron transfer, many students are misled into believing that all redox reactions necessarily involve an exchange of electrons. This, of course, is just not so. Even in the simple case of burning charcoal to produce carhon dioxide, it is doubtful that one could make a very strong case for the supposition that carhon has lost electrons and that oxygen has gained them. Carbon dioxide is a respectable covalent molecule and the assigned oxidation numbers of +4 for carbon and -2 for oxygen are only possible when we artificially assume that all bonds are ionic and assign bonding electrons to the more electronegative element involved. Perhaps a more convincing example is the reaction between the chlorate ion and sulfite ion to produce chloride and sulfate
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The evidence obtained by conducting the reaction using KC103 enriched with oxygen-18 indicates that the reaction proceeds by the direct transfer of an oxygen atom from C 1 0 3 to the S0s2-. There is no evidence of a transfer of electrons from one atom to another. There is, however, a change in oxidation number and the reaction is considered to be an appropriate example of oxidation and reduction. Authors of most introductory texts-even those who define oxidation as a loss of electrons-make the point that oxidation and reduction need not involve a transfer of electrons hut the point apparently gets lost in the minds of students. When I discuss redox reactions in staff meet-
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ings with graduate teaching assistants who are working in the introductory chemistry course that I teach, most of them are surprised when I suggest that electron transfer may not be involved. I get the same response with undergraduate seniors who are preparing to he high school chemistry teachers. Try it with your own biased sample before you conclude that we do not mislead students when we define oxidation and reduction in terms of electron transfer! Teachers who are accustomed to defining oxidation and reduction in terms of electron transfer react to my suggestions with such questions as, "But what about electrochemistry? Surely you want to discuss electron transfer there?" or "What do you do about balancing redox equations? Don't you use half-reactions? You must talk about electron transfer then!" I find no difficulty with these arguments. Of course I talk about electrons when I discuss electrochemistry. When I discuss a Daniell cell I point out that the zinc metal goes into solution and forms ZnZ+ ions and that in the process, it gives up electrons to the external circuit; that the CuZ+ ions in solution accept electrons from the external circuit to form neutral copper atoms which plate out on the copper electrode. But I do not say, "The zinc' was oxidized because it lost electrons." It is just as easy to say, "The zinc was oxidized because it increased in oxidation numher." It is incidental that this is accomplished through the loss of electrons. In like manner, I find no problem in balancing redox equations. I normally introduce redox reactions before electrochemistry is discussed and equations are balanced by the gain and loss in oxidation numher with no discussion of electrons. I do not break the reaction down into half-reactions a t this point because it appears too artificial for reactions that occur in a single beaker. However, this can be done if desired. When electrochemistry is discussed, I do use half-reactions. There is some logic in discussing what occurs in each half-cell, and it is a simple matter to use these half-reactions t o balance the equation for the overall reaction. The electrons shown in the equation for the half-reaction are not there because of some artificial device hut because there is evidence that they pass through the external circuit. Whether the mechanics of balancing the equation are carried out by balancing the gain and loss of electrons or by balancing the gain and loss in oxidation number makes little difference since the result is the same. I generally leave it u p to the student to use the method that he finds more convenient. However, I try to point out that when the method of balancing halfreactions using gain and loss of electrons is applied to other redox reactions, it is simply a mechanical tool for balancing the equation and does not necessarily indicate the chemistryof the process. Do we continue to define oxidation and reduction in terms of electron transfer simply because that is the way we were taught? Is there some logical or pedagogically sound reason that we should not discard this definition in favor of the more general one in terms of gain and loss in oxidation number? If there is, I have not found it. 1 have found what I consider to be sound reasons for the converse and until I hear arguments to the contrary, I will continue to say, "Oxidation is a gain in oxidation number and reduction is a decrease in oxidation number." I think you JDH should too.
