Redox reactions may involve proton transfers and other bond-breaking and bond-making processes, as well as electron transfers, and therefore the equations involved are much more difficult to deal with than those describing acid-base reactions. In order to be able to recognize redox reactions, we need a method for keeping a careful account of all the electrons. This is done by assigning oxidation numbers to each atom before and after the reaction.

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For example, in NO3– the nitrogen is assigned an oxidation number of +5 and each oxygen an oxidation number of –2. This arbitrary assignment corresponds to the nitrogen’s having lost its original five valence electrons to the electronegative oxygens. In NO2, on the other hand, the nitrogen has an oxidation number of + 4 and may be thought of as having one valence electron for itself, that is, one more electron than it had in NO3–.

This arbitrarily assigned gain of one electron corresponds to reduction of the nitrogen atom on going from NO3– to NO2. As a general rule, reduction corresponds to a lowering of the oxidation number of some atom. Oxidation corresponds to increasing the oxidation number of some atom. Applying the oxidation number rules to the following equation, we have

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api/deki/files/71193/7e3520554330bd1c1248fba1651c6194.png?revision=1&size=bestfit&width=500&height=35" />The only atoms which change are Mn, from +7 to +2, a reduction, and S, from +4 to +6, an oxidation. The reaction is a redox process. SO2 has been oxidized by MnO4–, and so MnO4–is the oxidizing agent. MnO4– has been reduced by SO2, and so SO2 is the reducing agent.

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b) The oxidation numbers

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show that no redox has occurred. This is an acid-base reaction because a proton, but no electrons, has been transferred.