In oxidation-reduction reactions, the substances involved gain or lose electrons and show different electron configurations before and after the reaction. Oxidation is the overall process by which a specie in a chemical reaction loses one or more electrons and increases its state of oxidation. An oxidant is a substance capable of oxidizing a chemical specie; it acquires the electron or electrons lost by the specie and is itself reduced in the overall process. Reduction is the overall process in which a specie in a chemical reaction gains one or more electrons and decreases its state of oxidation. A reductant is a substance capable of reducing a chemical specie; it loses the electrons gained by the specie and is itself oxidized in the overall process.
An ORP reaction, therefore, involves an electron exchange capable of doing work. This capability is expressed in terms of potential for a half-cell, or electron, reaction. Table 7.7.5 lists the potentials for standard conditions, that is, where reactants and products are at unit activity. The voltages in this table are referenced to the standard hydrogen electrode (SHE), which is assigned the value of 0.000 V.
Note that the reactions in Table 7.7.5 are written as reductions, which is the almost universally used convention. In this section, the term ox/red indicates the oxidized form on the left side of the equation and the reduced form on the right. For example, the standard potential for ferric iron, Fe3+, reduced to ferrous iron, Fe2+, is written as E° Ox/Red = +0.770 V.
ERed/Ox = -EOx/Red simply means that the polarity is reversed when the reaction is written as an oxidation reaction. For example, Fe2+ = Fe3+ + e-. ERed/Ox = -0.770 V.
For example, in a common industrial process, hexava-lent chromium is reduced with a ferrous sulfate solution. The half-reactions are:
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