redox (reduction-oxidation) reaction
A redox (reduction-oxidation) reaction is any of a large class of chemical reactions, including many familiar processes such as combustion, corrosion, and respiration. Oxidation was originally defined simply as the combination of an element or compound with oxygen, or the removal of hydrogen from a compound; and reduction as combination with hydrogen or removal of oxygen. In the modern theory has been generalized: oxidation is defined as loss of electrons, and reduction as gain of electrons. The two always go together: there is an oxidizing agent which is reduced, and a reducing agent which is oxidized. Thus in the reaction
Fe3+ + I - → Fe2+ + ½I2
the iron (III) gains ion gains an electron and is reduced to iron (II), and the iodide ion loses ion loses an electron and is oxidized to iodine.
The strength of a redox reagent, expressing its tendency to react, is measured by the electrode potential of the half-reaction, and so redox reagents may be ranked in an extended electrochemical series. In a covalent compound or complex ion, each atom is assigned an oxidation number, which is the charge it would have if all the bonds were ionic – the electrons in a covalent bond between two atoms are assigned to the atom with the higher electronegativity. Thus the sulfur atom in sulfur dioxide has an oxidation number of +4, and each of the two oxygen atoms has an oxidation number of -2. When sulfur is oxidized by oxygen, S + O2 → SO2, its oxidation number increases from 0 (for elements, by definition) to +4, corresponding to a virtual loss of electrons.
Disproportionation is the simultaneous oxidation-reduction of the same chemical substance. An example is the disproportionation of copper (I) chloride, involving oxidation to copper (II) chloride and reduction to metallic copper:
2CuCl → CuCl2 + Cl.