Literally, disproportionation is a chemical reaction in which a single metal ion species undergoes an internal redox reaction and forms two different products in stoichiometric quantities. It is to be distinguished from such processes as partial combustion, which leads to a significant variety of chemical species. A disproportionation leads to very specific compounds, of varying oxidation state and chemical composition.

Disproportionation is a reaction that is more characteristic of the heavier transition elements and the f-block elements. Thus, while the chemistry of platinum(II) and platinum(IV) is well developed, attempts at the synthesis of the platinum(III) species were foiled by the disproportionation reaction yielding the +II and +IV oxidation states. The synthesis of Pt(III) complexes requires extraodinary ligands and special conditions. Note that the same is not true of palladium, where a well-developed +III oxidation state has been observed, although this oxidation state is often achieved through the disproportionation of the Pd(II) species giving Pd(O) as the other product. Also, disproportionation is sometimes masked by the chemical composition of the product. TlBr3 undergoes a reduction to give a compound that has the formula "TlBr₂". This species is actually the compound Tl⁺[TlBr₄]^- in which both Tl(I) and Tl(III) are present.

With the actinides, disproportionation reactions are quite common. For example, plutonium(IV) in water spontaneously reacts via disproportionation to yield the plutonium(VI) oxide and plutonium(III):

3Pu^{+ 4} + 2H₂O PuO₂^{+ 2} + 2Pu⁺³ + 4H⁺

Indeed, in 1M HClO₄ at room temperature, all four oxidation states of plutonium, Pu(III) to Pu(VI) are observed in equilibrium.