Oxidation states

Iron can be Fe²⁺ or Fe³⁺; manganese ranges from +2 to +7. Transition metals show variable oxidation states. The reason: the (n–1)d and ns subshells are very close in energy, so electrons from both can be used in bonding.

Their oxidation states usually differ by one unit (Fe²⁺, Fe³⁺). In p-block metals they differ by two (Sn²⁺, Sn⁴⁺).

Manganese shows the most states, from +2 up to +7, in the middle of the series. At the ends there is little choice: scandium shows only +3 and zinc only +2.

The highest oxidation states turn up in oxides, fluorides and oxoanions, such as Mn₂O₇, MnO₄⁻ and CrO₄²⁻. Oxygen and fluorine are small and highly electronegative, so they can stabilise high oxidation states. Oxygen can also form multiple bonds to the metal.

To find an oxidation state, add up the charges. In Cr₂O₇²⁻: 2x + 7(–2) = –2, so x = +6.

The oxides follow a pattern:

• Low oxidation state (CrO, +2): basic and ionic.
• Middle state (Cr₂O₃, +3): amphoteric, reacting with both acids and bases.
• High state (CrO₃, +6): acidic and covalent.

Disproportionation is a reaction in which one species is both oxidised and reduced, giving a higher and a lower state. In water, copper(I) does this: 2Cu⁺(aq) → Cu²⁺(aq) + Cu(s). Cu²⁺ wins because its very large, negative hydration enthalpy more than pays for removing the second electron.