Types of redox reactions

Redox reactions come in four main types:

• Combination: two substances join to form one. C + O₂ → CO₂.
• Decomposition: one substance breaks into two or more. 2H₂O → 2H₂ + O₂. But not every decomposition is redox: in CaCO₃ → CaO + CO₂, no oxidation number changes.
• Displacement: an element replaces another in a compound. In metal displacement, a more reactive metal pushes out a less reactive one (Zn + CuSO₄ → ZnSO₄ + Cu). In non-metal displacement, sodium displaces hydrogen from water, and fluorine displaces chlorine: F₂ + 2Cl⁻ → 2F⁻ + Cl₂.
• Disproportionation: the same element is both oxidised and reduced. In 2H₂O₂ → 2H₂O + O₂, oxygen (at –1) goes down to –2 in water and up to 0 in O₂. Chlorine does the same in alkali: Cl₂ + 2OH⁻ → Cl⁻ + ClO⁻ + H₂O.

For disproportionation, the element must be in an oxidation state from which it can go both up and down, so it needs at least three possible oxidation states. Chlorine at +7 in ClO₄⁻ is already at its highest state; it cannot go up, so ClO₄⁻ cannot disproportionate.

Among the halogens, oxidising power falls down the group: F₂ > Cl₂ > Br₂ > I₂. So each halogen can displace those below it from their salts. Among the hydrohalic acids, reducing power rises: HF < HCl < HBr < HI.

A species with an element in an intermediate oxidation state (such as SO₂ or H₂O₂) can act either as an oxidant or as a reductant, depending on what it reacts with. One in its highest state can only go down, so it can act only as an oxidant; HNO₃, with nitrogen at +5, is an example.