Oxidation and Reduction

A redox (oxidation-reduction) reaction is a chemical reaction involving the transfer of electrons. For a redox reaction to occur, there must always be one species being oxidised (losing electrons), and another species being reduced (gaining electrons).

Tip: Use OIL RIG = Oxidation is Loss, Reduction is Gain, to help remember how electrons move for each process.

The oxidised species is the reductant or reducing agent. It is releasing electrons for the other another species to gain. The reduced species is the oxidant or oxidising agent. It is taking electrons of another species and causing oxidation.

Oxidation Numbers

Which species is being oxidised and which is reduced cannot be determined from simply looking at the balanced redox reaction. Oxidation numbers can help us determine if it is a redox reaction, and most importantly which species is being reduced and which is being oxidised.

Oxidation numbers themselves tell us the charge the species would have if its bonds were purely ionic (meaning no sharing of electrons). The number of this charge is what helps us determine the movement of electrons.

1. Monoatomic, diatomic or polyatomic elements by themselves have an oxidation number of 0 (e.g. Mg (s) or Ni (g)).

2. Monoatomic ions have an oxidation number equal to their charge (e.g. Mg⁺² (aq)).

3. When bonded to other elements, oxygen has an oxidation number of -2. If it is the oxygen of a peroxide (e.g. FeO2, ZnO2, Na2O2), its oxidation number is +2.

4. When bonded to other elements, hydrogen has an oxidation number of +1. If it is the hydrogen of a metal hydride (e.g. NaH, ZnH2, MgH2), its oxidation number is -1.

The oxidation number of atoms in a neutral compound always sum to 0. The oxidation number of atoms in a polyatomic ion must sum to its overall charge.

If an atom in a molecule shows an increase in oxidation number, it has lost electrons and has been oxidized.

If an atom in a molecule shows a decrease in an oxidation number, it has gained electrons and has been reduced.

Redox Equations

A redox reaction is made up of 2 half reactions, which represent the oxidation process and the reduction process.

Each individual half equation must be balanced before combining them together to form a redox equation. This requires the elements and overall charges on both sides to be equal.

The number of electrons lost must also be equal to the number of electrons gained. Therefore, the half equations should be multiplied by the lowest common multiple that allows the number of electrons added to each half equation to be equal.