Oxidation and Reduction

Oxidation and reduction were originally defined in terms of oxygen.

Oxidation meant gain of oxygen, reduction meant loss of oxygen.

Both definitions still work, but in modern chemistry we use a broader definition.

Oxidation is the loss of electrons.

Reduction is the gain of electrons.

A useful memory aid: OIL RIG — Oxidation Is Loss, Reduction Is Gain.

Oxidation and reduction always happen together: if one species loses electrons, another must gain them.

A reaction in which both processes happen is called a redox reaction.

The species that accepts electrons is called the oxidising agent. By gaining electrons, it is itself reduced.

The species that gives up electrons is called the reducing agent. By losing electrons, it is itself oxidised.

Strong oxidising agents include the halogens, KMnO₄ and K₂Cr₂O₇.

Strong reducing agents include reactive metals (e.g. group 1 and 2), and species like H₂ and CO.

Oxidation States (Oxidation Numbers)

Oxidation state is a way of tracking electrons in compounds, even when the bonding is covalent rather than ionic.

It tells you the imaginary charge an atom would carry if every bond it made were purely ionic — that is, if every shared electron were assigned to the more electronegative atom.

There are a few rules to find an oxidation state.

The oxidation state of an uncombined element is 0. So O₂, Na metal, Cl₂ are all 0.

The oxidation state of a simple ion is equal to its charge. Na⁺ is +1, S²⁻ is −2.

Group 1 metals are always +1 in compounds, group 2 are always +2, aluminium is always +3.

Hydrogen is +1 in most compounds, but −1 in metal hydrides (NaH, CaH₂).

Oxygen is −2 in most compounds, but −1 in peroxides (H₂O₂) and +2 in OF₂.

Fluorine is always −1 in compounds.

The sum of oxidation states in a neutral compound is 0.

The sum of oxidation states in an ion equals the charge of the ion.

These rules let you find unknown oxidation states.

In MnO₄⁻, each O is −2 and there are four of them, contributing −8 in total. The overall ion charge is −1, so Mn must be +7.

In Cr₂O₇²⁻, the seven O atoms give −14. The overall charge is −2, so the two Cr atoms together must be +12 — meaning each Cr is +6.

When an oxidation state increases, the atom has been oxidised. When it decreases, the atom has been reduced.

Tracking changes in oxidation states is the easiest way to identify what is being oxidised and what is being reduced in a complex reaction.

Half-Equations

A half-equation shows only one half of a redox reaction — either the oxidation or the reduction part.

Electrons are written explicitly as e⁻.

For the oxidation of magnesium:

Mg → Mg²⁺ + 2e⁻.

Magnesium loses two electrons. The electrons appear on the products side.

For the reduction of chlorine:

Cl₂ + 2e⁻ → 2Cl⁻.

Chlorine gains two electrons. The electrons appear on the reactants side.

To balance a more complex half-equation, follow a standard procedure.

First, balance the main atoms — everything except H and O.

Next, balance oxygens by adding H₂O.

Next, balance hydrogens by adding H⁺.

Finally, balance the charges by adding electrons.

Example: the reduction of manganate(VII) in acidic solution.

Start with MnO₄⁻ → Mn²⁺.

Oxygens: add 4H₂O on the right.

Hydrogens: add 8H⁺ on the left.

Charges: left side is (−1) + 8(+1) = +7; right side is (+2). To balance, add 5 electrons to the left.

Final equation: MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O.

Combining Half-Equations into Redox Equations

To write a full redox equation, combine the oxidation half-equation with the reduction half-equation so that the electrons cancel.

The trick: multiply each half-equation by whatever factor is needed to make the number of electrons equal on both sides.

Then add the two together and cancel the electrons.

Anything else that appears on both sides — H⁺, H₂O, spectator species — should also be cancelled to give the cleanest form.

Example: manganate(VII) reacting with iron(II) in acidic solution.

Reduction half: MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O.

Oxidation half: Fe²⁺ → Fe³⁺ + e⁻.

Multiply the oxidation half by 5 so each side has 5 electrons.

Add them: MnO₄⁻ + 8H⁺ + 5Fe²⁺ → Mn²⁺ + 4H₂O + 5Fe³⁺.

Electrons cancel and the equation is balanced for both atoms and charges.

This is the equation behind the classic potassium manganate(VII) titration used to determine the concentration of iron(II) in solution. The deep purple MnO₄⁻ ion is reduced to almost colourless Mn²⁺, so the end point is reached when one extra drop turns the solution permanently pink.

Common Redox Reactions to Know

Several redox reactions come up repeatedly in AQA exams.

Reactions of group 2 metals with water and acid are redox: the metal is oxidised, hydrogen is reduced.

Halogen displacement reactions are redox: a more reactive halogen oxidises the halide ion of a less reactive one.

Reactions of metals with acids are redox: the metal is oxidised to a cation, H⁺ is reduced to H₂ gas.

Disproportionation is a special case where the same element is simultaneously oxidised and reduced. The reaction of chlorine with cold dilute sodium hydroxide is the classic example: Cl₂ + 2NaOH → NaCl + NaClO + H₂O. Chlorine goes from 0 to −1 (reduced) and from 0 to +1 (oxidised) in the same reaction.

More advanced redox topics — electrode potentials, electrochemical cells and standard hydrogen electrodes — are covered in the A2 topic 3.1.11.