Definitions of oxidation and reduction in terms of electron transfer and oxidation number changes, rules for assigning oxidation numbers, identifying redox reactions, oxidising and reducing agents, tests, and everyday examples of redox chemistry.
Oxidation and reduction always occur together — you cannot have one without the other. This paired process, called a redox reaction, underpins rusting, combustion, bleaching, and the reactions that occur at electrodes during electrolysis.
Redox reactions are not confined to the laboratory. Several familiar processes involve the transfer of electrons:
| Process | Redox involvement |
|---|---|
| Rusting of iron | Iron is oxidised; oxygen is reduced |
| Bleaching of stains | Bleach acts as an oxidising agent, breaking down coloured molecules |
| Browning of cut fruit | Phenolic compounds in fruit are oxidised by oxygen in air |
| Food preservation with SO₂ or sodium sulfite | These substances act as reducing agents, preventing oxidation of food |
There are two complementary sets of definitions, both of which must be known.
| Process | Definition | Memory aid |
|---|---|---|
| Oxidation | Loss of electrons | OIL (Oxidation Is Loss) |
| Reduction | Gain of electrons | RIG (Reduction Is Gain) |
Together: OIL RIG
In the reaction between zinc and copper(II) sulfate:
Ionic equation:
Zinc: Zn → Zn²⁺ + 2e⁻ (loses electrons — oxidised)
Copper ions: Cu²⁺ + 2e⁻ → Cu (gains electrons — reduced)
Oxidation is an increase in oxidation number. Reduction is a decrease in oxidation number. This definition extends the concept to reactions where discrete electron transfer is less obvious.
An oxidation number (oxidation state) is the apparent charge an atom would have if all bonds in its compound were fully ionic. It is a bookkeeping tool for tracking electron shifts.
| Rule | Value |
|---|---|
| Uncombined element (e.g. Na, O₂, Cl₂) | 0 |
| Simple monatomic ion | Equals the ionic charge (e.g. Na⁺ = +1, Cl⁻ = −1) |
| Group I metals in compounds | Always +1 |
| Group II metals in compounds | Always +2 |
| Oxygen in compounds | Usually −2 (except in peroxides, where it is −1) |
| Hydrogen in compounds | Usually +1 (except in metal hydrides, where it is −1) |
| Sum of all oxidation numbers in a neutral compound | 0 |
| Sum of all oxidation numbers in a polyatomic ion | Equals the overall charge |
Find the oxidation number of manganese in KMnO₄.
K is +1. O is −2 (×4 = −8). The compound is neutral, so:
Manganese is +7 in potassium manganate(VII).
Find the oxidation number of sulfur in SO₄²⁻.
O is −2 (×4 = −8). The ion has charge 2−, so:
Sulfur is +6 in the sulfate ion.
A reaction is a redox reaction if at least one element changes oxidation number. The element whose oxidation number increases is oxidised; the element whose oxidation number decreases is reduced.
Iron: +2 → +3 (increase — oxidised)
Chlorine: 0 → −1 (decrease — reduced)
This is a redox reaction.
Copper: +2 → 0 (decrease — reduced)
Hydrogen: 0 → +1 (increase — oxidised)
CuO is reduced; H₂ is oxidised. This is a redox reaction.
An oxidising agent causes oxidation in another substance. To do this, it gains electrons — and so becomes reduced itself.
A reducing agent causes reduction in another substance. To do this, it loses electrons — and so becomes oxidised itself.
| Agent | What it does to the other substance | What happens to itself |
|---|---|---|
| Oxidising agent | Oxidises it (removes electrons) | Is itself reduced (gains electrons) |
| Reducing agent | Reduces it (supplies electrons) | Is itself oxidised (loses electrons) |
Common oxidising agents: oxygen, chlorine, acidified potassium manganate(VII), hydrogen peroxide, concentrated sulfuric acid.
Common reducing agents: hydrogen, carbon, carbon monoxide, reactive metals (e.g. zinc, iron), iron(II) ions.
Hydrogen peroxide is an example of a substance that can act as both an oxidising agent and a reducing agent, depending on the other reactant. When it reacts with a stronger oxidising agent, it gives up electrons and is oxidised (acts as reducing agent). When it reacts with a stronger reducing agent, it gains electrons and is reduced (acts as oxidising agent).
| Test | What it detects |
|---|---|
| Add acidified potassium manganate(VII) (purple): decolourises if a reducing agent is present | Reducing agent |
| Add acidified potassium dichromate(VI) (orange): turns green if a reducing agent is present | Reducing agent |
| Add potassium iodide solution: turns brown/black (iodine liberated) if an oxidising agent is present | Oxidising agent |
| Blue litmus paper: turns red in chlorine (oxidising agent decolourises the indicator) | Oxidising agent |