An oxidation–reduction (redox) reaction is an electron transfer between two species. It powers cellular respiration, iron corrosion, batteries, and industrial electrolysis.
Definitions
- A species that loses electrons is oxidized.
- A species that gains electrons is reduced.
This never happens in isolation: if someone loses electrons, someone else picks them up. Every oxidation is paired with a reduction.
The species that takes electrons is the oxidizing agent (it oxidizes the other). The one that donates them is the reducing agent (it reduces the other).
The redox couple
An oxidized species can be cycled back to its reduced form, and vice-versa. We write a redox couple as Ox / Red, for example:
- Cu²⁺ / Cu: Cu²⁺ + 2 e⁻ ⇌ Cu
- Fe³⁺ / Fe²⁺: Fe³⁺ + e⁻ ⇌ Fe²⁺
- H⁺ / H₂: 2 H⁺ + 2 e⁻ ⇌ H₂
Each half-equation shows how the oxidant takes electrons to become its conjugate reductant.
Building a full redox equation
Combine two half-equations by balancing electrons. Example: a copper strip dipped into silver nitrate solution.
Half-equations: - Ag⁺ + e⁻ ⇌ Ag (reduction) - Cu ⇌ Cu²⁺ + 2 e⁻ (oxidation)
To balance electrons, multiply the first by 2:
2 Ag⁺ + 2 e⁻ → 2 Ag Cu → Cu²⁺ + 2 e⁻
Sum: 2 Ag⁺ + Cu → 2 Ag + Cu²⁺
Copper is oxidized, silver is reduced. The strip gets coated with metallic silver and the solution turns blue (Cu²⁺).
Oxidation number
For trickier cases (no obvious monatomic ions), each atom is assigned an oxidation number (o.n.), a fictitious charge. A redox reaction changes the o.n. of at least one element.
Base rules: - Oxygen is usually −II. - Hydrogen is +I (except in hydrides, where it is −I). - The sum of o.n. in a neutral species is 0; in an ion, it equals the charge.
Why it matters
Redox is everywhere: - Batteries: we capture the flow of electrons in a wire. - Corrosion: iron loses electrons to atmospheric O₂. - Metabolism: glucose is oxidized to produce cellular energy. - Electrolysis: a current forces a non-spontaneous redox.
Understanding redox is understanding a large half of useful chemistry.