Rusting

Rusting is the oxidation of iron to hydrated iron(III) oxide. It requires both water and oxygen to occur, and neither on its own is enough. Malay: pengaratan · Chinese: 生锈.

Rusting is a redox process. The iron atoms lose electrons to form iron(II) ions (Fe to Fe²⁺ plus 2e-), which are then oxidised further; oxygen dissolved in the water gains those electrons, and the product is the reddish-brown hydrated iron(III) oxide we call rust. Because both water and oxygen are needed, iron kept fully dry or fully sealed from air does not rust, while salt, acids and contact with a less reactive metal all speed rusting up.

Example. The classic test uses three test tubes: an iron nail with both air and water rusts; a nail in boiled water under oil, with no air, does not; and a nail in dry air over a drying agent does not. This shows that water and oxygen together are the necessary conditions.

Confusion to avoid. Do not treat rusting as a general word for corrosion. Rusting is specifically the corrosion of iron; corrosion is the wider term for the oxidation of any metal by its surroundings. Also do not say rusting needs only oxygen or only water; the mark depends on stating that both are required.

Preventing it. Because rusting is oxidation of iron, it is prevented by keeping out water and oxygen, or by sacrificial protection with a more reactive metal, or by galvanising. These methods build directly on the reactivity series.

A closer look at the chemistry. Iron acts as an anode where it is oxidised, and a nearby spot on the same surface acts as a cathode where oxygen and water are reduced, so a single piece of iron can behave like a tiny voltaic cell. Dissolved salts and impurities make these local cells work faster, which is why iron rusts more quickly near the sea and on dirty metal than on clean iron in dry air.

Explaining rusting and its prevention, with the conditions correctly stated, is a standard question, and it draws together oxidation, half-equations and the reactivity series in SPM Chemistry.