The chemistry of rusting
Now look at the process that happens when iron rusts. Iron atoms give up electrons and become positive ions — they are oxidised.
Fe (s) → Fe2+ (aq) + 2e
Metal atoms are oxidised like this only when there is something nearby that can take the electrons released. When oxygen from the atmosphere and water or water vapour are present together, they take the electrons and are reduced.
Fe (s) → Fe2+ (aq) + 2e
2H2O (l) + O2 (g) + 4e → 4OH- (aq)
the half reactions of rusting: (1) and (2)
The number of electrons given out in (1) must balance the number taken in (2), so (1) is multiplied by 2 and added to (2). The Fe2+ and OH- ions formed together make iron(II) hydroxide.
2Fe (s) → 2Fe2+ (aq) + 4e
2H2O (l) + O2 (g) + 2Fe (s) → 2Fe(OH)2
(1) × 2 = (3); (2) + (3)
So rusting too is an electrochemical process, like the ones studied in 12.1. Reaction (1) is the anode reaction, because it is an oxidation; reaction (2) is the cathode reaction, because it is a reduction. The Fe(OH)2 formed reacts further with air to form hydrated ferric oxide (Fe2O3.H2O).
4Fe(OH)2 (s) + O2 (g) → 2(Fe2O3.H2O) (s) + 2H2O (l)
Rust
The hydrated ferric oxide formed — rust — is reddish-brown. Because the number of water molecules bound to the ferric oxide can vary, the chemical formula of rust is more properly written Fe2O3.xH2O.
Rust is not the same as iron oxide
Rust holds water in its formula. That is why it is flaky and crumbles away, exposing fresh iron underneath to rust in turn — unlike the tight oxide layer on aluminium, which protects the metal.
Rusting is a tiny electrochemical cell
Iron is oxidised at anode areas, oxygen and water are reduced at cathode areas, and the products end up as rust, Fe2O3.xH2O.