What rusting is
Rusting is the corrosion of iron. It is a slow redox reaction that needs both oxygen and water present. Rust is hydrated iron(III) oxide, Fe2O3·xH2O, a flaky brown solid. Because iron is oxidised, rusting is an electrochemical process: one part of the metal acts as the anode and another as the cathode.
Key idea
Anode (oxidation): Fe → Fe2+ + 2e−
Cathode (reduction): O2 + 2H2O + 4e− → 4OH−
The Fe2+ is then further oxidised by oxygen to Fe3+, forming rust.
Preventing rust
Rust can be prevented by keeping out oxygen and water (painting, oiling, plastic coating) or by sacrificial protection. In sacrificial protection a more reactive metal such as zinc or magnesium is attached to the iron; the more reactive metal loses electrons in preference, so the iron is protected even if the coating is scratched.
Worked example
Galvanised iron is coated with zinc. Because zinc is more electropositive than iron, zinc becomes the anode: Zn → Zn2+ + 2e−. The iron becomes the cathode and does not rust. Attaching blocks of magnesium to ship hulls works the same way.
A classic experiment uses three test tubes: iron nails in ordinary water with air present do rust; nails in water that has been boiled and sealed under a layer of oil do not rust (no oxygen); and nails in a tube containing a drying agent do not rust (no water). This confirms that both oxygen and water are needed. Alloying iron with chromium and nickel produces stainless steel, which resists rusting because a thin, tough oxide layer forms on its surface.
Remember
- Rusting needs oxygen AND water together.
- Salt and acid speed up rusting (better electrolyte).
- A more reactive metal protects iron (sacrificial protection).
- Tin plating fails if scratched, because tin is less reactive than iron.