What is corrosion? Explain the rusting of iron as an electrochemical process, and give three methods of preventing it.

1. 1 · Define. Corrosion is the slow coating of a metal with its oxide or other salts by the action of air, water and other substances. Chemically it is oxidation of the metal.
2. 2 · Set up the cell. A water drop with dissolved CO₂ (H₂CO₃ supplies H⁺) is the electrolyte. One spot of the iron is the anode and another is the cathode.
3. 3 · Write the reactions.
   • Anode: 2Fe(s) → 2Fe²⁺(aq) + 4e⁻ (E° = −0.44 V)
   • Cathode: O₂(g) + 4H⁺(aq) + 4e⁻ → 2H₂O(l) (E° = +1.23 V)
   • Overall: 2Fe + O₂ + 4H⁺ → 2Fe²⁺ + 2H₂O; E°cell = 1.67 V
4. 4 · Show rust forming. Air oxidises Fe²⁺ to Fe³⁺, which forms rust, Fe₂O₃·xH₂O.
5. 5 · Draw. Iron strip with a water drop; label the anode spot, the cathode spot, electron flow through the iron and the rust.
6. 6 · Give three methods of prevention.
   • Barrier protection: paint, oil, grease or chemicals such as bisphenol
   • Covering with another metal: galvanising with zinc, or tin coating
   • Sacrificial electrode: Mg or Zn joined to the iron corrodes instead
7. 7 · Link to real life. Galvanised buckets and roof sheets, and Mg blocks on buried pipelines, use these ideas.
8. 8 · Close in one line. So rusting is a natural galvanic cell, and preventing it means breaking that cell.

Scoring tip: Write the E° value beside each half-reaction and work out 1.67 V. It shows the examiner why rusting happens on its own.