Explain how rusting of iron is envisaged as setting up of an electrochemical cell.
Rusting of iron is considered as an electrochemical phenomenon occurring in presence of water and air. At a particular spot on iron, oxidation takes place and that spot behaves as anode: 2 Fe (s) → 2 Fe²⁺ + 4 e⁻ (E° = -0.44 V).
Electrons released at anodic spot move through the metal to another spot behaving as cathode where oxygen is reduced in presence of H⁺: O₂(g) + 4 H⁺(aq) + 4 e⁻ → 2 H₂O(l) (E° = 1.23 V).
The overall reaction: 2Fe(s) + O₂(g) + 4H⁺(aq) → 2Fe²⁺(aq) + 2 H₂O(l) (E°cell = 1.67 V). Ferrous ions are further oxidised by atmospheric oxygen to ferric ions forming rust (hydrated ferric oxide, Fe₂O₃·xH₂O).
Explanation
The textbook explains rusting as an electrochemical process where different spots on iron surface act as anode and cathode. The anodic spot undergoes oxidation releasing electrons, while the cathodic spot reduces oxygen. The H⁺ ions come from H₂CO₃ formed by dissolution of CO₂ from air into water. The overall cell potential of 1.67 V indicates spontaneity of the process.