Rusting of Iron

  1. In the rusting of iron, iron acts as the reducing agent and oxygen gas as the oxidising agent.
  2. The process or rusting of iron can be explained by using figure below.
  3. When the surface of the iron is exposed to water droplets, the centre of the water droplets undergoes the process of oxidation and is known as the anode.
  4. The edge of the water droplets undergoes a process of reduction and is known as the cathode. (The edge of the water droplets acts as the cathode because of the concentration of soluble oxygen is higher on the edge of the water droplets than in the centre.)
  5. At the anode, the metal iron undergoes oxidation to form the iron(II) ion with the loss of electrons.
    Fe → Fe2+ + 2e
  6. Electrons that are free at the anode flow through the metal iron to the cathode area where soluble oxygen in the water accepts electrons to form hydroxide ions.
    O2 + 2H2O + 4e → 4OH-
  7. The iron(II) ions are then combines with the hydroxide ion to form iron(II) hydroxide.
    Fe2+ + OH-  Fe(OH)2
  8. Iron(II) hydroxide is then oxidised by oxygen to form iron(III) hydroxide.
    4Fe(OH)2 + 2H2O + O2 → 4Fe(OH)3
  9. The iron(III) hydroxide is then decomposed to form hydrated iron(III) oxide, Fe2O3xH2O by oxygen in the air.
    4Fe(OH)3  Fe2O3xH2O
  10. The hydrated iron(III) oxide is brown in colour and is known as rust.
  11. The overall equation for the rusting of iron is
    4Fe + 3O2 +2xH2O → 2Fe2O3xH2O

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