Explain the rusting of iron with the help of the electrochemical theory of corrosion and a diagram.

Electrochemical (wet) theory: when a metal is in contact with moisture or an electrolyte, many tiny galvanic cells form on its surface. The anodic areas dissolve (oxidation), and the electrons flow through the metal to the cathodic areas, where they are consumed (reduction). The ions migrate through the electrolyte, and the corrosion product forms between the anode and cathode.

Rusting of iron (in neutral, aerated water – oxygen absorption mechanism):

At the anode: \( \mathrm{Fe \rightarrow Fe^{2+} + 2e^-} \)

At the cathode: \( \mathrm{\tfrac12 O_2 + H_2O + 2e^- \rightarrow 2OH^-} \)

Net: \( \mathrm{Fe^{2+} + 2OH^- \rightarrow Fe(OH)_2} \)

With excess oxygen: \( \mathrm{4Fe(OH)_2 + O_2 + 2H_2O \rightarrow 4Fe(OH)_3} \), i.e. \( \mathrm{Fe_2O_3\cdot xH_2O} \) (yellowish-brown rust).

With limited oxygen: \( \mathrm{3Fe(OH)_2 + \tfrac12 O_2 \rightarrow Fe_3O_4\cdot 3H_2O} \) (black magnetite).

                 O2 (air)
                   |
      water drop / electrolyte containing dissolved O2
   ________________________________________________
  /   cathode       rust           anode       cathode \
 /  O2+H2O+2e-->2OH-  Fe(OH)3   Fe -> Fe2+ + 2e-        \
=================== IRON ================================
          e-  <-------  e-  ------->  e-   (flow through metal)
 (well-aerated edge = cathode; less-aerated centre = anode)

Rust usually forms between the anode and cathode, where Fe²⁺ and OH⁻ meet. The less-aerated part (the centre of a drop, or under dirt) becomes the anode and corrodes (differential aeration).

In acidic medium (hydrogen evolution mechanism): anode: Fe → Fe²⁺ + 2e⁻; cathode: 2H⁺ + 2e⁻ → H₂↑.