The chemistry of corrosion

Posted on

The rusting of iron can be summarised by the equation

4 Fe(s) + 3 O2(g) + H2O(l)  ⇾  2 Fe2O3•xH2O(s)

In reality, the process takes place via a series of half reactions that take place in the iron itself and the water droplet.

  1. Iron metal is oxidised to iron(II) ions in the centre of the water droplet, where the dissolved oxygen concentration is lower.  The Fe2+ ions pass into solution.

Fe(s) ⇾ Fe2+(aq) + 2e–      E⦵ = – 0.44V

  1. In regions of the water droplet with a higher concentration of dissolved oxygen (the edges), oxygen is reduced to hydroxide ions by electrons that flow from the oxidation of iron through the metal. 

O2(g) + 2H2O(l) + 4e– ⇾ 4OH–(aq)       E⦵ = +0.40V

  1. Fe2+ and OH– ions react to form Fe(OH)2, a green precipitate, away from the surface of the iron so there is no protective oxide layer formed. 
  1. Fe2+(aq) ions are oxidised to Fe3+(aq), which then react with OH– ions and water to form the insoluble hydrated oxide known as rust. 

Sea water has a greater conductivity because of the high concentration of dissolved ions and so iron rusts faster in the sea than in rain water. 

Preventing corrosion

  • oil / paint / grease will all help prevent water and oxygen from reaching the surface of the iron
  • galvanising is a process in which the iron surface is coated in zinc which forms an impermeable oxide layer

Zn2+(aq) + 2e– ⇌ Zn(s) E⦵ = – 0.76V

Fe2+(aq) + 2e– ⇌ Fe(s) E⦵ = – 0.44V

O2(g) + 2H2O(l) + 2e– ⇌ 4OH–(aq) E⦵= + 0.40V

If scratched, the iron (Fe) is oxidised to Fe2+ by O2(g) / H2O(l) / OH–(aq) half cell which has the more positive E⦵. The Fe2+ is then reduced back to Fe by the Zn2+(aq) / Zn(s) half cell which has the less positive E⦵ of the two and so forms an oxidation half cell.

  • cathodic protection is where the iron is attached by a wire to a block of a more easily oxidised metal such as magnesium

Mg2+(aq) + 2e– ⇌ Mg(s) E⦵ = – 2.37V

Fe2+(aq) + 2e– ⇌ Fe(s) E⦵ = – 0.44V

The magnesium is a sacrificial anode, reacting with in preference to iron with O2(g) / H2O(l), and it needs to be replaced periodically.

Practice question

A rusted iron bar found on a beach was covered in pits containing sea water and a greenish precipitate that had turned orange at the surface. 

(a) Give the formulae of the green and orange precipitates.

(b) Give the half equations for the two redox reactions that take place when iron rusts.

(c) Explain why the iron bar rusts faster in sea water than if exposed only to rain water.

    Answer

    (a)  Green precipitate is Fe(OH)2 and orange precipitate is Fe2O3•xH2O

    (b)  Fe(s) ⇾  Fe2+(aq) + 2e- / O2(g) + 2H2O(l) + 4e-  ⇾  4OH-(aq)

    (c)  Seawater contains a greater concentration (not more!) of dissolved ions than rainwater so it is a better conductor.