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.
- 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
- 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
- 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.
- 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.
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