Corrosion as a Redox Reaction

  1. Corrosion of metal = metal loses its electrons to form positive ions.
    Example
    Corrosion of iron
    Fe → Fe2+ + 2e

    Corrosion of Magnesium
    Mg → Mg2+ + 2e

  2. The higher the position of a metal in the electrochemical series, the more electropositive (reactive) the metal is. The metal has a greater tendency to give away electrons to form the metal ion, that is the metal is more easily corroded. 
  3. For instance, the metal magnesium corrodes more easily than copper because magnesium is more electropositive than copper.
  4. Corrosion of metal is a redox reaction as the metal loses electrons to oxygen and water, which act as the oxidising agents to receive the electrons.
  5. Corrosion of iron is also called rusting.
    Fe (s) → Fe2+ (aq) + 2e
  6. Rusting of iron can only occur if both oxygen and water are present.


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

Electrochemical Corrosion

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Controlling Rusting

Steps to Control Rusting

The use of a metal which is less electropositive

  1. The plating of iron with a thin layer of a metal which is less electropositive such as tin, silver or copper will prevent the iron underneath it to react with water and air, and so prevents the iron from rusting.
  2. However, the rusting of iron will occur faster if the protective layer is scratched. This is because iron is more electropositive than tin, silver or copper. The plating of iron by tin is used a lot in the making of tinned food.

The use of a metal which is more electropositive

  1. Metals which are more electropositive are used as a sacrificial metal to prevent corrosion on metals which are less electropositive. The metal which is more electropositive corrodes and acts as the anode.
  2. The less electropositive acts as the cathode and is protected from corroding. This method is known as the cathode protection or electrochemical protection.
Cover by paint, oil and grease
  1. A layer of paint, oil/grease, or plastic that is used to cover the surface of iron from contact with air and water in the atmosphere. So, rusting can be avoided. For example,
    1. Protection by a layer of paint usually is used for iron and steel objects like cars, ships, bridges and other things which do not undergo friction during use.
    2. Protection by a layer of oil/grease is used for part of machinery that move.
    3. Protection by plastic is used for daily items at home like clothes hangers and fencing.
    4. Alloying of iron can prevent rusting. For example, when iron is alloyed with chromium and nickel to form stainless steel, the layer of chromium oxide that is hard, strong and difficult to crack on the surface of the iron alloy prevents the iron from rusting.