Home arrow Knowledgebase arrow Rusting of iron
Rusting of iron

Explanation of the Rusting of Iron

When iron rusts a spontaneous redox reaction occurs, between the oxygen and iron. If water is added the rusting occurs more rapidly.


Iron (s) + Oxygen(g)----------> Iron (III) oxide(s)

Image of Rusting of iron

When water is added to iron, an electrochemical cell is created that has a distinct anode and cathode. If an iron nail is placed in agar or gel in which ferric cyanide ions and phenolphthalein indicator have been placed, the ends of the nail turn blue and the middle of the nail turns red. The blue colour is caused by ferricyanide indicator reaction with the iron ions, and the red colour (pink) is due to reaction between hydroxide ions and phenolphthalein. How are these ions produced.

At one spot on the nail (the Anodic site of our electrochemical cell) Iron loses electrons (is oxidized) to form iron (II) ions.
Fe (s) ------> Fe2+ (aq) + 2e-

At another spot on the nail the oxygen in the air combines with water and forms hydroxide ions.
1/2 O2 (g) + H2O (l) + 2e- --------> 2OH- (aq)

In the presence of oxygen the iron further oxidizes at the anode (loses electrons) to become iron (III) ions.
Fe 2+ (aq) ------> Fe3+ (aq) + e-

The iron (III) ions and the hydroxide combine to form rust ( flaky brown substance) .
2 Fe 3+ (aq) + 6OH- (aq) -----> Fe2O3 (s) + 3 H2O (l)

Notice that water is required for the reaction at the cathode but produced in the overall reaction. It therefore is acting like a homogeneous catalyst, to speed up the rusting of iron. In essence the water and oxygen make it easier for iron to rust.

As with an electrochemical cells the electrons flow from the anode to the cathode. Oxygen and water are both need to speed the rusting process in metals.