Mechanism of Corrosion

Various theories had been advanced to explain corrosion from time to time. The modern view, known as the electrochemical theory of corrosion, appears to be more sound. This theory is described below by taking the example of rusting of iron.

Water containing oxygen a and carbon dioxide acts as an electrolyte, and helps in the flow of electrons.

The formation of rust on the surface of iron occurs through the following steps.

(i) At the anodic region. Iron in contact with water forms anode and gets oxidized to Fe2+.

Anodic region: Fe(s) →Fe2+ (aq) + 2e–           E0 Fe2+ /Fe  = –0.44 V

The released electrons move to another portion of the iron sheet. This portion of the iron sheet serves as cathode.

(ii)At the cathodic region. At this cathodic portion of the surface, oxygen in the presence of H+ ions (produced due to the ionization of water molecules) get reduced to form H2O.

Cathodic region:

O2(g) + 4H+(aq) + 4e– →2H2O(l) EoO2/H+ ,H2O  = + 1.23 V

The overall reaction of the local cell is the sum of the cathodic and anodic reactions.

Overall cell reaction: 2Fe(s) + O2(g) + 4H+(aq) →2Fe2+(aq) + 2H2O(l) E0 cell = 1.67 V

Fe2+ ions move through the water on the surface of the iron sheet. Presence of electrolytes (e.g., sodium chloride, SO2, CO2 etc.) in water helps in carrying more current through the local cell on the surface of iron. The increased flow of current enhances the rate of corrosion. The Fe2+ ions are further oxidized by atmospheric oxygen to Fe3+, and form hydrated iron(III) oxide, expressed as Fe2O3.xH2O. The hydrated ferric oxide is called rust.

4Fe2+ (aq) + O2(g) + 4H2O(l)  → 2Fe2O3(s) + 8H+
Fe2O3(s) + xH2O (l) → Fe2O3 .xH2Orust

H+ ions produced in the above reaction help further in the rusting of iron. Impurities present in iron also enhance rusting by setting a number of localized cells. Pure iron does not rust.

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