Corrosion

Corrosion

We are familiar with the rusting of iron. Have you ever noticed a green film formed on copper and brass vessels?. In both, the metal is oxidised by oxygen in presence of moisture. This redox process which causes the deterioration of metal is called corrosion. As the corrosion of iron causes damages to our buildings, bridges etc....it is important to know the chemistry of rusting and how to prevent it. Rusting of iron is an electrochemical process.

Electrochemical mechanism of corrosion

The formation of rust requires both oxygen and water. Since it is an electrochemical redox process, it requires an anode and cathode in different places on the surface of iron. The iron surface and a droplet of water on the surface as shown in figure (9.15) form a tiny galvanic cell. The region enclosed by water is exposed to low amount of oxygen and it acts as the anode. The remaining area has high amount of oxygen and it acts as cathode. So based on the oxygen content, an electro chemical cell is formed. corrosion occurs at the anode i,e,. in the region enclosed by the water as discussed below.

At anode (oxidation): Iron dissolves in the anode region

2Fe(s) → 2Fe²⁺ (aq) + 4e^-                           E = 0.44V.

The electrons move through the iron metal from the anode to the cathode area where the oxygen dissolved in water, is reduced to water.

At Cathode (reduction)

The reaction of atmospheric carbon dioxide with water gives carbonic acid which furnishes the H⁺ ions for reduction.

O₂ (g) + 4H⁺ (aq) + 4e^- → 2H₂O (l )                        E = 1.23V

The electrical circuit is completed by the migration of ions through water droplet.

The overall redox reactions is,

2Fe(s) + O₂(g) + 4H⁺ (aq) → 2Fe²⁺ (aq) + 2H₂O(l )           E = 0.444 +1.23 = 1.67V

The positive emf value indicates that the reaction is spontaneous.

Fe²⁺ ions are further oxidised to Fe³⁺ , which on further reaction with oxygen to form rust.

4Fe²⁺ (aq)+O₂(g)+4H⁺ (aq) → 4Fe³⁺ (aq)+2H₂O(l)

2Fe³⁺ (aq)+4H₂O(l) → Fe₂O₃.H₂O(s) + 6H⁺ (aq)

Other metals such as aluminium, copper and silver also undergo corrosion, but at a slower rate than iron. For example, let us consider the reduction of aluminium,

Al(s) → Al³⁺ (aq)+3e^−

Al³⁺ , which reacts with oxygen in air to forms a protective coating of Al₂O₃ . This coating act as a protective film for the inner surface. So,further corrosion is prevented.

Protection of metals form corrosion

This can be achieved by the following methods.

i. Coating metal surface by paint.

ii. Galvanizing - by coating with another metal such as zinc. zinc is stronger reducing agent than iron and hence it can be more easily corroded than iron. i.e., instead of iron, the zinc is oxidised.

iii. Cathodic protection - In this technique, unlike galvanising the entire surface of the metal to be protected need not be covered with a protecting metal. Instead, metals such as Mg or zinc which is corroded more easily than iron can be used as a sacrificial anode and the iron material acts as a cathode. So iron is protected, but Mg or Zn is corroded.

Passivation - The metal is treated with strong oxidising agents such as concentrated HNO₃. As a result, a protective oxide layer is formed on the surface of metal.

Alloy formation - The oxidising tendency of iron can be reduced by forming its alloy with other more anodic metals.

Example, stainless steel - an alloy of Fe and Cr .