When metals are exposed to atmospheric conditions, they react with air or water in the environment to form undesirable compounds (usually oxides). This process is called corrosion.
The least active metals such as gold, platinum and palladium are attacked by environment.
For example: Silver tarnishes, copper develops a green coating, lead or stainless steel lose their lustre due to corrosion.
Corrosion is a process of deterioration of a metal as a result of its reaction with air or water (environment) surrounding it.
In case of iron, corrosion is called rusting. Chemically rust is hydrated form of ferric oxide, Fe2O3 · xH2O. Rusting of iron is generally caused by moisture, carbon dioxide and oxygen present in air.
Rusting takes place only when iron is in contact with moist air.
Factors which affect Corrosion
The main factors which affect corrosion are:
(i) Position of metals in e.m.f. series: The reactivity of metal depends upon its position in the electrochemical series. More the reactivity of metal, the more will be the possibility of the metal getting corroded.
(ii) Presence of impurities in metals: The impurities help in setting up voltaic cells, which increase the speed of corrosion.
(iii) Presence of electrolytes: Presence of electrolytes in water also increases the rate of corrosion.
(iv) Presence of CO2 in water: Presence of CO2 in natural water increases rusting of iron. Water containing CO2 acts as an electrolyte and increases the flow of electrons from one place to another.
(v) Presence of protective coatings: When the iron surface is coated with layers of metals more active than iron, then the rate of corrosion is retarded.
Mechanism of Rusting of Iron
Non-uniform surface of metal or impurities present in iron behave like small electric cells in the presence of water containing dissolved oxygen or carbon dioxide. A film of moisture with dissolved CO2, constitutes electrolytic solution covering the metal surface at various places.
In the small electrolytic cells, pure iron acts as anode while cathodes are impure portions.
The overall rusting involves the following steps:
Oxidation occurs at the anodes of each electrochemical cell. Therefore, at each anode neutral iron atoms are oxidised to ferrous ions.
Fe(s) ——-> Fe2+ (aq) +2e¯
The metal atoms in the lattice pass into the solution as ions, leaving electrons on the metal itself. These electrons move towards the cathode region through the metal.
At the cathodes of each cell, the electrons are taken up by hydrogen ions (reduction takes place). The H+ ions are obtained either from water or from acidic substances (e.g CO2) in water:
H2O ——>H+ + OH¯
CO2 + H2O ——-> H+ + HCO3¯
H+ + e¯ ———-> H
The hydrogen atoms on the iron surface reduce dissolved oxygen.
4H++O2 ———> 2H2O
Therefore, the overall reaction at cathode of different electrochemical cells may be written as:
4 H+ + O2 +4e¯ ——>2 H2O
Oxidation half reaction
Fe(s) ——-> Fe2+ (aq) + 2e¯] x 2 (EΦ= -0.44 V)
Reduction half reaction
4H+ + O2 +4e¯ ——> 2H2O (EΦ=1.23 V)
Overall cell reaction
2Fe(s) +4H+ +O2 ——–>2Fe2+ (aq) + 2H2O (EΦ =1.67 V)
The ferrous ions are oxidised further by atmospheric oxygen to form rust.
4Fe2+(aq)+O2 (g) + 4H2O ——> 2Fe2O3 + 8H+
Fe2O3 + x H2O ——> Fe2O3·xH2O
Salt water accelerates corrosion. This is mainly due to the fact that salt water increases the electrical conduction of electrolyte solution formed on the metal surface. Therefore, rusting becomes more serious problem where salt water is present.
Prevention of Corrosion
This can be prevented or retarded by the methods given below:
1) Barrier protection
In this method a barrier is placed between iron and atmospheric air. The barrier protection can be achieved by any of the following methods :
(i) the surface is coated with paint or some chemicals.
(ii) the surface is protected by applying a thin film of oil or grease.
(iii) the metal is electroplated with metals like tin, nickel, zinc, chromium, aluminium, etc.
Many vehicles such as cycles, motors, cars made from iron sheets are protected from rusting by paints. Many iron articles are electroplated with coating of other metals.
2) Sacrificial Protection
In this method iron is protected from rusting by covering it with a layer of a metal more active than iron. The active metal loses electrons in preference to iron and goes into ionic state. Therefore, the covering metal is consumed with time, but long as it is present on the surface of iron, the latter is not rusted. This type of process in which rusting of iron is protected is called sacrificial protection.
Zinc is commonly used for covering iron surfaces. The process of covering iron with zinc is called galvanization. The galvanized iron materials maintain their lustre due to the coating of invisible layer of basic zinc carbonate, ZnCO3.Zn(OH)2, on the zinc film.
If some scratches occur on the protective zinc film on coated iron, even then iron will not be rusted. This is due to the fact that because of scratches, both zinc and iron get exposed to oxidation but zinc undergoes oxidation in preference to iron. This is so because the reduction potential of zinc is less than the reduction potential of iron.
Zn2+ (aq) +2e¯ ⇔ Zn(s) EΦ= – 0.76 V
Fe2+ +2e¯ ⇔ Fe(s) EΦ= -0.44 V
Fe2+ +2e ⇔ Fe(s) EΦ = -0.44 V
3) Electrical protection
Fe2+ (aq) + 2e¯ ———>Fe(s) EΦ= -0.44 V