Atmospheric Corrosion

Atmospheric corrosion can be defined as the corrosion of materials exposed to air and its pollutants, rather than immersed in a liquid. It can further be classified into dry, damp, and wet categories. If the metal is exposed to rain, it may corrode while it is wet at the rate appropriate to immersion in impure, well-aerated water, but the rate will fall when it dries.

A fundamental requirement for electrochemical corrosion processes is the presence of an electrolyte. Thin-film "invisible" electrolytes tend to form on metallic surfaces under atmospheric exposure conditions after a certain critical humidity level is reached. The critical humidity of iron is 60 percent in an atmosphere free of sulphur dioxide. The anodic oxidation reaction involves the dissolution of the metal, while the cathodic reaction is often assumed to be the oxygen reduction reaction.

For iron these reaction are as following:

Anodic reactions involve oxidation of metal to its ions, e.g. for iron this reaction occurs.

Anode Reaction: 2Fe = 2Fe^{+ +}+ 4e (1)

The cathodic process involves reduction and several reactions are possible. In acidic water, where hydrogen ions (H⁺) are plentiful, the following reaction occurs.

2H⁺ + 2e = H₂ (2)

In alkaline solutions, where hydrogen ions are rare, the reduction of water will occur to yield alkali and hydrogen.

2H₂O + 2e = H₂ + 2OH^- (3)

However, unless the water is deaerated reduction of oxygen is the most likely process, again producing alkali at the surface of the metal.

O₂ + 2H₂O + 4e = 4OH^- (4)

Reactions (1) and (2) are shown schematically in Fig 1 where anodic and cathodic sites are nearby on the surface of a piece of metal. It is possible to change the rate of these two reactions by withdrawing electrons or supplying additional electrons to the piece of metal. It is an established principle that if a change occurs in one of the factors under which a system is in equilibrium, the system will tend to adjust itself so as to annul, as far as possible, the effect of that change. Thus, if it withdraws electrons from the piece of metal the rate of reaction (1) will increase to attempt to offset its action and the dissolution of iron will increase, whereas reaction (2) will decrease. Conversely, if we supply additional electrons from an external source to the piece of metal, reaction (1) will decrease to give reduced corrosion and reaction (2) will increase. The latter case will apply to cathodic protection. Thus, to prevent corrosion we have to continue to supply electrons to the steel from an external source to satisfy the requirements of the cathodic reaction.

Relatively complex reaction sequences have been proposed for the corrosion product formation and breakdown processes to explain observed atmospheric corrosion rates for different classes of metals