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Corrosion of Metals

CORROSION OF METALS

Metals are frequently used in almost all walks of our present day life. Bridges, railways, buildings, vehicles, industries, household articles, all involve the use of metals. But we commonly observe that certain metals (except those which are least reactive like Au, Pt, Pd etc.) are slowly eaten up on long exposure to atmosphere. For example:

• silver gets tarnished,

• copper develops green coating on its surface,

• iron rusts, and

• aluminium, zinc, lead loses its lustre.

In fact, such metals react with the gases or moisture present in the environment to form undesirable compounds. This process in general is referred to as corrosion of metals. Corrosion may, thus, be defined as the process of slow conversion of metals into their undesirable compounds (usually oxides) by reaction with moisture and other gases present in the atmosphere.

 

CORROSION, AN OXIDATION-REDUCTION PROCESS

Corrosion of metals is basically an oxidation-reduction process. Most metals undergo slow oxidation with atmospheric oxygen and moisture to form oxide layer on their surface. In certain metals like iron this oxide .layer does not stick to the surface and gradually falls off exposing fresh metal surface to further corrosion: In this way, the metal is slowly eaten away weakening the structure in which it has been used.

certain metals like aluminium, copper, chromium zinc, lead react with atmospheric gases and moisture to form hydroxide, oxide or trioxocarbonate(IV) layers which generally stick to their surface and their surface get tarnished. This layer, however, becomes a sort of protective layer which prevent further corrosion. But at higher temperatures even this layer can also crack exposing fresh surface for further corrosion.

Iron is generally used in buildings and other huge structures and rusting of iron (corrosion) cause damage to these structures. Sometimes this damage is so severe that it leads to the collapse of structure. In this unit, we shall study mechanism of corrosion of iron and the method used for its prevention.

 

FACTORS WHICH AFFECT CORROSION

The factors which affect the rate of corrosion are:

(i)                 Reactivity of the metal. The more active metals are more prone to corrosion.

(ii)               Presence of impurities. Presence of impurities helps in setting up a corrosion cell and makes the corrosion to occur rapidly. For example, pure iron does not rust.

(iii)             Air and moisture. Air and moisture are quite helpful in corrosion. The presence of gases like CO2 and SO2 in air makes it still rapid. For example, no rusting is caused if iron is kept in vacuum.

(iv)             Strains in metal. Strains in metal also help in corrosion. For example, in iron articles, rusting is more pronounced on the areas having bends, dents, scratches, nicks and cuts.

(v)               Presence of electrolytes. The presence of electrolytes also makes the corrosion process faster. For example, iron rusts more rapidly in saline water in comparison to pure water.

Let us understand the mechanism of corrosion by studying the most familiar example of rusting of iron. Chemically rust is hydrated iron (III)’ oxide, Fe2O3 . XH2O. It is generally caused by moisture, CO2, O2 of air. Rust is a non-sticking brown-coloured material which can be easily removed by scratching. There are a number of theories about the mechanism of rusting. The most widely accepted theory is electrochemical theory which is being discussed here.

 

ELECTROCHEMICAL THEORY OF RUSTING

According to this theory, the impure iron surface behaves like a small electrochemical cell in presence of water containing dissolved oxygen or carbon dioxide. Such a cell is also called corrosion cell or corrosion couple. In these miniature corrosion cells, pure iron acts as anode and impure surfaces act as cathode. Moisture having dissolved oxygen or carbon dioxide in it constitutes electrolytic solution. At anode, oxidation of Fe atoms takes place. Thus, Fe atoms pass into solution as Fe2+ ions leaving behind electrons in the metal which are pushed into cathodic area.

At cathode, the electrons are picked up by the H+ ions which are produced either from H2O or from H2CO3 (formed by dissolution of CH2 in moisture).

CO2 + H2O à H+ + HCO 3

2H+ (aq) + 2e à 2H

The H atoms, thus, formed, reduce the dissolved – oxygen as

2H+ + 1/2  O2 à H2O

The net reduction process at the cathodic area is

The net reaction of the corrosion cell can be obtained adding equations (i) and (ii).

Fe + 2H+ + 1/2  O2 à Fe2+ + H2O ;

The iron (II) ions so formed move through water come at the surface of iron object where these are further oxidised to iron (III) state by atmospheric oxygen constitute rust which is hydrated iron (III) oxide.

2Fe2+ + 1/2 O2 + 2H2O à Fe2O3 + 4H+

Fe2O3 + xH2O àFe2O3 xH2O