Corrosion refers to the loss or conversion into other insoluble compounds of a surface layer of a solid in contact with a fluid. Corrosion can also be defined as the deterioration of metals by an electrochemical process. The economic damage of corrosion causes a lot of tremendous economic damage to building, bridges, ships, cars, equipment and metallurgical plants, river and sea vessel, underground pipelines and other structures. Examples of corrosion are;
• Formation of rust on iron, oxygen gas and water must be present for iron to rust Fe → Fe2+ + 2e- , 4Fe2+ (aq) + O2 (g) + (4+2x) H2O → 2Fe2O3. X H2O + 8H2+(aq)
• Coinage metal such as copper and silver also corrode but much slowly Cu(s) →Cu2+(aq) + 2e-, Ag(s) → Ag+(aq) + e-
• Deterioration of copper and brass
• Tarnish on silver
Aluminium is another metal which corrode. Aluminium has a much greater tendency to oxides than iron does because aluminium has a more negative standard reduction potential than iron, base on this fact, we expect to see air plane slowly corrode away in rainstorms and soda cans transform into piles of corroded aluminium. This process do not occur because the layer of insoluble aluminium oxide Al2O3 that forms on its surface when the metal is exposed to air serves to protect the aluminium underneath from further corrosion.
Electrochemical Methods of Preventing Corrosion
The electrochemical methods of avoiding corrosion are:
• Application of external voltage or current or use of a sacrificial anode to set the voltage of the material in the passive zone (anodic protection) or at a sufficiently negative potential such that the rate of corrosion is slow (cathodic protection).
• Removal of the reducible and aggressive species in solution e.g increase of pH, removal of hydrogen or in humid atmospheres and reduction of humidity.
• Avoidance of bimetallic contacts where this can lead to enhanced corrosion.
• Avoidance of mechanical stress
• Selection of a bulk material with high corrosion resistance
• Coating of metal with a suitable, sufficiently thick and homogenous protective film, example oxide and paint
The fundamental electrochemical processes that occur during formation of rust on iron are;
Oxygen gas and water must be present for iron to rust. A region of the metal surface serves as the anode, where oxidation half reaction occurs; Fe → Fe2+ + 2e-, the electrons given up by ions to reduce atmospheric oxygen to water at the cathode. Half reaction at the cathode is O2(g) +4H+ + 4e- → 2H2O(l), the overall reaction is 2Fe(s) + O2(g) + 4H+(aq) → 2Fe2+(aq) + 2H2O the reaction occurs in acidic medium, the H+ are supplied by the reaction of atmospheric carbon dioxide with water to form carbonic acids.
CO2 + H2O → H2CO3 (carbonic acid)
Fe2+ ions formed at the anode are further oxidised by oxygen
4Fe2+(aq) + O2(g) + (4 + 2x)H2O → 2Fe2O3.xH2O + 8H+(aq)
The chemical formula of the substance referred to as rust is Fe2O3.XH2O
Other various ways of preventing rust metals are:
• Barrier protection: in this method a barrier film is introduced between iron and atmospheric oxygen and moisture. Barrier protection can be achieved by one of the following ways: by painting the surface, by coating the surface with a thin film of oil or grease and by electroplating iron with some non-corrosive metal such as nickel, chromium, copper e.t.c.
• Sacrificial protection: in this method surface of iron is covered with layer of more active like zinc. This active metal loses electrons (undergoes oxidation) in preference to iron and hence, prevents the rusting of iron. Zinc metal is generally used for protecting iron and the process is called Galvanization. Zinc, magnesium and aluminium powder dissolved in paints can also be applied as protective layers.
• Use of anti-rust solutions: the alkaline phosphate and alkaline chromate solutions act as anti-rust solutions. When iron articles are dipped into a boiling and strongly alkaline solution of sodium phosphate, a protective insoluble film of iron phosphate is formed on them. The film protects the article from rusting.
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