Economic Importance Of Rusting Of Iron

• Heat:- An increase in temperature always increases rate of chemical reactions, including rusting. 

• Humidity:- Corrosion starts when the relative humidity of the air

exceeds around 65%. Many areas has a higher humidity in winter (80-95%) than in summer (60-80%). In consequence, iron rusts five times faster in winter as it does in summer. "Many water have lime and carbonic acid in equilibrium. This is called equilibrium water, where there is sufficient carbon dioxide in solution to stabilize the carbonate. 

• Contact with a less reactive metal:- If iron and copper plates

are joined together and put in water containing dissolved oxygen, iron loses electrons more readily than copper. Hence, iron forms the anode and copper the cathode of an electrochemical cell. In this case, iron rusts even more quickly than when there was no copper.  Other factors that speed up rusting include the presence of sharply pointed regions in the iron piece, or a high concentration of dissolved oxygen in water. 

2.2 PROTECTION FROM RUSTING:- Iron can be protected from rusting by applying a protective layer. Since air and water are necessary for rusting to occur, so any method which can keep out one or both of them from iron will prevent rusting.

The protective methods are as follows:- 

• Coating with paint, oil or grease:- Objects that are unlikely to

be scratched can be coated with paint (or lacquer, or enamel). For example, bridges, ships and car bodies are painted. Moving parts of a machine are protected by applying oil or grease. 

• Coating with another metal:- Iron can be coated with a thin

layer of another metal which is resistant to corrosion. Galvanized iron is iron coated with zinc e.g roofs, buckets and dustbins are made from galvanized iron. 3) Using Alloys Of 

Iron can be protected from rusting by coating it with another metal that does not corrode so easily. Galvanizing (zinc-plating) provides a good example of cathodic protection. Zinc is a more reactive metal, it loses electron in preference to iron, thereby preventing the formation of iron (ii) ions (Fe2+ (aq) ions). This is because, zinc is above iron in the electrochemical series, it serve as the anode. When zinc layer is broken, the zinc dissolves instead of the iron, which being the cathode, and remains intact. When the zinc layer is undamaged, the iron is covered up and is protected from rusting. In case the coating is partly damaged, the iron, though exposed, is still protected. Zinc, being more reactive than iron, will form zinc ions

At anode (zinc):- Zn (s)                            Zn2+ (aq) + 2e- 

The electrons flow from zinc (anode) to iron (cathode), where the following half-reaction take place: 

At cathode:- O2 (g) + 2H2O (l) + 4e-                4 OH- (aq) 

Since iron is forced to accept the electrons, it cannot corrode, as corrosion involves giving off electrons. Rusting of iron will only set in after the entire zinc layer has corroded. Thus, zinc form a more efficient protective covering for iron.

 Cathodic protection is also used to prevent rusting in underground iron pipelines. Bags containing magnesium turnings are connected to the buried pipelines at intervals. The magnesium corrodes instead of the iron. The magnesium should therefore be replaced from time to time. 

• Electrical Protection:- Sometimes rusting can be prevented by

using electricity. For example, the negative terminal of the car battery is always connected to the car body. This supplies electrons to the iron body, preventing it from losing electrons. In some piers, the steel structures are protected electrically by connecting them to the negative terminal of a d.c. source.