Question

Rust is hydrated iron (III), $F{e_2}{O_3}.x{H_2}O$ ( reddish brown powder). How is it obtained?

Answer 1

Hint: Try to recall that rust is a reddish-brown iron oxide and is formed by electrochemical corrosion of metals. Also, it is known during rusting both oxidation and reduction occurs at separate locations on the iron. Now by using this you can easily answer the given question.

Complete step by step solution:
It is known to you that rusting is basically deterioration of metal and is simply a Oxidation process.
The special characteristic of the rusting process is that the oxidation and reduction steps occur at separate locations on the surface of iron metal.
This is possible because iron is a conductive metal, so the electrons can flow through it from the anodic to the cathodic regions.
When iron is exposed to oxygen and moisture, it gets oxidized and turns into iron oxide which is also known as “rust”.
The presence of water is necessary in order to transport ions to and from the metal.
Rusting of iron can be regarded as a short-circuited electrochemical cell in which half-cell reactions are as follows:
Cathodic reaction: ${O_2}(g) + 4{H^ + }(aq) + 4{e^ - } \to 2{H_2}O(l)$
Anodic reaction: $Fe(s) \to F{e^{2 + }}(aq) + 2{e^ - }$
Atmospheric oxidation: $F{e^{2 + }}(aq) + 2{H_2}O(l) + \dfrac{1}{2}{O_2}(g) \to F{e_2}{O_3}(s) + 4{H^ + }$.
Hence, from above we can conclude that rust, a reddish-brown powder when iron is exposed to oxygen and any amount of moisture and oxidation of iron into iron oxide takes place.

Note: It should be remembered that during rusting, both cathodic and anodic steps must take place for rusting of iron to occur. Prevention of either one will stop rusting of iron.
Also, you should remember that rusting can be prevented by the process of galvanization, in which sacrificial coating of zinc is done on the surface of iron metal.