Redox, Rusting and Iron

Redox Reactions

Redox reactions involve both oxidation and reduction reactions happening simultaneously.
An oxidation reaction involves the loss of electrons.
A reduction reaction involves the gain of electrons.
Redox reactions are seen in many everyday processes such as rusting and photosynthesis.

Oxidation States

Oxidation state refers to the number of electrons that may be gained, lost, or shared by an atom during a chemical reaction.
This helps to understand and determine whether a substance has been oxidised or reduced in a reaction.
In the context of rusting, iron has an oxidation state of 0 in its metallic form, but when rusted (becomes iron(III) oxide), the iron has an oxidation state of +3.

Rusting and Iron

Rusting is an example of a redox reaction, specifically it is an oxidation reaction.
Rusting primarily involves iron, water, and oxygen, forming hydrated iron(III) oxide (rust).
The process can be represented as: 4 Fe(s) + 3 O2(g) + 6 H2O(l) -> 4 Fe(OH)3(s)
Rusting of iron is a slow reaction under normal circumstances, but factors such as temperature, surface area and presence of electrolytes can speed up the reaction.
The damage due to rusting is a significant problem in industries, particularly those involving structures made from iron and steel.

Preventing Rusting

There are various ways to prevent rusting of iron, such as oiling, painting, galvanising, and using a sacrificial metal.
In a sacrificial method, a more reactive metal (like zinc or magnesium) is attached to the iron. These metals react with oxygen before iron, preventing it from rusting.
Galvanising involves coating iron with a layer of zinc, which also acts as a sacrificial layer.
Rusting can also be prevented by designing structures that allow water to drain away, reducing the contact of iron with water.