Redox Reactions

• Oxidation can be described as a reaction in which hydrogen is lost or oxygen is gained.
• Reduction can be defined as a reaction in which hydrogen is gained or oxygen is lost.

For example, CuO_(s) + H_2(g) Cu_(s) + H₂O_(l)

• In the above reaction copper oxide loses oxygen, so it has been reduced. Hydrogen gains oxygen so it has been oxidised.
• Oxidation and Reduction always occur together in what is called a redox reaction.

Redox in terms of Electron Transfer

• Redox reactions can also be described in terms of electron transfers.
• Reduction is the gain of electrons, and oxidation is the loss of electrons.

Consider the following reaction:

Mg_(s) + Cl_2(g) MgCl_2(s)

• The above reaction shows the reaction between Magnesium and Chlorine.
• There are two processes occurring in the reaction above:

1. Oxidation reaction:     The magnesium is losing electrons:
   
   Mg Mg²⁺ + 2e^-
   
2. Reduction reaction:     The chlorine is gaining electrons:
   
   Cl₂ + 2e^- 2Cl^-

• These two equations are called half equations.
• The reducing agent is an electron donor, in the previous reaction it is the Magnesium.
• The oxidising agent is an electron acceptor (thief), in the previous reaction it is the Chlorine.

Oxidation Numbers

• Oxidation numbers are assigned to an atom or an ion to describe its relative state of oxidation or reduction. Using the oxidation numbers, it is possible to deduce whether a redox reaction has occurred and to identify the oxidising agent and reducing agent.
• Rules for assigning Oxidation numbers:

• If the oxidation number increases, the element is oxidised.
• If the oxidation number decreases, the element is reduced.
• The total change in an oxidation number increasing and decreasing in a reaction, must be the same in both directions.

Useful books for revision: