Oxidation and reduction in terms of electron transfer and oxidation numbers, redox reactions, the principles of electrolysis, and the industrial use of electrolysis to extract and purify metals.

What this dot point is asking

CCEA wants you to define oxidation and reduction in terms of electron transfer, use oxidation numbers to identify redox reactions, write half-equations, describe the principles of electrolysis, and explain how electrolysis is used industrially to extract and purify metals. Redox underpins respiration, the extraction of reactive metals and the operation of cells, so it connects chemistry to biology and industry.

Oxidation, reduction and oxidation numbers

Oxidation numbers are a bookkeeping tool for electrons. Key rules: an uncombined element has oxidation number 0; a simple ion has the oxidation number of its charge; oxygen is usually minus 2 and hydrogen usually plus 1; and the oxidation numbers in a neutral compound sum to 0, or to the overall charge in an ion. By assigning oxidation numbers before and after a reaction, you can spot redox: an element whose oxidation number increases has been oxidised, and one whose oxidation number decreases has been reduced. This is useful when electron loss and gain are not obvious from the formulae.

Half-equations and redox reactions

A redox reaction can be split into two half-equations, one showing the oxidation (electrons lost, written on the right) and one showing the reduction (electrons gained, written on the left). For example, when zinc displaces copper from copper(II) sulfate solution, zinc atoms lose two electrons to form zinc ions (oxidation) and copper(II) ions gain two electrons to form copper atoms (reduction). The electrons lost by the reducing agent are exactly the electrons gained by the oxidising agent, so the half-equations balance when the electrons match. Displacement of a less reactive metal by a more reactive one is a common redox reaction tested at this level.

Electrolysis and its industrial use

Electrolysis is used industrially where chemical reduction is not practical. Aluminium is extracted by electrolysing molten aluminium oxide (dissolved in cryolite to lower the melting point): aluminium ions are reduced to aluminium metal at the cathode, and oxide ions are oxidised at the anode. This is energy-intensive but the only economic way to extract such a reactive metal. Copper is purified by electrolysis: an impure copper anode dissolves and pure copper is deposited at the cathode, leaving impurities behind. These examples show electrolysis turning electrical energy into useful chemical change in industry.

Examples in context

Example 1. Extracting reactive metals. Aluminium is too reactive to be extracted by heating its oxide with carbon, so it is obtained by electrolysis of molten aluminium oxide. The huge electrical demand explains why aluminium smelters are built near cheap electricity, linking redox chemistry to industrial economics and energy.

Example 2. Redox in respiration. In aerobic respiration, glucose is oxidised (losing electrons and hydrogen) while oxygen is reduced to water (gaining electrons). The controlled transfer of electrons releases the energy used to make ATP, showing that the redox concept from this unit underlies the biology of energy release in the Human Body Systems unit.

Try this

Q1. Define oxidation and reduction in terms of electrons. [2 marks]

• Cue. Oxidation is the loss of electrons; reduction is the gain of electrons.

Q2. State at which electrode reduction occurs in electrolysis and why. [2 marks]

• Cue. At the cathode (negative electrode); positive ions gain electrons there, which is reduction.

Q3. Deduce the oxidation number of chlorine in the chloride ion. [1 mark]

• Cue. Minus 1 (a simple ion has the oxidation number of its charge).