IGCSE Chemistry 0620 - Redox Reactions

Objective: Identify Oxidising Agents and Reducing Agents in Redox Reactions

What are Redox Reactions?

Redox reactions, short for reduction-oxidation reactions, involve the transfer of electrons between chemical species. These reactions always occur together: one species loses electrons (oxidation), and another species gains electrons (reduction).

Key Terms

• Oxidation: The loss of electrons by a substance. Oxidation always occurs when another substance is being reduced.
• Reduction: The gain of electrons by a substance. Reduction always occurs when another substance is being oxidised.
• Reducing Agent: The substance that causes reduction by donating electrons. It gets oxidised in the process.
• Oxidising Agent: The substance that causes oxidation by accepting electrons. It gets reduced in the process.

Oxidation Number (Assigning Oxidation Numbers)

Oxidation numbers are a way of keeping track of electron transfer in redox reactions. They are hypothetical charges that atoms would have if all bonds were completely ionic. Rules for assigning oxidation numbers:

1. The oxidation number of an atom in a monatomic ion is equal to its charge.
2. The oxidation number of oxygen is usually -2 (except in peroxides like $H_2O_2$ where it's -1) and fluorine is always -1.
3. The oxidation number of hydrogen is usually +1 (except in $لح_2$ where it's -1).
4. The oxidation number of metals is usually positive.
5. The sum of the oxidation numbers in a neutral compound is zero.
6. The sum of the oxidation numbers in a polyatomic ion is equal to the charge of the ion.

Identifying Oxidising and Reducing Agents

To identify the oxidising and reducing agents, we need to look at the change in oxidation numbers during the reaction.

Oxidising Agent: The species whose oxidation number decreases during the reaction. It is reduced.

Reducing Agent: The species whose oxidation number increases during the reaction. It is oxidised.

Examples of Redox Reactions

Reaction	Oxidising Agent	Reducing Agent	Change in Oxidation Number
$Zn(s) + Cu^{2+}(aq) \rightarrow Zn^{2+}(aq) + Cu(s)$	$Cu^{2+}(aq)$ (oxidation number decreases from +2 to 0)	$Zn(s)$ (oxidation number increases from 0 to +2)	Zn is oxidised (loss of 2 electrons), Cu$^{2+}$ is reduced (gain of 2 electrons)
$2Mg(s) + O_2(g) \rightarrow 2MgO(s)$	$O_2(g)$ (oxidation number decreases from 0 to -2)	$Mg(s)$ (oxidation number increases from 0 to +2)	Mg is oxidised (loss of 2 electrons), O2 is reduced (gain of 2 electrons)
$Cl_2(g) + 2Br^-(aq) \rightarrow 2Br_2(aq)$	$Cl_2(g)$ (oxidation number increases from 0 to +1)	$Br^-(aq)$ (oxidation number decreases from -1 to 0)	Br- is oxidised (loss of 1 electron), Cl2 is reduced (gain of 1 electron)

Practice

Consider the following reaction: $Fe^{2+}(aq) + MnO_4^-(aq) \rightarrow Fe^{3+}(aq) + Mn^{2+}(aq)$

1. Identify the oxidising agent.
2. Identify the reducing agent.
3. Explain the change in oxidation numbers.

Answers to Practice

1. The oxidising agent is $MnO_4^-$ because its oxidation number decreases from +7 to +2.
2. The reducing agent is $Fe^{2+}$ because its oxidation number increases from +2 to +3.
3. In the reaction, iron loses two electrons (oxidation) and manganese gains two electrons (reduction). This is reflected in the change in oxidation numbers of the reactants and products.