1.9.3 Redox Equations

What is a Redox Reaction?

A redox reaction involves both oxidation (loss of electrons) and reduction (gain of electrons) occurring simultaneously. During a redox reaction, one species is oxidized and another is reduced.

Example of a Redox Reaction

Consider the reaction:

6HI + H_2SO_4 \rightarrow 3I_2 + S + 4H_2O

Oxidation: Iodine in $HI$ is oxidized from an oxidation state of -1 to 0 in $I_2$ .
Reduction: Sulfur in $H_2SO_4$ is reduced from an oxidation state of +6 to 0 in elemental sulfur. Thus, oxidation and reduction occur together in the same reaction, making it a redox process.

Special Case: Disproportionation Reaction

In a disproportionation reaction, the same element is both oxidized and reduced in a single reaction.

Example:

Cl_2 + H_2O \rightarrow HCl + HOCl

In this reaction:

Chlorine is oxidized from 0 in $Cl_2$ to +1 in $HOCl$ .
Chlorine is also reduced from 0 in $Cl_2$ to -1 in $HCl$ . This is a disproportionation reaction because chlorine undergoes both oxidation and reduction.

Oxidizing and Reducing Agents

Oxidizing agents are electron acceptors. They cause other substances to lose electrons (be oxidized) and are themselves reduced.
Reducing agents are electron donors. They cause other substances to gain electrons (be reduced) and are themselves oxidized.

Half-Equations in Redox Reactions

Half-equations are used to represent the oxidation or reduction process separately, showing the transfer of electrons. This is important for balancing redox reactions.

Writing Half-Equations

Oxidation Half-Equation:

Represents the loss of electrons.

\text{Example: } Fe \rightarrow Fe^{3+} + 3e^-

Reduction Half-Equation:

Represents the gain of electrons.

\text{Example: } Cl_2 + 2e^- \rightarrow 2Cl^-

Combining Half-Equations

To form the overall balanced redox equation, the oxidation and reduction half-equations must be combined. The number of electrons in each half-equation must be equal, allowing the electrons to cancel out.

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Example: Iron and Chlorine Reaction

2Fe + 3Cl_2 \rightarrow 2FeCl_3

Step 1: Oxidation Half-Equation:

2Fe \rightarrow 2Fe^{3+} + 6e^-

Step 2: Reduction Half-Equation:

3Cl_2 + 6e^- \rightarrow 6Cl^-

By combining these equations, the electrons cancel out, resulting in the balanced overall redox equation:

2Fe + 3Cl_2 \rightarrow 2FeCl_3

Balancing More Complex Redox Reactions

In more complicated reactions, such as when an oxidizing or reducing agent contains oxygen or hydrogen, water molecules ( $H_2O$ ) and hydrogen ions ( $H^+$ ) may be needed to balance the equation.

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Example: Manganate(VII) to Manganese(II) In an acidic solution, $MnO_4^-$ is reduced to $Mn^{2+}$ :

Step 1: Reduction Half-Equation:

MnO_4^- \rightarrow Mn^{2+}

Step 2: Add H^+ and H_2O to balance oxygen and hydrogen:

MnO_4^- + 8H^+ \rightarrow Mn^{2+} + 4H_2O

Step 3: Balance the charges by adding electrons:

MnO_4^- + 8H^+ + 5e^- \rightarrow Mn^{2+} + 4H_2O

Redox Equations Simplified Revision Notes for A-Level AQA Chemistry