Oxidising and reducing agents

Table of Contents

1. Oxidising agents and reducing agents
2. Redox reactions
   1. Oxidation numbers
3. Oxidising agents and reducing agents
   1. Oxidants and reductants
   2. Common oxidants and reductants
4. Half-equations
   1. Writing half-equations
   2. Combining half-equations
5. Reduction potentials
   1. Source of reduction potentials
   2. Using standard electrode potentials
   3. Predicting reaction tendencies
6. Redox reactions in action

1. Additional resources:
   1. Glossary
   2. Print content

Combining half-equations

Half-equations are also helpful to use when writing balanced equations for redox reactions. The method used involves identifying the correct half-equations to use, and combining the half-equations in a way that causes the electrons involved to cancel out. Again you will have encountered this type of problem in Stage 2 Chemistry so you can review your notes before tackling this section.

The method to use is shown below.

Method for combining half-equations

This method can be used to combine any two half-equations.

Step 1: Write down the two half-equations showing the oxidation and reduction processes.

Step 2: If required, multiply the equations in order to obtain the same number of electrons on each side of the equations.

Step 3: Cancel the electrons and combine the equations by adding all the remaining species.

Step 4: If required, cancel out any species that are present on both sides of the equation.

Redox reaction example

Reaction of potassium permanganate (KMnO₄) with iron(II) sulfate (FeSO₄). (This reaction can be used to analyse the amount of iron in iron tablets.) Work through the following steps.

Disproportionation

Many substances have the ability to act as reductants and oxidants depending on the conditions and/or the other chemicals present.

For example, chlorine (Cl₂) can be oxidised to hypochlorite (ClO^-) as shown below.

Cl₂ + 2 H₂O → 2 ClO^- + 4 H⁺ + 2 e^-

In a different situation chlorine (Cl₂) can be reduced to chloride (Cl^-) as shown below.

Cl₂ + 2 e^- → 2 Cl^-

Sometimes a species is oxidised and reduced in the same reaction. This process is called disproportionation. The overall redox reaction is found by combining the two half-equations in the same way as shown in the previous examples.

In your notes, combine the two half-equations above to produce an equation that shows the disproportionation of chlorine.

Click here to reveal the answer.

Now complete the following problems to practise combining half-equations .

The answers can be accessed here .