Oxidation-Reduction (Redox) Reactions (College Board AP® Chemistry): Revision Note

Oxidation-Reduction Reactions

Half-reactions

• During redox reactions, reduction and oxidation occur at the same time

• However, it is convenient to consider them as separate processes

• For example, consider the following net ionic equation

Sn²⁺ (aq) + 2Fe³⁺ (aq) → Sn⁴⁺ (aq) + 2Fe²⁺ (aq)

• Sn²⁺ oxidizes into Sn⁴⁺ by losing two electrons. Therefore, the oxidation half-reaction can be written as:

Oxidation:

Sn²⁺ (aq)  → Sn⁴⁺ (aq) + 2e^–

• Fe³⁺ is reduced to Fe²⁺ by gaining one electron. Therefore the reduction half-reaction can be written as:

Reduction:

Fe³⁺ (aq) + 1e^– → Fe²⁺ (aq)

• The number of electrons gained in the reduction half-reaction must be equal to the number of electrons lost in the oxidation half-reaction, so the iron half-reaction is multiplied by two

2Fe³⁺ (aq) + 2e^– → 2Fe²⁺ (aq)

• If this condition is met, the electrons cancel out when both half-reactions are added

• When both half-reactions are added, the result must be the net ionic equation

Balancing chemical equations by the method of half-reactions

• For balancing a redox reaction using the half-reactions, follow these steps:

1. Divide the equation into half-reactions (one oxidation and one reduction)

2. Balance each half-reaction using the following rules:

   1. First, balance elements other than H and O

   2. Second, balance O atoms by adding H₂O

   3. Third, balance H atoms by adding H⁺ (this method is for reactions in acidic solution)

   4. Finally, balance charges by adding e^-

3. Multiply half-reactions by integers to make the number of electrons in both half-reactions the same

4. Add both half-reactions, and cancel chemical species that appear on both sides

5. Check if the atoms are balanced