Chapter 8: Redox Reactions

Chapter 8: Redox Reactions

Oxidation and Reduction

- Oxidation: Loss of electrons/increase in oxidation number.

- Reduction: Gain of electrons/decrease in oxidation number.

- Redox reactions always involve both processes happening together.

Oxidation Numbers

- A way to keep track of electron transfer.

- Rules:

  - Elements in their natural state = 0.

  - Oxygen usually = –2 (except in peroxides = –1).

  - Hydrogen usually = +1 (except in metal hydrides = –1).

  - Group 1 metals = +1, Group 2 metals = +2.

- A change in oxidation number shows whether oxidation or reduction has occurred.

Identifying Redox Reactions

- Look for changes in oxidation numbers.

- Example:  

  Zn + Cu²⁺ → Zn²⁺ + Cu  

  - Zn: 0 → +2 (oxidised).  

  - Cu: +2 → 0 (reduced).

Redox Equations

- Balanced equations must show both electron transfer and atom balance.

- Steps:

  1. Write half-equations for oxidation and reduction.

  2. Balance atoms and charges.

  3. Combine half-equations to form an overall balanced redox equation.

Disproportionation

- A single species is both oxidised and reduced.

- Example:  

  Cl₂ + H₂O → HCl + HClO  

  - Cl₂ → Cl⁻ (reduced).  

  - Cl₂ → Cl⁺ (oxidised).

Oxidising and Reducing Agents

- Oxidising agent: Accepts electrons, gets reduced.

- Reducing agent: Donates electrons, gets oxidised.

- Examples:

  - Oxidising agents: KMnO₄, K₂Cr₂O₇.

  - Reducing agents: Zn, Fe²⁺, SO₂.

Applications of Redox

- Extraction of metals (e.g., blast furnace).

- Electrochemical cells and batteries.

- Bleaching and disinfecting (oxidising agents).

- Analytical titrations (redox titrations).