Identify redox reactions as reactions involving gain and loss of electrons

Chemical Reactions – Redox (Reduction‑Oxidation)

Learning Objective

By the end of this lesson you will be able to identify redox reactions as those in which electrons are transferred – i.e. one species is oxidised (loses electrons) and another is reduced (gains electrons). You will also be able to name the oxidising and reducing agents and write the half‑reactions required for balancing.

1. Key Definitions (Cambridge IGCSE 0620)

• Oxidation – loss of electrons (historically described as gain of oxygen).
• Reduction – gain of electrons (historically described as loss of oxygen).
• A redox reaction always contains both an oxidation and a reduction step.
• Oxidising agent – the species that is reduced; it causes another species to be oxidised.
• Reducing agent – the species that is oxidised; it causes another species to be reduced.

2. Assigning Oxidation Numbers – The “Roman‑Numeral” Method

The oxidation number (ON) is a bookkeeping charge that tracks electron transfer. Use the three fundamental rules first; then apply the common‑value table.

1. Zero rule: An element in its free (uncombined) state has ON = 0 (e.g. Na, O₂, Cl₂).
2. Ionic rule: For a monatomic ion the ON equals the ionic charge (e.g. Na⁺ = +1, Cl⁻ = –1).
3. Sum rule: The algebraic sum of the ONs in a neutral compound is 0; in a polyatomic ion it equals the ion’s overall charge.

2.1 Quick‑Reference Table of Common Oxidation Numbers

3. Step‑by‑Step Procedure to Identify a Redox Reaction

1. Write the balanced molecular equation.
2. Assign oxidation numbers to every atom on both sides using the rules above.
3. Compare the oxidation numbers:
   • If an element’s ON increases → it is **oxidised** (loss of electrons).
   • If an element’s ON decreases → it is **reduced** (gain of electrons).
4. Check electron balance: total electrons lost must equal total electrons gained.
5. Identify agents: the species that is oxidised is the reducing agent; the species that is reduced is the oxidising agent.
6. Optional historical check: look for a gain or loss of oxygen atoms (useful for quick mental verification).

4. Oxidising and Reducing Agents – Quick Reference

5. Laboratory Indicators of Redox Processes

• Colour change of transition‑metal ions
  • Permanganate (KMnO₄) – deep purple → Mn²⁺ (pale pink or colourless) → reduction.
  • Iodide (I⁻) – colourless → I₂ (brown‑violet) → oxidation.
  • Cu²⁺ (blue) → Cu⁺ (colourless) or Cu (metallic red) → reduction.
• Gas evolution
  • H₂ from acids reacting with active metals (reduction of H⁺).
  • O₂ from decomposition of H₂O₂ (oxidation of O⁻¹ to O⁰).
  • Cl₂ when chlorides are oxidised (e.g. Cl⁻ → Cl₂).
• Precipitate formation can accompany redox, e.g. Ag⁺ + Br⁻ → AgBr(s) while Ag⁺ is reduced to Ag⁰ in a separate step.

6. Worked Examples (Cambridge‑style)

Example 1 – Combustion of Methane

$$\mathrm{CH_4 + 2\,O_2 \;\longrightarrow\; CO_2 + 2\,H_2O}$$

Carbon is oxidised, oxygen is reduced – a classic redox reaction.

Example 2 – Zinc + Hydrochloric Acid

$$\mathrm{Zn + 2\,HCl \;\longrightarrow\; ZnCl_2 + H_2}$$

Zn is the reducing agent; H⁺ is the oxidising agent.

Half‑Reaction Method (Zn + HCl)

• Oxidation: