Define an oxidising agent as a substance that oxidises another substance and is itself reduced

Redox Chemistry – Oxidising Agents (Cambridge IGCSE 0620)

1. Quick Overview of Redox Reactions

• Redox = reduction‑oxidation – two processes that occur simultaneously.
• One species **loses electrons** (is **oxidised**); another **gains electrons** (is **reduced**).
• Electrons are conserved – they are transferred from the reducing agent to the oxidising agent.

2. Core Definitions (AO1)

3. Formal Definition of an Oxidising Agent

An oxidising agent appears in the **reduction half‑reaction** of a redox equation and therefore occurs on the right‑hand side (products) of the overall balanced reaction because it has gained electrons.

4. How to Identify the Oxidising Agent (AO2)

1. Assign oxidation numbers to every atom in the reactants and products.
2. Compare the numbers:
   • Species whose oxidation number **increases** are being oxidised → they are the reducing agents.
   • Species whose oxidation number **decreases** are being reduced → they are the oxidising agents.
3. Alternatively, write the two half‑reactions. The species that appears in the **reduction** half‑reaction is the oxidising agent.

5. Why a Substance Acts as an Oxidising Agent (AO3)

Three complementary explanations are accepted in the IGCSE exam:

• Oxidation‑state change: a large negative change (e.g. Mn +VII → Mn +II, Δ = ‑5) shows a strong tendency to accept electrons.
• Electron affinity / electronegativity: atoms or ions that strongly attract electrons (high electron affinity) are good oxidisers (e.g. Cl₂).
• Standard electrode potential (E°): a high positive E° for the reduction half‑reaction indicates a powerful oxidising agent.
  Example: MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O E° = +1.51 V (very strong).

6. Half‑Reaction Method – Step‑by‑Step (AO2 + AO3)

6.1 General Procedure (Acidic Medium)

1. Write the unbalanced skeletal equation.
2. Separate into oxidation and reduction half‑reactions.
3. Balance all atoms **except** H and O.
4. Balance O by adding H₂O.
5. Balance H by adding H⁺.
6. Balance charge by adding electrons (e⁻).
7. Equalise the number of electrons transferred in the two half‑reactions (multiply if necessary).
8. Add the half‑reactions together and cancel species that appear on both sides.

6.2 Adapting to Basic Medium

1. Follow steps 1‑7 as if the reaction were in acidic medium.
2. For every H⁺ present, add the same number of OH⁻ to both sides. H⁺ + OH⁻ → H₂O
3. Combine any resulting H₂O molecules and cancel where possible.

6.3 Worked Example – Acidic Medium

Reaction: Zn + Cu²⁺ → Zn²⁺ + Cu

1. Oxidation: Zn → Zn²⁺ + 2e⁻
2. Reduction: Cu²⁺ + 2e⁻ → Cu
3. Electrons are already equal (2 e⁻). Adding the two half‑reactions gives the balanced overall equation:

Zn + Cu²⁺ → Zn²⁺ + Cu

Oxidising agent: Cu²⁺ (it is reduced).

6.4 Worked Example – Basic Medium

Reaction to balance: ClO⁻ + H₂O₂ → Cl⁻ + O₂ (basic solution)

1. Oxidation half‑reaction (H₂O₂ → O₂) H₂O₂ → O₂ + 2H⁺ + 2e⁻
2. Reduction half‑reaction (ClO⁻ → Cl⁻) ClO⁻ + 2e⁻ → Cl⁻
3. Electrons are already equal (2 e⁻). Add the two half‑reactions:

H₂O₂ + ClO⁻ → O₂ + 2H⁺ + Cl⁻

4. Convert to basic medium: add 2 OH⁻ to both sides to cancel the 2 H⁺:

H₂O₂ + ClO⁻ + 2OH⁻ → O₂ + 2H₂O + Cl⁻

5. Combine H₂O where possible (2 H₂O on the right can be left as written). Final balanced equation (basic):

ClO⁻ + H₂O₂ + 2OH⁻ → Cl⁻ + O₂ + 2H₂O

Oxidising agent: ClO⁻ (it is reduced).

7. Common Oxidising Agents in the IGCSE Syllabus (AO1 + AO2)

8. Redox in the Wider Cambridge Syllabus

• Electrolysis – Redox underlies the decomposition of molten salts and aqueous solutions (e.g., electrolysis of water, NaCl). The anode reaction is oxidation; the cathode reaction is reduction.
• Energy changes – Redox reactions involve transfer of electrical energy; a positive cell potential (E°) corresponds to a spontaneous reaction (ΔG = ‑nFE°).
• Everyday examples – Rusting of iron, combustion of fuels, bleaching of fabrics, and the action of disinfectants are all redox processes.

9. Redox Titrations – Using an Oxidising Agent (AO2 + AO3)

• Principle: A standard solution of a strong oxidising (or reducing) agent reacts stoichiometrically with the analyte. The volume of titrant required gives the amount of analyte.
• Typical system: 0.020 M potassium permanganate (KMnO₄) titrated against an unknown Fe²⁺ solution in acidic medium.
• Overall reaction (acidic): MnO₄⁻ + 5Fe²⁺ + 8H⁺ → Mn²⁺ + 5Fe³⁺ + 4H₂O
• Interpretation: KMnO₄ is the **oxidising agent** (it is reduced to Mn²⁺). The colour change from colourless Mn²⁺ to faint pink MnO₄⁻ marks the end point.
• Calculation tip:

  (M·V)₍KMnO₄₎ × 5 = (M·V)₍Fe²⁺₎

  (5 mol e⁻ are transferred per mole of MnO₄⁻).

10. Safety Note (AO1)

• Strong oxidising agents (KMnO₄, K₂Cr₂O₇, conc. HNO₃, Cl₂) can ignite or explode when in contact with organic material.
• Store them in a cool, dry, well‑ventilated cupboard away from reducing agents, flammable liquids and metals.
• Wear safety goggles, gloves, and use a fume cupboard for volatile or concentrated reagents.

11. Exam Tips (AO2 + AO3)

1. Never guess the oxidising agent – always check oxidation numbers.
2. When asked for the “oxidising agent”, write its **reduction half‑reaction** as part of your answer.
3. In titration questions, identify which species is the titrant and confirm it is being reduced (oxidising agent) or oxidised (reducing agent).
4. For balancing, start with half‑reactions; this avoids missing O or H atoms.
5. After balancing, verify:
   • Mass balance (same number of each atom on both sides)
   • Charge balance (total charge equal on both sides)
6. When an AO3 explanation is required, mention at least one of the three accepted reasons (oxidation‑state change, electron affinity, or standard electrode potential) and, where possible, give a numerical E° value.

12. Practice Questions

1. In the reaction Fe + CuSO₄ → FeSO₄ + Cu:
   • Identify the oxidising agent.
   • Write its reduction half‑reaction.
2. Which of the following is NOT a useful oxidising agent in the IGCSE syllabus?
   
   • A. NaClO
   • B. Dilute H₂SO₄
   • C. K₂Cr₂O₇
3. Balance the redox reaction in basic medium: ClO⁻ + H₂O₂ → Cl⁻ + O₂.
4. A 25.0 cm³ sample of an unknown Fe²⁺ solution is titrated with 0.020 M KMnO₄. The end point is reached after 15.6 cm³ of KMnO₄. Calculate the concentration of Fe²⁺ (M).

13. Answers & Explanations

1. Oxidising agent: Cu²⁺ (from CuSO₄).
   Reduction half‑reaction: Cu²⁺ + 2e⁻ → Cu
2. Not an oxidising agent: B. Dilute H₂SO₄. It is a weak acid and does not readily accept electrons under normal IGCSE conditions.
3. Balancing in basic medium (as shown in section 6.4):

   ClO⁻ + H₂O₂ + 2OH⁻ → Cl⁻ + O₂ + 2H₂O

4. Reaction: MnO₄⁻ + 5Fe²⁺ + 8H⁺ → Mn²⁺ + 5Fe³⁺ + 4H₂O (1 mol KMnO₄ reacts with 5 mol Fe²⁺).
   Moles of KMnO₄ = 0.020 M × 0.0156 L = 3.12 × 10⁻⁴ mol.
   Moles of Fe²⁺ = 5 × 3.12 × 10⁻⁴ = 1.56 × 10⁻³ mol.
   Concentration of Fe²⁺ = 1.56 × 10⁻³ mol / 0.0250 L = **0.062 M**.