Define reduction in terms of: (a) gain of electrons (b) a decrease in oxidation number

Redox (Oxidation‑Reduction) Reactions

Learning objectives

By the end of this lesson you will be able to:

• State the two textbook definitions of reduction and of oxidation.
• Identify the oxidising and reducing agents in a redox equation (including colour‑change tests required by the IGCSE syllabus).
• Write the reduction and oxidation half‑reactions in acidic and basic media.
• Balance complete redox equations using the half‑reaction method.
• Explain how a change in oxidation number reflects the gain or loss of electrons.
• Recognise how redox principles are applied in electrolysis and in hydrogen‑oxygen fuel cells.

1. Definitions

1.1 Reduction

In a redox reaction reduction is the process by which a chemical species:

1. Gains one or more electrons (e⁻); electrons appear on the left‑hand side of the half‑reaction.
2. Consequently its oxidation number (or oxidation state) becomes lower – i.e. it becomes less positive or more negative.

1.2 Oxidation

Oxidation is the opposite process. A species is oxidised when it:

1. Loses one or more electrons (electrons appear on the right‑hand side of the half‑reaction).
2. Its oxidation number therefore increases – it becomes more positive or less negative.

1.3 Quick reminder – Roman numerals

In the Cambridge syllabus oxidation numbers are written as Roman numerals in formulas (e.g. Fe III, Cu II). When you assign oxidation numbers, record them in this form – it is required for AO1 marks.

2. Oxidising and Reducing Agents

The species that causes the other to be reduced is the oxidising agent; the species that causes the other to be oxidised is the reducing agent.

Oxidising agents (common)	Typical colour change (acidic test)
Potassium permanganate, KMnO₄	Colourless → deep purple (MnO₄⁻)
Hydrogen peroxide, H₂O₂	Colourless → pale yellow (O₂ evolution)
Potassium dichromate, K₂Cr₂O₇	Orange → green (Cr³⁺)
Concentrated nitric acid, HNO₃	No permanent colour – brown fumes of NO₂

Reducing agents (common)	Typical colour change (acidic test)
Zinc metal, Zn	No colour change; Zn²⁺ is colourless
Hydrogen gas, H₂	Colourless → colourless (H⁺ formed)
Carbon monoxide, CO	Colourless → colourless (CO₂ formed)
Iron(II) sulphate, FeSO₄	Pale green → colourless (Fe³⁺ formed)

3. How electron transfer relates to oxidation numbers

When an atom, ion or molecule accepts electrons, the extra negative charge reduces the effective charge on that element. This is recorded as a lower (more negative) oxidation number. Conversely, loss of electrons raises the oxidation number. The two textbook descriptions – “gain of electrons” and “decrease in oxidation number” – are therefore equivalent.

4. Writing half‑reactions

4.1 General steps (acidic medium)

1. Assign oxidation numbers to every element in reactants and products.
2. Identify which species is reduced (oxidation number falls) and which is oxidised (oxidation number rises).
3. Write separate half‑reactions for the reduction and oxidation processes, placing electrons on the appropriate side.
4. Balance each half‑reaction:
   • Balance all atoms except H and O.
   • Balance O by adding H₂O.
   • Balance H by adding H⁺.
   • Balance charge by adding electrons.
5. Multiply the half‑reactions by whole numbers so that the electrons cancel, then add them together.

4.2 Converting to basic medium

If the reaction occurs in a basic solution, after completing the acidic‑medium steps:

1. Add the same number of OH⁻ to both sides of each half‑reaction to neutralise all H⁺.
2. Combine OH⁻ with H⁺ to form H₂O.
3. If water appears on both sides, cancel it.

5. Worked examples

Example 1 – Reduction of copper(II) oxide by hydrogen (acidic medium)

Overall reaction:

$$\mathrm{CuO(s) + H_2(g) \;\longrightarrow\; Cu(s) + H_2O(l)}$$

Species	Cu	O	H
CuO (reactant)	+II	–II	–
Cu (product)	0	–	–
H₂ (reactant)	–	–	0
H₂O (product)	–	–II	+I

Reduction (Cu):

$$\mathrm{Cu^{2+} + 2e^- \;\longrightarrow\; Cu}$$

Oxidation (H₂):

$$\mathrm{H_2 \;\longrightarrow\; 2H^+ + 2e^-}$$

Adding the two half‑reactions gives the overall equation shown above.

Example 2 – Reduction of iron(III) oxide by carbon monoxide (acidic medium)

Overall reaction:

$$\mathrm{Fe_2O_3(s) + 3CO(g) \;\longrightarrow\; 2Fe(s) + 3CO_2(g)}$$

Species	Fe	O	C
Fe₂O₃ (reactant)	+III	–II	–
Fe (product)	0	–	–
CO (reactant)	–	–II	+II
CO₂ (product)	–	–II	+IV

Reduction (Fe³⁺ → Fe):

$$\mathrm{Fe^{3+} + 3e^- \;\longrightarrow\; Fe}$$

Oxidation (CO → CO₂):

$$\mathrm{CO + \tfrac{1}{2}O_2 \;\longrightarrow\; CO_2 + 2e^-}$$

Multiplying the oxidation half‑reaction by 3 and the reduction half‑reaction by 2 gives the balanced overall equation.

Example 3 – Balancing in basic medium: MnO₄⁻ → MnO₂

Unbalanced skeletal equation (basic solution):

$$\mathrm{MnO_4^- \;\longrightarrow\; MnO_2}$$

1. Assign oxidation numbers: Mn +VII → Mn +IV (reduction).
2. Write the reduction half‑reaction in acidic medium:
   $$\mathrm{MnO_4^- + 4H^+ + 3e^- \;\longrightarrow\; MnO_2 + 2H_2O}$$
3. Convert to basic medium: add 4 OH⁻ to each side:
   $$\mathrm{MnO_4^- + 4H^+ + 4OH^- + 3e^- \;\longrightarrow\; MnO_2 + 2H_2O + 4OH^-}$$ $$\Rightarrow\; \mathrm{MnO_4^- + 4H_2O + 3e^- \;\longrightarrow\; MnO_2 + 6H_2O}$$ Cancel water to obtain the final basic‑medium half‑reaction:
   $$\boxed{\mathrm{MnO_4^- + 2H_2O + 3e^- \;\longrightarrow\; MnO_2 + 4OH^-}}$$

6. Redox in practice

6.1 Electrolysis of molten salts

When molten NaCl is electrolysed:

• At the cathode: Na⁺ + e⁻ → Na(l) (reduction of Na⁺).
• At the anode: 2Cl⁻ → Cl₂(g) + 2e⁻ (oxidation of Cl⁻).
• The overall reaction: 2NaCl(l) → 2Na(l) + Cl₂(g).

6.2 Hydrogen‑oxygen fuel cell

Overall cell reaction (acidic electrolyte):

$$\mathrm{2H_2(g) + O_2(g) \;\longrightarrow\; 2H_2O(l)}$$

Half‑reactions:

• Reduction (at the cathode): O₂ + 4H⁺ + 4e⁻ → 2H₂O.
• Oxidation (at the anode): 2H₂ → 4H⁺ + 4e⁻.

These illustrate how redox underpins modern energy technologies.

7. Practice tasks

For each reaction underline the species that is reduced and write its reduction half‑reaction (include electrons on the left). Use the appropriate medium (acidic unless the question states “basic”).

1. \(\displaystyle \mathrm{Zn(s) + CuSO_4(aq) \rightarrow ZnSO_4(aq) + Cu(s)}\)
2. \(\displaystyle \mathrm{2Al(s) + 3Cl_2(g) \rightarrow 2AlCl_3(s)}\)
3. \(\displaystyle \mathrm{MnO_2(s) + 4HCl(aq) \rightarrow MnCl_2(aq) + Cl_2(g) + 2H_2O(l)}\)
4. \(\displaystyle \mathrm{PbO_2(s) + 4H^+(aq) + 2e^- \rightarrow Pb^{2+}(aq) + 2H_2O(l)}\) (identify the reduction step in the overall process)

8. Quick checklist for the exam (AO1 + AO2)

• Write oxidation numbers (Roman numerals) for every element in reactants and products.
• Identify the species whose oxidation number falls – that species is reduced (oxidising agent).
• Identify the species whose oxidation number rises – that species is oxidised (reducing agent).
• Write the half‑reaction for the reduction, placing electrons on the left.
• Balance the half‑reaction (atoms, then charge). Use H₂O and H⁺ for acidic media; add OH⁻ to convert to basic media.
• Make sure the total electrons lost = total electrons gained before adding the half‑reactions.
• Check that the final overall equation is balanced for mass and charge.