Redox Reactions – Identifying Oxidising and Reducing Agents (IGCSE 0620)
1. Core Definitions (Cambridge syllabus)
- Oxidation – loss of electrons (modern definition).
Historically described as “loss of oxygen” or “gain of hydrogen”.
- Reduction – gain of electrons (modern definition).
Historically described as “gain of oxygen” or “loss of hydrogen”.
- Oxidising agent (oxidant) – the substance that causes oxidation; it is itself reduced.
- Reducing agent (reductant) – the substance that causes reduction; it is itself oxidised.
The electron‑transfer definitions are the ones examined in the Cambridge IGCSE, although the historic wording may appear in the textbook.
2. Writing Oxidation Numbers (Roman‑numeral convention)
When you label oxidation numbers in the exam, write them as Roman numerals (e.g. FeIII, MnVII). This is a specific requirement of the Cambridge syllabus.
2.1. Full Set of Oxidation‑Number Rules
- Element in its standard state = 0.
- Mono‑atomic ion: oxidation number = ionic charge.
- Oxygen: usually –2.
- Peroxides (e.g. H2O2) = –1.
- Super‑oxides (e.g. KO2) = –½.
- Hydrogen: usually +1.
- In metal hydrides (e.g. NaH) = –1.
- Halogens: –1 unless bonded to a more electronegative element (e.g. ClO4⁻, where Cl is +7).
- Alkali metals (Group 1) = +1; alkaline‑earth metals (Group 2) = +2.
- Sum rule: the algebraic sum of the oxidation numbers in a neutral molecule is 0; in an ion it equals the overall charge.
- Transition‑metal ions may exhibit several oxidation states (e.g. Fe II/III, Mn VII). Use Roman numerals when writing them.
3. Step‑by‑Step Method to Identify Oxidising and Reducing Agents
- Write the correctly balanced chemical equation.
- Assign oxidation numbers (Roman numerals) to every atom on both sides.
- Atoms whose oxidation numbers **increase** are **oxidised**.
- Atoms whose oxidation numbers **decrease** are **reduced**.
- The **species that is oxidised** is the reducing agent.
The **species that is reduced** is the oxidising agent.
4. Worked Examples (with Roman‑numeral notation)
| Reaction |
Oxidation‑Number Changes |
Oxidising Agent |
Reducing Agent |
| $$\mathrm{2Mg + O_2 \rightarrow 2MgO}$$ |
Mg: 0 → II (increase)
O: 0 → -II (decrease)
|
O2 (oxygen is reduced to O-II) |
Mg (magnesium is oxidised to MgII) |
| $$\mathrm{Cu + 2Ag^{+} \rightarrow Cu^{2+} + 2Ag}$$ |
Cu: 0 → II (increase)
Ag: +I → 0 (decrease)
|
Ag+ (silver ion gains electrons) |
Cu (copper loses electrons) |
| $$\mathrm{H_2SO_4 + Zn \rightarrow ZnSO_4 + H_2}$$ |
Zn: 0 → II (increase)
H (in H2SO4): +I → 0 (decrease)
|
H2SO4 (the H⁺ ions are reduced to H2) |
Zn (zinc metal is oxidised) |
| $$\mathrm{2KMnO_4 + 5Fe^{2+} + 8H^{+} \rightarrow 2Mn^{2+} + 5Fe^{3+} + 4H_2O + 2K^{+}}$$ |
Mn: VII → II (decrease)
Fe: II → III (increase)
|
MnO4⁻ (permanganate ion is reduced) |
Fe2+ (ferrous ion is oxidised) |
| $$\mathrm{2Al + 3Cl_2 \rightarrow 2AlCl_3}$$ |
Al: 0 → III (increase)
Cl: 0 → -I (decrease)
|
Cl2 (chlorine is reduced) |
Al (aluminium is oxidised) |
5. Half‑Reaction Method (useful for complex equations)
This method is optional but highly effective when the overall redox changes are not obvious.
- Separate the overall reaction into an oxidation half‑reaction (electrons on the right) and a reduction half‑reaction (electrons on the left).
- Balance each half‑reaction for all atoms except O and H.
- Balance O by adding H2O, then balance H by adding H⁺ (acidic medium) or OH⁻ (basic medium).
- Balance the charge by adding electrons.
- Multiply the half‑reactions so that the number of electrons transferred is the same, then add them together; electrons cancel.
Example – copper‑silver reaction
Oxidation: $$\mathrm{Cu \rightarrow Cu^{2+} + 2e^-}$$
Reduction: $$\mathrm{2Ag^{+} + 2e^- \rightarrow 2Ag}$$
Combined (electrons cancel): $$\mathrm{Cu + 2Ag^{+} \rightarrow Cu^{2+} + 2Ag}$$
6. Practice Questions (space for working is provided in the exam booklet)
- For the reaction $$\mathrm{2Al + 3Cl_2 \rightarrow 2AlCl_3}$$, identify the oxidising and reducing agents.
- In the reaction $$\mathrm{Na_2S_2O_3 + I_2 \rightarrow Na_2S_4O_6 + 2I^-}$$, determine which species is oxidised and which is reduced.
- Write the half‑reactions for the redox process in $$\mathrm{Fe^{2+} + MnO_4^- + H^{+} \rightarrow Fe^{3+} + Mn^{2+} + H_2O}$$ and identify the oxidising and reducing agents.
7. Summary Checklist (Cambridge‑style)
- Assign oxidation numbers using the complete set of rules (including peroxides, super‑oxides, metal hydrides, and transition‑metal states).
- Write oxidation numbers as Roman numerals (e.g. FeIII, MnVII).
- Identify increases → oxidation; decreases → reduction.
- Remember: oxidised species = reducing agent; reduced species = oxidising agent.
- Use the half‑reaction method for any reaction that is not immediately recognisable as redox.
- Check the sum rule: the total of all oxidation numbers equals 0 for a neutral compound and equals the ionic charge for an ion.