Redox Reactions
Syllabus Unit: Physical Chemistry
Why it matters
Chapter overview
This topic covers electron-transfer chemistry from its classical (oxygen/hydrogen) and modern (electron-transfer) definitions, the oxidation-number bookkeeping rules used to track electron flow, the two systematic methods for balancing complex redox equations, and the electrochemical cells that harness spontaneous redox reactions to generate electricity.
1Oxidation & Reduction
Oxidation and reduction always occur together — one species loses electrons exactly as another gains them, which is why every redox reaction is really two half-reactions in disguise.
Learning Objectives
- State the classical (oxygen/hydrogen-transfer) and modern (electron-transfer) definitions of oxidation and reduction.
- Identify the oxidising agent and reducing agent in a given redox reaction.
- Explain why oxidation and reduction must occur simultaneously (conservation of electrons).
- Recognise disproportionation reactions, where the same element is simultaneously oxidised and reduced.
Key Terms
Two historical lenses describe the same phenomenon, converging on the modern electron-transfer definition used throughout chemistry today.
Classical vs modern definitions
- Classical oxidation — gain of oxygen or loss of hydrogen, e.g. 2Mg + O2 → 2MgO.
- Classical reduction — loss of oxygen or gain of hydrogen.
- Modern (electronic) oxidation — loss of electrons, e.g. Zn → Zn²⁺ + 2e⁻.
- Modern (electronic) reduction — gain of electrons, e.g. Cu²⁺ + 2e⁻ → Cu.
Disproportionation
A special redox reaction where the SAME element, in the same starting species, is simultaneously oxidised and reduced into two different products — e.g. 2H2O2 → 2H2O + O2, where oxygen goes from -1 in H2O2 to -2 in H2O and 0 in O2.