Topic 9 covers electron transfer reactions: oxidation, reduction, electrochemical cells, and electrolysis. SL introduces oxidation states and simple electrochemistry. HL adds standard electrode potentials and spontaneity predictions.
Oxidation: loss of electrons, increase in oxidation state. Reduction: gain of electrons, decrease in oxidation state. OIL RIG mnemonic. Oxidising agent: causes oxidation (is itself reduced). Reducing agent: causes reduction (is itself oxidised). Oxidation states: assign using rules (elements = 0, O = −2, H = +1 in most compounds).
Two half-cells connected by a salt bridge and external wire. Oxidation at anode (−), reduction at cathode (+). Salt bridge maintains electrical neutrality and completes the circuit. Cell notation: anode | anode solution || cathode solution | cathode. Example: Zn(s) | Zn²⁺(aq) || Cu²⁺(aq) | Cu(s). Electrons flow from more reactive to less reactive metal.
Uses electrical energy to drive non-spontaneous reactions. Electrolyte: molten ionic compound or aqueous solution. Cations migrate to cathode (reduction). Anions migrate to anode (oxidation). Factors affecting products in aqueous electrolysis: position in electrochemical series, concentration, and nature of electrodes.
E° measured against standard hydrogen electrode (SHE, 0.00 V). Cell EMF: E°cell = E°cathode − E°anode. Spontaneous if E°cell > 0 (ΔG° = −nFE°cell < 0). More positive E°: stronger oxidising agent. More negative E°: stronger reducing agent. Predicting feasibility of reactions using electrochemical series.
Voltaic (galvanic) cell: spontaneous reaction produces electrical energy (like a battery). E°cell > 0. Anode is negative, cathode is positive. Electrolytic cell: uses external electrical energy to drive a non-spontaneous reaction (like electroplating). E°cell < 0. Anode is positive, cathode is negative (connected to + and − terminals of power supply respectively). Both have oxidation at the anode and reduction at the cathode.
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