How chemical reactions make electricity and electricity drives reactions
In a galvanic cell, such as the Daniell cell, a spontaneous redox reaction produces electricity. Oxidation occurs at the anode and reduction at the cathode, and a salt bridge completes the circuit. Each half-cell has an electrode potential, measured against the standard hydrogen electrode, which is set at 0 V. The cell emf is E°cell = E°cathode − E°anode, and a positive value means the reaction is spontaneous, linked by ΔG° = −nFE°.
The Nernst equation, E = E° − (0.0591/n) log Q at 298 K, shows how emf depends on concentration. Conductance of electrolytes is measured as molar conductivity, which increases on dilution; Kohlrausch’s law lets the limiting molar conductivity of weak electrolytes be calculated. In electrolysis, electricity forces a non-spontaneous reaction; Faraday’s laws say the mass deposited is proportional to charge passed, with 1 F = 96,500 C per mole of electrons. Corrosion of iron is an electrochemical process.
Why it matters
Batteries in phones, cars and inverters, from dry cells to lithium-ion cells, are electrochemical cells. Fuel cells powered spacecraft and are being developed for clean vehicles. Electrolysis is used to extract aluminium, refine copper and electroplate jewellery. Understanding corrosion helps protect ships and pipelines through galvanising and sacrificial protection.
Key terms
- Electrode potential
- The tendency of an electrode to gain or lose electrons, measured in volts.
- Salt bridge
- A tube of electrolyte that connects two half-cells and keeps them electrically neutral.
- Nernst equation
- An equation that relates a cell’s emf to the concentrations of the species involved.
- Faraday constant
- The charge carried by one mole of electrons, about 96,500 coulombs.