Electricity driving chemistry
An electrolytic cell uses an external direct current to force a non-spontaneous redox reaction. The process is called electrolysis: an electrolyte (molten or aqueous ionic compound) is decomposed by electricity. Two electrodes, the anode (positive) and cathode (negative), dip into the electrolyte.
During electrolysis, cations move to the cathode where they gain electrons (reduction); anions move to the anode where they lose electrons (oxidation).
Worked example
Electrolysis of molten lead(II) bromide, PbBr2:
Cathode: Pb2+ + 2e− → Pb (grey metal forms)
Anode: 2Br− → Br2 + 2e− (brown vapour forms)
Factors in aqueous electrolysis
In aqueous solutions, water also provides H+ and OH−, so selective discharge depends on three factors: the position of ions in the electrochemical series (lower ions are discharged more easily), the concentration of ions, and the type of electrode (reactive electrodes such as copper take part in the reaction).
Remember
- Cathode = negative = reduction (cations gain electrons).
- Anode = positive = oxidation (anions lose electrons).
- PANC: Positive Anode, Negative Cathode.
- A concentrated halide favours discharge of the halide over OH−.
Electrolysis is used for extracting reactive metals, purifying copper and electroplating objects with protective or decorative metal layers.
An electrolytic cell is the opposite of a voltaic cell: rather than a reaction producing electricity, electricity forces a non-spontaneous reaction to happen. In the purification of copper, an impure copper anode dissolves (Cu → Cu2+ + 2e−) and pure copper is deposited on the cathode. In electroplating, the object to be coated is made the cathode and the plating metal is made the anode, so metal ions travel from anode to cathode and build up an even, protective layer.