Electrolysis is a non-spontaneous chemical process powered by electrical energy.
The reactions happening at the electrodes in the electrolysis of a molten ionic compound are redox in nature. The cations are reduced, gaining electrons to form the metal element and the anions are oxidised, losing electrons to form the non-metal element.
The overall result is that the ionic compound is turned back into elements.
e.g. CuCl2(s) ⇾ Cu(s) + Cl2(g)
The electrons flow around the circuit and charged ions are free to move through the molten liquid electrolyte, hence the circuit is complete.
A typical practical set up for electrolysis looks like this:
Metals that are more reactive than carbon are extracted from their ores by electrolysis and once extracted, many metals are purified or refined by electrolysis.
Extraction of aluminium
Aluminium is found as an ore, bauxite, which is a mixture of minerals mainly containing different forms of aluminium hydroxide and iron oxides. The aluminium compounds are separated and converted into aluminium oxide, Al2O3, and the aluminium is then extracted by electrolysis.
Key points:
- electrolysis requires two inert electrodes such as carbon / graphite or platinum
- one is connected to the negative terminal of the battery forming the cathode and one is connected to the positive terminal forming the anode
- the molten (or aqueous) ionic salt is the electrolyte
- the cathode attracts cations which are reduced e.g. Al3+ + 3e– ⇾ 3Al
- the anode attracts anions which are oxidised e.g. 2O2- ⇾ O2 + 4e–
All of which makes predicting the products of electrolysis of molten salts straightforward, we just need to remember that most of the non-metal elements formed at the anode are diatomic (O2, Cl2, Br2, I2) so we need to balance our oxidation half equation appropriately …
E.g. in the electrolysis of molten zinc bromide, ZnBr2, zinc is formed at the cathode and bromine at the anode:
Zn2+ + 2e– ⇾ Zn (reduction)
2Br– ⇾ Br2 + 2e– (oxidation)
N.B. if oxygen is released at the anode and the anode is made of carbon / graphite, it will react so that the product is actually CO2 and the anode will need to be replaced periodically.
Practice questions
- (a) Predict the products formed at the anode and cathode for the electrolysis of each of the following molten ionic salts. The experiments were carried out with graphite electrodes.
| Ionic salt | Product at anode | Equation | Product at cathode | Equation |
|---|---|---|---|---|
| AgCl | 2Cl- ⇾ Cl2 + 2e- | |||
| PbO | carbon dioxide, CO2 | |||
| CuI2 | ||||
| K2S |
(b) Write equations for the reactions happening at each electrode.
2. Explain why electrolysis of zinc chloride will not happen in the apparatus shown below:
Answers
- (a) and (b)
| Ionic salt | Product at anode | Equation | Product at cathode | Equation |
|---|---|---|---|---|
| AgCl | chlorine, Cl2 | 2Cl– ⇾ Cl2 + 2e– | silver, Ag | Ag+ + e– ⇾ Ag |
| PbO | carbon dioxide, CO2 | 2O2- ⇾ O2 + 4e– C + O2 ⇾ CO2 | lead, Pb | Pb2+ + 2e– ⇾ Pb |
| CuI2 | iodine, I2 | 2I– ⇾ I2 + 2e– | copper, Cu | Cu2+ + 2e– ⇾ Cu |
| K2S | sulfur, S | S2- ⇾ S + 2e– | potassium, K | K+ + e– ⇾ K |
2. Ions are not free to move in solid zinc chloride which is why the electrolyte must be molten (or aqueous).