Introduction to Electrolysis
Electrolysis is a process where electrical energy is used to drive a non-spontaneous chemical reaction. The electrical energy provided is converted into chemical energy, driving the endothermic electrolytic reactions.
Use this page to revise the following concepts within introduction to Electrolysis:
Components of an Electrolytic Cell
Electrolyte
This is typically a molten substance or an aqueous solution containing free-moving ions. These ions can gain or lose electrons, allowing the current to flow through the external circuit .
Anions migrate to the anode and cations migrate to the cathode .
Electrodes
There are two electrodes where the electrolysis reactions occur, one acts as the anode and the other as the cathode.
The electrode connected to the negative terminal of the power source is the cathode, because it is where electrons are supplied. At the cathode, reduction takes place, with the substance gaining electrons.
The electrode connected to the positive terminal of the power source is the anode because it is where electrons are drawn away. At the anode, oxidation occurs, meaning electrons are lost by the substance being oxidised.
Inert electrodes (such as graphite or platinum) can be used in the electrolysis process. These electrodes do not participate in the chemical reactions themselves but provide a surface for the ions to undergo oxidation or reduction.
External power source
An external source, such as a battery or power pack, supplies the energy needed to drive the reaction. The flow of electrons is moving from the anode to the cathode.
Predicting Reactions Based on the Electrochemical Series
The electrochemical series lists common reduction half reaction, in order of their Standard electrode potential, \(E^0\), or reduction potential. The following is an excerpt from an electrochemical series.
| Reactions | Standard electrode potential \((E_0)\) at \(25^{\circ}C\) |
|---|---|
| \(Cl_2(s) + 2e^- \rightleftharpoons 2Cl^-(aq)\) | \(+1.36 V\) |
| \(O_2(g) + 4H^+(aq) + 4e^- \rightleftharpoons 2H_2O(l)\) | \(+1.23 V\) |
| \(I_2(s) + 2e- \rightleftharpoons 2I^-(aq)\) | \(+0.54 V\) |
| \(Cu^{2+} (aq) + 2e^- \rightleftharpoons Cu(s)\) | \(+0.34 V\) |
| \(Ni^{2+} (aq) + 2e^- \rightleftharpoons Ni(s)\) | \(-0.25 V\) |
| \(Co^{2+} (aq) + 2e^{-} \rightleftharpoons Co(s)\) | \(-0.28 V\) |
| \(Fe^{2+} (aq) + 2e^- \rightleftharpoons Fe(s)\) | \(-0.44 V\) |
| \(Zn^{2+} (aq) + 2e^- \rightleftharpoons\) Zn(s)\) | \(-0.76 V\) |
| \(2H_2O(l) + 2e^- \rightleftharpoons H2(g) + 2OH^-(aq)\) | \(-0.83 V\) |
| \(Na^+(aq) + e^- \rightleftharpoons Na(s)\) | \(-2.71 V\) |
| \(K^+(aq) + e^- \rightleftharpoons K(s)\) | \(-2.93 V\) |
When conducting electrolysis, you must consider the competing reactions that can occur at each electrode, as multiple chemicals may be present. In aqueous solutions, for example, water may also participate in the reaction.
The reaction that has the highest reduction potential in the electrochemical series is most likely to occur at the cathode , as reduction is favoured for the strongest oxidising agent.
The reaction with the lowest reduction potential, the strongest reducing agent, is most likely to occur at the anode .
