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WAChemistrySyllabus dot point

How can an external power source drive a non-spontaneous redox reaction in an electrolytic cell?

Describe the operation of electrolytic cells, predict the products of electrolysis, and explain industrial applications

A focused answer to the WACE Year 12 Chemistry dot point on electrolytic cells, how an external voltage drives non-spontaneous reactions, predicting products of molten and aqueous electrolysis, and industrial uses, with a worked example and common exam mistakes.

Generated by Claude Opus 4.77 min answer

Reviewed by: AI editorial process; not yet individually human-reviewed

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What this dot point is asking

An electrolytic cell does the opposite of a galvanic cell: it uses electrical energy from an external source to drive a reaction that would not happen on its own (a reaction with a negative Ecell∘E^\circ_{cell}).

Electrode signs are reversed

The power supply pushes electrons to one electrode and pulls them from the other.

  • Cathode (reduction): connected to the negative terminal of the supply, so it is the negative electrode.
  • Anode (oxidation): connected to the positive terminal, so it is the positive electrode.

Predicting products of electrolysis

For a molten ionic compound the only species present are the cation and anion, so the cation is reduced at the cathode and the anion is oxidised at the anode. Electrolysis of molten sodium chloride gives sodium metal and chlorine gas.

For an aqueous solution water is also present and can compete:

Cathode possibilities: Mn++ne−→Mor2H2O+2e−→H2+2OH−\text{Cathode possibilities: } \text{M}^{n+} + n\text{e}^- \rightarrow \text{M} \quad \text{or} \quad 2\text{H}_2\text{O} + 2\text{e}^- \rightarrow \text{H}_2 + 2\text{OH}^-

Anode possibilities: anion oxidationor2H2O→O2+4H++4e−\text{Anode possibilities: anion oxidation} \quad \text{or} \quad 2\text{H}_2\text{O} \rightarrow \text{O}_2 + 4\text{H}^+ + 4\text{e}^-

The species more easily reduced (more positive potential) wins at the cathode; the species more easily oxidised wins at the anode. So electrolysis of dilute sodium chloride solution tends to give hydrogen and oxygen (from water) rather than sodium and chlorine, though high chloride concentration can favour chlorine at the anode (an overpotential effect noted in some courses).

Industrial applications

  • Electroplating: depositing a thin metal layer (such as silver or chromium) onto an object made the cathode.
  • Electrolytic refining: purifying copper by making impure copper the anode and pure copper the cathode.
  • Extraction of reactive metals: aluminium is extracted by electrolysis of molten aluminium oxide because it is too reactive to reduce chemically.

Why this matters

Electrolysis underpins major industries: aluminium and chlorine production, copper refining, and electroplating. The dot point connects directly to standard electrode potentials (to predict products) and to Faraday's laws (to calculate amounts), and it is the natural contrast to galvanic cells.

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