QCE Chemistry - Unit 3 - Oxidation and reduction
Electrolytic Cells | QCE Chemistry
Understand how electrolytic cells use electrical energy to force non-spontaneous redox reactions in QCE Chemistry.
Updated 2026-05-17 - 5 min read
QCAA official coverage - Chemistry 2025 v1.3
Exact syllabus points covered
- Identify that electrolytic cells use an external electrical potential difference to provide the energy to allow a non-spontaneous redox reaction to occur.
- Identify the essential components of an electrolytic cell, including source of electric current and conductors, positive and negative electrodes, and the electrolyte.
- State the factors that affect the products in electrolysis.
- Determine the products of the electrolysis of a molten salt.
- Explain the products of the electrolysis of aqueous solutions, e.g. dilute and concentration sodium chloride(aq) and copper sulfate(aq).
- Describe that electrolytic cells can be used in small-scale and industrial situations, including metal plating and the purification of copper.
- Calculate moles of electrons, current, time, mass of substance or volume of gas produced or used during electrolysis. (Formula: q= 𝑛(e−)×F or q =I×t).
- Analyse data to determine the relative amounts of product produced at each electrode in electrolysis.
Electrolytic cells use electrical energy to drive a non-spontaneous redox reaction. Instead of a reaction producing a voltage, an external power supply pushes electrons in the direction needed to force the reaction.
Electrolysis is used in electroplating, metal purification, extraction of reactive metals, and decomposition of molten or aqueous ionic substances.
Original Sylligence diagram for electrolysis cell.
Electrode signs
In an electrolytic cell:
- anode is positive
- cathode is negative
- oxidation occurs at the anode
- reduction occurs at the cathode
The power supply pulls electrons away from the anode and pushes electrons onto the cathode.
Predicting products
The first question is whether the electrolyte is molten or aqueous.
For molten salts, only the ions from the salt are present and mobile. For molten sodium chloride:
Cathode: $\mathrm{Na^+(l)} + e^- \rightarrow \mathrm{Na(l)}$
Anode: $2\mathrm{Cl^-(l)} \rightarrow \mathrm{Cl_2(g)} + 2e^-$
Overall: $2\mathrm{NaCl(l)} \rightarrow 2\mathrm{Na(l)} + \mathrm{Cl_2(g)}$
For aqueous solutions, water is also present and can be oxidised or reduced. This means the expected products are not always the metal and non-metal from the salt.
For aqueous sodium chloride, sodium ions are very difficult to reduce, so water is reduced instead:
Cathode: $2\mathrm{H_2O(l)} + 2e^- \rightarrow \mathrm{H_2(g)} + 2\mathrm{OH^-(aq)}$
At the anode, dilute chloride solutions may form oxygen from water, while concentrated chloride solutions are more likely to form chlorine gas. QCE questions usually give enough context for you to decide.
For aqueous electrolysis, product prediction depends on the competing species at each electrode. At the cathode, metal ions compete with water for reduction. Very reactive metal ions such as $\mathrm{Na^+}$ and $\mathrm{K^+}$ are usually not reduced in water; water is reduced instead to form hydrogen gas. At the anode, halide ions can compete with water. Dilute halide solutions often produce oxygen from water, while concentrated halide solutions are more likely to produce the halogen.
That concentration detail is easy to overlook. If a question says "concentrated sodium chloride", expect chlorine gas at the anode. If it says "dilute sodium chloride", oxygen may be the more appropriate product unless the data supplied says otherwise.
Molten versus aqueous
Solid ionic compounds do not conduct well because their ions are locked in a lattice. Molten ionic compounds conduct because the ions are mobile. Aqueous ionic compounds also conduct because the ions can move through water.
This gives three different cases:
- solid salt: ions are present but not mobile, so electrolysis does not proceed normally
- molten salt: only salt ions are mobile, so the salt ions form the products
- aqueous salt: salt ions and water are present, so water may be part of the reaction
Electroplating
Electroplating uses electrolysis to coat an object with a thin layer of metal. The object being plated is connected as the cathode because metal ions need to gain electrons and deposit as solid metal.
For silver plating:
Cathode: $\mathrm{Ag^+(aq)} + e^- \rightarrow \mathrm{Ag(s)}$
If a silver anode is used, it can replace the silver ions being removed from solution:
Anode: $\mathrm{Ag(s)} \rightarrow \mathrm{Ag^+(aq)} + e^-$
The cathode gains mass as the metal layer forms.
Electroplating setup matters. The object to be coated must be the cathode, because the metal ions need to gain electrons on its surface. The electrolyte should contain ions of the plating metal. If a metal anode is used, it can dissolve to replace metal ions removed from solution, keeping the ion concentration more stable.
Comparing electrolytic and galvanic cells
Electrolytic and galvanic cells use the same redox language, but the energy direction is different.
- Galvanic: chemical energy -> electrical energy; spontaneous; $E_{\mathrm{cell}}$ positive under standard conditions.
- Electrolytic: electrical energy -> chemical energy; non-spontaneous; requires a power supply.
- Galvanic anode: negative.
- Electrolytic anode: positive.
- In both: anode is oxidation and cathode is reduction.