QCE Chemistry - Unit 3 - Oxidation and reduction

Electrochemical Cells | QCE Chemistry

Understand how electrochemical cells separate redox reactions into half-cells in QCE Chemistry.

Updated 2026-05-18 - 4 min read

An electrochemical cell is a system where a redox reaction is connected to an electric circuit. The main idea is simple: electrons are made to travel through a wire instead of moving directly from the reductant to the oxidant in the same beaker.

This is why electrochemical cells are useful. They let chemical changes produce electrical energy, or they let electrical energy force chemical changes to happen.

The two big cell types

There are two types you need to separate clearly:

• Galvanic cells use a spontaneous redox reaction to produce electrical energy.
• Electrolytic cells use electrical energy to force a non-spontaneous redox reaction.

The confusing part is electrode charge. In a galvanic cell, the anode is negative and the cathode is positive. In an electrolytic cell, the anode is positive and the cathode is negative.

Cell components

Each half-cell contains a redox pair, usually a metal electrode with its ions in solution, such as $\mathrm{Zn(s)}/\mathrm{Zn^{2+}(aq)}$, or an inert electrode with aqueous ions. Inert electrodes such as graphite or platinum conduct electrons but are not consumed by the reaction.

The external wire lets electrons move between electrodes. The internal pathway, usually a salt bridge or electrolyte, lets ions move so charge does not build up.

The electrodes are where electron transfer happens, but they are not always reactants. A zinc electrode in a zinc ion solution participates in the reaction because zinc atoms can leave the metal as ions. A graphite electrode in a solution of ions often just provides a conductive surface. When you draw or label a cell, say what the electrode material is; "electrode" by itself is usually too vague for a strong practical answer.

Why the salt bridge matters

If electrons leave one half-cell and enter another, the solutions quickly become charged unless ions can move. A salt bridge fixes this by allowing spectator ions to migrate.

In a typical galvanic cell:

• anions from the salt bridge move toward the anode, where positive ions are being produced
• cations from the salt bridge move toward the cathode, where positive ions are being consumed
• electrons move through the wire, not through the salt bridge

If the salt bridge is missing, the cell may work only briefly or not at all. At the anode, oxidation can produce excess positive ions in solution. At the cathode, reduction can remove positive ions from solution. Without ion migration, charge separation builds up and opposes further electron flow.

Observations in practical work

Electrochemical cell questions often connect equations to observations. If a metal electrode is being oxidised, atoms leave the electrode as ions, so its mass may decrease. If metal ions are being reduced, metal may plate onto an electrode, so its mass may increase. Bubbles suggest a gas product, such as $\mathrm{H_2(g)}$, $\mathrm{O_2(g)}$ or $\mathrm{Cl_2(g)}$.

When you record observations, separate starting observations from changes during the reaction. A good practical note might include initial colour, electrode appearance, whether the solution is clear or cloudy, whether bubbles form, whether a precipitate appears, and how the final electrode surfaces look. Those details make it much easier to justify which half-reactions occurred.

Direction of movement

There are three movements to keep separate:

• electrons move through the external wire
• cations move through the salt bridge or electrolyte toward the cathode
• anions move through the salt bridge or electrolyte toward the anode

Electrons do not move through the salt bridge, and ions do not move through the wire. If a diagram asks for arrows, label each arrow with the particle moving, not just the direction.

Common mistake

Quick check

Sources

• QCAA Chemistry subject page
• QCAA Chemistry 2025 syllabus
• OpenStax Chemistry 2e: Review of Redox Chemistry
• OpenStax Chemistry 2e: Galvanic Cells