QCE Chemistry - Unit 3 - Oxidation and reduction

Galvanic Cells | QCE Chemistry

Learn how galvanic cells produce electrical energy from spontaneous redox reactions in QCE Chemistry.

Updated 2026-05-17 - 4 min read

Galvanic cells convert chemical energy into electrical energy using a spontaneous redox reaction. They are also called voltaic cells. Batteries are everyday examples, although school diagrams usually show the cell as two separated half-cells.

How a galvanic cell works

If an oxidant and reductant are mixed directly, electrons transfer during collisions and the energy is mostly released as heat. In a galvanic cell, the two half-reactions are separated, so electrons are forced to travel through an external wire.

Electrode signs

In a galvanic cell:

• anode is negative
• cathode is positive
• electrons flow from anode to cathode
• conventional current is opposite to electron flow

Oxidation still occurs at the anode and reduction still occurs at the cathode.

For the zinc-copper cell:

Anode: $\mathrm{Zn(s)} \rightarrow \mathrm{Zn^{2+}(aq)} + 2e^-$

Cathode: $\mathrm{Cu^{2+}(aq)} + 2e^- \rightarrow \mathrm{Cu(s)}$

Overall: $\mathrm{Zn(s)} + \mathrm{Cu^{2+}(aq)} \rightarrow \mathrm{Zn^{2+}(aq)} + \mathrm{Cu(s)}$

The zinc electrode supplies electrons, so it is the negative anode. Copper(II) ions accept electrons, so copper is deposited at the positive cathode.

The cell keeps running only while there is a usable oxidant and reductant and a complete circuit. As zinc atoms oxidise, the zinc electrode can lose mass. As copper(II) ions are reduced, copper metal can plate onto the cathode. At the same time, ion concentrations change, so the voltage may slowly change rather than staying fixed at the standard value.

Salt bridge and observations

The salt bridge completes the circuit by allowing ions to move. In the zinc half-cell, $\mathrm{Zn^{2+}}$ builds up, so anions move toward the anode. In the copper half-cell, $\mathrm{Cu^{2+}}$ is removed from solution, so cations move toward the cathode.

You can often connect this to visible changes:

• anode metal may lose mass as it forms ions
• cathode may gain mass as metal plates onto it
• solution colour may fade if coloured ions are consumed
• voltage may change as concentrations change during the reaction

In a practical task, the salt bridge should contain ions that stay out of the chemistry. Potassium nitrate is a common school example because $\mathrm{K^+}$ and $\mathrm{NO_3^-}$ usually behave as spectators in these cells. If the salt bridge ions react or form a precipitate, the cell is no longer testing only the intended redox reaction.

Drawing a galvanic cell

A high-quality galvanic cell diagram should show:

• the two half-cells and their solutions
• the anode and cathode labels
• electrode polarity
• the direction of electron flow through the wire
• the salt bridge or porous barrier
• ion movement through the salt bridge
• the voltmeter or external circuit
• the relevant half-equations

Cell notation

A zinc-copper galvanic cell can be written as:

$ \mathrm{Zn(s)}\ |\ \mathrm{Zn^{2+}(aq)}\ ||\ \mathrm{Cu^{2+}(aq)}\ |\ \mathrm{Cu(s)} $

The left side is usually the oxidation half-cell. The double line represents the salt bridge.

Rules for reading notation:

• single vertical line $|$ separates phases or interfaces
• double vertical line $||$ represents the salt bridge
• anode is usually written on the left
• cathode is usually written on the right

If a half-cell uses an inert electrode, include the inert conductor in the notation. For example, platinum may be written when all reacting species are aqueous or gaseous and no metal electrode participates directly.

Predicting voltage

For standard galvanic cells:

$ E_{\mathrm{cell}} = E^\circ_{\mathrm{cathode}} - E^\circ_{\mathrm{anode}} $

Use the standard reduction potentials as written in the data book. The more positive reduction potential is more likely to be reduced, so it usually becomes the cathode in a spontaneous galvanic cell.

Common mistake

Quick check

Sources

• QCAA Chemistry subject page
• QCAA Chemistry 2025 syllabus
• OpenStax Chemistry 2e: Galvanic Cells
• OpenStax Chemistry 2e: Electrode and Cell Potentials