QCE Chemistry - Unit 3 - Oxidation and reduction

Standard Electrode Potential | QCE Chemistry

Learn how standard electrode potentials predict redox direction and galvanic cell voltage in QCE Chemistry.

Updated 2026-05-18 - 4 min read

Standard electrode potentials compare how strongly different species tend to gain electrons under standard conditions. They are measured as reduction potentials and compared with the standard hydrogen electrode, which is assigned $0.00\ \mathrm{V}$.

Reading the table

Electrode potential tables are written as reduction half-equations. That means electrons are normally on the left:

$ \text{oxidised form} + e^- \rightarrow \text{reduced form} $

The more positive the $E^\circ$ value, the stronger the tendency to be reduced. A species with a high reduction potential is a strong oxidant because it causes another species to lose electrons.

For example, if copper(II) has a more positive reduction potential than zinc(II), $\mathrm{Cu^{2+}}$ is more likely to accept electrons than $\mathrm{Zn^{2+}}$. In a zinc-copper galvanic cell, $\mathrm{Cu^{2+}}$ is reduced and $\mathrm{Zn(s)}$ is oxidised.

Strongest oxidant and strongest reductant

The table is not only for voltage calculations. It is also a ranking tool.

For species on the left-hand side of the reduction half-equations, the lower down the table they are, the stronger they are as oxidising agents. They want electrons more strongly. For species on the right-hand side, the higher up the table they are, the stronger they are as reducing agents. They are more willing to donate electrons.

So when a question gives several possible reactants, do not just pick the first metal and ion you recognise. List the species present, identify the strongest oxidant and strongest reductant available, then check that the predicted reaction is spontaneous.

Standard conditions

The symbol $E^\circ$ only applies under standard conditions. In school chemistry, this usually means:

• aqueous solutions are $1.0\ \mathrm{mol\ L^{-1}}$
• gases are at standard pressure
• temperature is $25\,^\circ\mathrm{C}$ or $298\ \mathrm{K}$
• solids are pure

If the experiment uses different concentrations or temperatures, the measured voltage can differ from the calculated standard value.

This is why practical galvanic-cell results rarely match the data-book value perfectly. A voltmeter reading may drift as the reaction proceeds because ion concentrations are changing. Electrodes may have oxide layers, the salt bridge may add resistance, and the solutions may not be exactly $1.0\ \mathrm{mol\ L^{-1}}$. In an investigation, those are not excuses; they are evaluation points.

Calculating cell potential

For a galvanic cell:

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

Use reduction potentials for both values. The cathode is where reduction occurs. If $E_{\mathrm{cell}}$ is positive under standard conditions, the reaction is spontaneous as written.

There is another equivalent method:

$ E^\circ_{\mathrm{cell}} = E^\circ_{\mathrm{reduction}} + E^\circ_{\mathrm{oxidation}} $

If you use this method, you reverse the anode half-equation and change the sign of its potential. Do not mix the two methods in the same line of working.

Predicting whether a reaction occurs

For a spontaneous galvanic reaction under standard conditions, the calculated $E_{\mathrm{cell}}$ should be positive. If the value is negative, the reaction is not spontaneous as written. The reverse reaction may be spontaneous instead, or the reaction may need electrical energy supplied in an electrolytic cell.

When more than two species are present, choose the strongest available oxidant and strongest available reductant, then check that the overall reaction gives a positive $E_{\mathrm{cell}}$.

Common mistake

Quick check

Sources

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