WAChemistrySyllabus dot point

How does the standard electrode potential series let us predict whether a redox reaction will occur and how much voltage it produces?

Use the standard electrode potential series to predict the spontaneity of redox reactions and calculate standard cell potentials

A focused answer to the WACE Year 12 Chemistry dot point on standard electrode potentials, the reference hydrogen electrode, predicting spontaneity, and calculating standard cell EMF from the potential series, with a worked example and common exam mistakes.

Generated by Claude Opus 4.77 min answerUpdated 2026-05-29

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

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What this dot point is asking

The standard electrode potential ( $E^\circ$ ) of a half-reaction measures its tendency to gain electrons (to be reduced) under standard conditions, relative to a reference. Because we cannot measure a single electrode in isolation, all values are quoted against the standard hydrogen electrode, which is assigned a potential of exactly 0 volts.

Reading the series

In the SCSA data booklet the half-reactions are written as reductions and listed by $E^\circ$ . The key interpretations:

A more positive $E^\circ$ means the species on the left is a stronger oxidising agent (it gains electrons readily). Fluorine and permanganate sit near the top.
A more negative $E^\circ$ means the species on the right is a stronger reducing agent (it loses electrons readily). Lithium and potassium sit near the bottom.

Calculating standard cell potential

For a complete cell, identify which half-reaction is reduction (cathode) and which is oxidation (anode), then:

E^\circ_{cell} = E^\circ_{cathode} - E^\circ_{anode}

where both values are taken as the tabulated reduction potentials. A positive $E^\circ_{cell}$ means the reaction is spontaneous (it occurs in a galvanic cell). A negative value means it is non-spontaneous and would need an external power source (electrolysis).

Importantly, $E^\circ$ is an intensive property: it does not depend on the amount of substance, so you never multiply a potential by the coefficients used to balance electrons.

Predicting spontaneity

To predict whether a given oxidising agent will react with a given reducing agent, find both half-reactions in the table. If the oxidising agent's half-reaction is higher (more positive) than the reducing agent's, the reaction is spontaneous; the combined $E^\circ_{cell}$ will be positive.

Calculating a cell potential

Question

A cell is built from a zinc electrode and a copper electrode. Given $E^\circ(\text{Cu}^{2+}/\text{Cu}) = +0.34$ V and $E^\circ(\text{Zn}^{2+}/\text{Zn}) = -0.76$ V, find the standard cell potential and state whether the cell reaction is spontaneous.

Solution

The copper half-reaction has the more positive potential, so copper ions are reduced (cathode) and zinc is oxidised (anode).

E^\circ_{cell} = E^\circ_{cathode} - E^\circ_{anode} = (+0.34) - (-0.76) = +1.10 \text{ V}

Because $E^\circ_{cell}$ is positive, the reaction $\text{Zn} + \text{Cu}^{2+} \rightarrow \text{Zn}^{2+} + \text{Cu}$ is spontaneous, which is why this combination forms a working galvanic cell.

Why this matters

The potential series is the predictive engine of electrochemistry. It tells you which metal will displace another, which way electrons flow in a cell, what voltage a cell delivers, and which products form during electrolysis. These limitations and predictions appear throughout the redox section of the examination.

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