Standard electrode potentials: WACE Year 12 Chemistry

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:

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.

Limitations: thermodynamics, not kinetics

A crucial caveat the best answers include is that $E^\circ$ predicts only whether a reaction is thermodynamically feasible, not whether it will actually proceed at a useful rate. A positive $E^\circ_{\text{cell}}$ says the reaction can happen, but a high activation energy may make it immeasurably slow (hydrogen and oxygen have a very positive cell potential yet do not react without a spark or catalyst). Standard potentials also apply strictly to standard conditions; changing concentration, temperature or using non-standard pressures shifts the real cell voltage away from $E^\circ_{\text{cell}}$. In aqueous electrolysis, an additional overpotential can mean the predicted product is not the one observed, which is why concentrated chloride solutions liberate chlorine rather than oxygen at the anode.

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.