Standard electrode potentials and the standard hydrogen electrode, the electrochemical series, calculating cell potentials, using electrode potentials to predict the feasibility of redox reactions, and the chemistry of electrochemical cells and fuel cells.

What this dot point is asking

CCEA wants you to define standard electrode potentials and the standard hydrogen electrode, use the electrochemical series, calculate cell potentials, use electrode potentials to predict the feasibility of redox reactions, and describe electrochemical cells and fuel cells.

Half-cells and the need for a reference

A redox reaction can be split into two half-reactions, one oxidation and one reduction, each carried out in its own half-cell. A typical metal/metal-ion half-cell is a metal rod dipping into a solution of its ions; an ion/ion half-cell (such as $\text{Fe}^{3+}/\text{Fe}^{2+}$) uses an inert platinum electrode dipping into a solution of both ions. Each half-cell sets up its own electrode potential, but a potential is only meaningful as a difference between two points, so a single half-cell cannot be measured in isolation. Chemists agree a reference, the standard hydrogen electrode, against which every other half-cell is measured.

The standard hydrogen electrode is hydrogen gas at $100\ \text{kPa}$ bubbled over a platinised platinum electrode in $1\ \text{mol dm}^{-3}$ $\text{H}^+$ at $298\ \text{K}$. The half-cell being measured is connected to it through a high-resistance voltmeter and a salt bridge, and the reading is the standard electrode potential of that half-cell.

The electrochemical series and cell potential

A half-cell with a very positive $E^{\ominus}$ (such as $\text{F}_2/\text{F}^-$ at $+2.87\ \text{V}$) contains a strong oxidising agent that is readily reduced. A half-cell with a very negative $E^{\ominus}$ (such as $\text{Li}^+/\text{Li}$ at $-3.04\ \text{V}$) contains a strong reducing agent. When two half-cells are joined, the more positive one runs as the reduction (positive electrode) and the more negative one is forced into oxidation (negative electrode). A salt bridge completes the circuit by allowing ions to flow and balancing charge, while electrons flow through the external wire from the negative to the positive electrode.

Predicting feasibility

A positive cell potential indicates a feasible reaction, although it does not guarantee a fast one. Feasibility is a thermodynamic statement (the reaction is energetically favourable, since $\Delta G^{\ominus} = -nFE^{\ominus}_{\text{cell}}$ is negative when $E^{\ominus}_{\text{cell}}$ is positive); a reaction can be feasible but immeasurably slow if it has a high activation energy. The predictions also assume standard conditions, so changing concentration or temperature can shift a value far enough to alter the outcome, especially when two potentials are close.

Cells and fuel cells

Electrochemical cells convert chemical energy to electrical energy. In a simple cell the spontaneous redox reaction drives electrons through the external circuit, and the cell runs until the reactants are used up. A fuel cell is different: it generates a voltage continuously while fuel (such as hydrogen) and oxygen are supplied from outside, so it never needs recharging. In the hydrogen fuel cell, hydrogen is oxidised at the negative electrode and oxygen is reduced at the positive electrode; in alkaline conditions the electrode reactions are $\text{H}_2 + 2\text{OH}^- \rightarrow 2\text{H}_2\text{O} + 2e^-$ and $\frac{1}{2}\text{O}_2 + \text{H}_2\text{O} + 2e^- \rightarrow 2\text{OH}^-$, giving an overall reaction $\text{H}_2 + \frac{1}{2}\text{O}_2 \rightarrow \text{H}_2\text{O}$. The only product is water, so fuel cells are attractive for clean power, though making and storing the hydrogen has its own energy and safety costs.

Examples in context

Electrode potentials explain everyday electrochemistry. A car battery (lead-acid) and a portable lithium-ion cell both rely on a positive cell potential to drive current, and the electrochemical series predicts which metal corrodes when two are in contact: in galvanising, zinc ($E^{\ominus} = -0.76\ \text{V}$) protects iron ($E^{\ominus} = -0.44\ \text{V}$) because the more negative zinc is oxidised in preference, acting as a sacrificial anode. The same reasoning tells us chlorine ($E^{\ominus} = +1.36\ \text{V}$) can oxidise bromide to bromine but cannot oxidise fluoride, which CCEA links to the trend in halogen reactivity. Hydrogen fuel cells power some buses and were used on spacecraft, producing drinkable water as a by-product.

Try this

Q1. State the value assigned to the standard hydrogen electrode. [1 mark]

• Cue. $0\ \text{V}$.

Q2. State the condition on cell potential for a redox reaction to be feasible. [1 mark]

• Cue. The cell potential must be positive.