Electrochemical cells and cell potentialsEdexcel International A Level Chemistry: Revision notes

An electrochemical cell joins two half-cells so that a redox reaction drives electrons round an external circuit. Each half-cell contains a species and its ion (or two ions of one element with an inert platinum electrode).

• The more negative half-cell loses electrons: oxidation occurs at this negative electrode.
• Electrons flow through the wire to the more positive half-cell, where reduction occurs.
• A salt bridge (filter paper soaked in potassium nitrate solution) completes the circuit by letting ions move.
• A high-resistance voltmeter measures the cell voltage without drawing current.

For Zn and Cu: Zn(s) → Zn²⁺(aq) + 2e⁻ at the negative electrode, and Cu²⁺(aq) + 2e⁻ → Cu(s) at the positive electrode.

The standard cell potential, E°cell, is the emf of a cell under standard conditions. It is found by combining two standard electrode potentials:

E°cell = E°(positive electrode) − E°(negative electrode), that is, E°(right) − E°(left) in the cell diagram.

Worked example, Zn|Zn²⁺ (−0.76 V) and Cu²⁺|Cu (+0.34 V): E°cell = (+0.34) − (−0.76) = +1.10 V.

A positive E°cell means the cell reaction as written is feasible. Do not multiply E° values by the stoichiometric numbers: for 2Ag⁺ + Cu → 2Ag + Cu²⁺, E°cell = (+0.80) − (+0.34) = +0.46 V, not 1.26 V.

Cell diagrams use conventional notation. A single line | is a phase boundary, and a double line ‖ is the salt bridge.

• Write the oxidation (negative) half-cell on the left and the reduction (positive) half-cell on the right.
• Put the electrode at the outside, then the species in the order they react across the cell.
• For two ions of one element, use a comma and an inert Pt electrode.

Examples:

Zn(s) | Zn²⁺(aq) ‖ Cu²⁺(aq) | Cu(s)

Pt(s) | Fe²⁺(aq), Fe³⁺(aq) ‖ Ag⁺(aq) | Ag(s)

With this arrangement E°cell = E°(right) − E°(left) is positive for a feasible reaction.

Each electrode reaction is an equilibrium, so its potential depends on temperature, concentration and pressure. E° values are only quoted at 298 K, 100 kPa and 1.00 mol dm⁻³, so that every half-cell can be compared fairly.

If conditions change, the equilibrium shifts (Le Chatelier). For Ag⁺ + e⁻ ⇌ Ag, a lower [Ag⁺] shifts the equilibrium left, making the electrode less positive. The measured emf then differs from E°cell.

Practical points: use a high-resistance voltmeter so no current flows, use fresh solutions, clean the electrodes, and use a KNO₃ salt bridge rather than KCl when silver ions are present (AgCl would precipitate).

Standard electrode potentials are also called standard reduction potentials, because every half-equation is written as a reduction: oxidised form + ne⁻ ⇌ reduced form.

Listing half-cells in order of E° gives the electrochemical series:

• Most negative E°: the reduced form (e.g. Mg, Li) is a strong reducing agent.
• Most positive E°: the oxidised form (e.g. Cl₂, MnO₄⁻) is a strong oxidising agent.

A species can reduce any oxidised form that lies below it (more positive). Example: Mg (−2.37 V) reduces Ni²⁺ (−0.25 V), and the cell has E°cell = +2.12 V.