Part 1: Half-cells, standard potentials and cell notation

Explain how separate redox reactions produce a voltage and represent a cell using the standard hydrogen electrode as a reference.

Keep the electron and ion pathways separate

In a simple metal/metal-ion half-cell, a conducting metal is in contact with a solution containing its ions. Connect two suitable half-cells with an external conductor and a salt bridge. During discharge, oxidation releases electrons at one electrode and reduction consumes them at the other. Electrons travel through the external circuit; ions carry charge through the solutions and salt bridge.

For a zinc–copper cell, zinc loses electrons and dissolves as Zn²⁺ while Cu²⁺ gains electrons and deposits copper. The zinc electrode is the negative terminal and the copper electrode is the positive terminal during spontaneous discharge. Oxidation occurs at the anode and reduction at the cathode. Do not transfer terminal-sign rules uncritically to an externally driven electrolytic cell.

The salt bridge completes the ionic circuit

A suitable soluble electrolyte lets ions move between the half-cells and prevents charge build-up that would stop sustained electron flow. Potassium nitrate is a common choice if compatible with the specific solutions. The bridge’s ions must not precipitate, complex or otherwise react significantly with the cell species. A metal wire is not a substitute for this ionic connection.

As Zn²⁺ forms, anions migrate towards that half-cell; as Cu²⁺ is removed, cations migrate towards the copper half-cell. The exact ion movements depend on the half-reactions. No electrons travel through the salt solution.

A potential needs a reference half-cell

An individual electrode potential cannot be measured on its own. Its standard value is measured relative to the standard hydrogen electrode, assigned 0.00 V. The standard hydrogen electrode uses platinum in contact with hydrogen gas at 100 kPa and an aqueous solution with [H⁺] = 1.00 mol dm⁻³, at 298 K in the AQA convention. Platinum is an inert conductor and catalytic surface; it is not consumed.

Standard electrode potentials also require the other half-cell’s specified dissolved species at 1.00 mol dm⁻³ and any participating gas at 100 kPa. Use ionic concentrations, not automatically the same concentration of every salt formula unit. These are the concentration-based conditions expected at this level.

Read a cell from left to right

A single vertical line marks a phase boundary and a double line represents the salt bridge or liquid junction. Separate species in the same phase with a comma, not another phase boundary. Include a conducting inert electrode such as Pt where the redox species do not supply their own solid conductor.

For spontaneous discharge drawn with the negative electrode on the left and positive electrode on the right, oxidation occurs on the left and reduction on the right. For any specified cell orientation, Ecell = Eright − Eleft using tabulated reduction potentials. Reversing the notation changes the sign of the cell voltage.

Tables list reduction directions

IUPAC electrode-potential tables write half-equations as reductions with electrons on the left. A more positive standard reduction potential indicates a stronger tendency for the oxidised species to accept electrons relative to the reference. A more negative value makes the reduced species a stronger reducing agent in a suitable pairing.

Name the actual chemical species: Ag⁺ is the oxidising agent in Ag⁺ + e⁻ → Ag; Ag metal is its reduced partner. Do not call the whole half-cell an electron donor without specifying the reaction direction.

Quick checks

Original Finesse questions. Reveal the indicative worked solutions after attempting each question; these are not official AQA mark allocations.

Sources