‹ Chemistry simulationsChemistry · A Level 9701 · 24 ElectrochemistrySix simulations covering Topic 24: an electrolysis cell where you choose the electrolyte and watch which ion is discharged at each electrode; Q = It and the Faraday constant turned into a mass of copper or a volume of hydrogen, and the copper experiment that gives the Avogadro constant; the standard hydrogen electrode wired to a second half-cell to measure E°; a cell built from any two half-cells with the electron flow, the positive electrode and the overall equation; a feasibility checker for any oxidising agent against any reducing agent; and the Nernst equation, with E redrawn as the concentrations change.
An electrolyte conducts because its ions can move. Pass a direct current through it and the cations drift to the negative electrode (the cathode) where they gain electrons, while the anions drift to the positive electrode (the anode) where they lose them. In a molten salt there is only one choice at each electrode. In solution, water joins in, and which ion is discharged depends on the electrode potentials and on the concentration.
The amount of product at an electrode depends only on how much charge has passed. Charge is current multiplied by time, one mole of electrons carries one Faraday (96 500 C), and the half-equation tells you how many moles of electrons each mole of product needs. Set the current and time, pick the product, and follow the working. Then run the copper experiment backwards to measure the Avogadro constant.
A single half-cell has no measurable voltage on its own; a voltmeter needs two electrodes. Chemists agreed on a reference: hydrogen gas at 1 atm bubbling over platinum in a solution containing 1.00 mol dm−3 of H+ at 298 K, defined as 0.00 V. The standard electrode potential of any other half-cell is the voltage it gives when connected to this reference under standard conditions.
Connect any two half-cells and the one with the more positive E° becomes the positive electrode: reduction happens there, and it pulls electrons through the wire from the other electrode, where oxidation happens. E°cell = E°(right) − E°(left). If that comes out negative, the electrons are simply flowing the other way; swap the half-cells and the same number appears with a plus sign.
Lay the half-equations out on a ladder of E° values. The species on the left of a half-equation high up the ladder are strong oxidising agents; the species on the right of one low down are strong reducing agents. An oxidising agent can oxidise any reducing agent that sits lower on the ladder, because E°cell for that pairing is positive. Pick a pair and test it, then read the caveats.
E° is the electrode potential at 1.00 mol dm−3. Change the concentrations and the potential shifts: more of the oxidised species makes the half-cell more willing to accept electrons, so E moves positive; more of the reduced species moves it negative. The Nernst equation puts a number on it, E = E° + (0.059 / z) log([oxidised] / [reduced]), where z is the number of electrons in the half-equation.
Six interactive simulations for Cambridge A Level Chemistry 9701 topic 24 Electrochemistry: an animated electrolysis cell for molten and aqueous electrolytes with the half-equations at each electrode and the role of concentration and electrode potential in deciding which ion is discharged, Q = It and the Faraday constant worked through to the mass of metal or volume of gas produced together with the copper electrolysis that gives the Avogadro constant, the standard hydrogen electrode in 3D connected through a salt bridge to a metal or ion-ion half-cell to measure a standard electrode potential, an electrochemical cell built from any two Data Booklet half-cells showing E cell, the direction of electron flow and the overall equation, a feasibility checker comparing an oxidising agent with a reducing agent including the kinetic and non-standard-condition limitations, and the Nernst equation with E plotted against the log of the concentration ratio.
Six simulations covering Topic 24: an electrolysis cell where you choose the electrolyte and watch which ion is discharged at each electrode; Q = It and the Faraday constant turned into a mass of copper or a volume of hydrogen, and the copper experiment that gives the Avogadro constant; the standard hydrogen electrode wired to a second half-cell to measure E°; a cell built from any two half-cells with the electron flow, the positive electrode and the overall equation; a feasibility checker for any oxidising agent against any reducing agent; and the Nernst equation, with E redrawn as the concentrations change.
Electrode potentials and electrolysis is a free, browser-based simulation for Cambridge International AS & A Level Chemistry 9701 (topic 24 Electrochemistry). The course is assessed in Paper 1 (Multiple Choice (AS)), Paper 2 (AS Structured), Paper 3 (Advanced Practical), Paper 4 (A Level Structured), Paper 5 (Planning, Analysis, Evaluation). Syllabus for examination in 2025-2027. There are 6 tabs, each a self-contained experiment with its own controls and live readings, and nothing to install: it runs on a laptop, tablet or classroom projector.
Full syllabus map: A Level Chemistry 9701 topic 24.
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