Electrochemistry
Electrode potentials, electrolysis, cells and redox feasibility.
Before you begin
Electrode potentials, electrolysis, cells and redox feasibility.
Learning goals
- Standard Electrode Potentials and the SHE
- Cell Potentials and Spontaneity
- Writing Redox Equations from Half-equations
- ΔG = −nFE
- Electrolysis Predictions and Faraday’s Law
- Batteries and Fuel Cells
Syllabus statements covered
- define the terms: — standard electrode (redox) potential
- define the terms: — standard cell potential
- describe the standard hydrogen electrode
- describe methods used to measure the standard electrode potentials of: — metals or non-metals in contact with their ions in aqueous solution
- describe methods used to measure the standard electrode potentials of: — ions of the same element in different oxidation states
- calculate a standard cell potential by combining two standard electrode potentials
- use standard cell potentials to: — explain/deduce the direction of electron flow from a simple cell
- use standard cell potentials to: — predict the spontaneity of a reaction
- understand the limitations in the use of standard cell potentials to predict the spontaneity of a reaction
- predict qualitatively how the value of an electrode potential varies with the concentration of the aqueous ion
- describe and explain redox processes in terms of electron transfer and/or of changes in oxidation number (oxidation state)
- construct redox equations using the relevant half-equations (see also Section 13)
- state and apply the relationship ∆G⦵ = −nFE⦵ to electrochemical cells, including the calculation of E⦵ for combined half reactions
- state the relationship, F = Le, between the Faraday constant, the Avogadro constant and the charge on the electron
- predict the identity of the substance liberated during electrolysis from the state of electrolyte (molten or aqueous), position in the redox series (electrode potential) and concentration
- calculate: — the quantity of charge passed during electrolysis
- calculate: — the mass and/or volume of substance liberated during electrolysis
- explain, in terms of the electrode reactions, the industrial processes of: — the anodising of aluminium
- explain, in terms of the electrode reactions, the industrial processes of: — the electrolytic purification of copper [technical details are not required]
- state the possible advantages of developing other types of cell, e.g. the H2/O2 fuel cell and improved batteries (as in electric vehicles) in terms of smaller size, lower mass and higher voltage
Lessons
Work through them in order.
- Standard Electrode Potentials and the SHEDefine standard electrode potential and measure a half-cell against the standard hydrogen electrode.
- Cell Potentials and SpontaneityCalculate a cell potential and use it to predict whether a reaction is feasible.
- Writing Redox Equations from Half-equationsCombine two half-equations into a balanced redox equation for the feasible direction.
- ΔG = −nFEUse ΔG = −nFE to link cell potential, electrons transferred and maximum electrical work.
- Electrolysis Predictions and Faraday’s LawPredict products, calculate yield and explain industrial cells.
- Batteries and Fuel CellsUse electrode potentials to explain how batteries and fuel cells work.
Practise and check
Recommended nextElectrochemistry topic check
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Topic reference
Electrochemistry links redox equations to electrode potentials, cell voltages, Gibbs energy and electrolysis. The lessons build the half-cell and cell model first, then apply it to redox equations, energy, electrolysis and practical devices.
Be comfortable with:
- Redox and Half-equations: electron transfer, oxidation number and charge balance.
- Mole and Avogadro Constant: amount of substance and particle counting.
- Entropy and Gibbs Free Energy: feasibility and the meaning of Δ G.
- Gaseous State: gas volumes when an electrolysis product is measured away from room conditions.
Before you start, be able to convert minutes to seconds and J to kJ, and identify which species gains or loses electrons.
Quick Reference
Galvanic cell vs electrolysis cell
Two-panel comparison of electrode signs, electron direction, ion direction, and oxidation-reduction labels.
Galvanic (voltaic): spontaneous
Anode: negative, oxidation
Cathode: positive, reduction
e- flow: anode -> cathode (external circuit)
Ion flow: cations -> cathode, anions -> anode
Electrolysis: non-spontaneous
Anode: positive, oxidation
Cathode: negative, reduction
e- flow: anode -> cathode in the cell path
Power supply drives the reaction direction
Exam rule: never swap the definitions - anode = oxidation, cathode = reduction.
| Question evidence | First move | Essential check |
|---|---|---|
| one unknown half-cell | compare it with the SHE | 298 K, 1 bar, 1.0 mol dm⁻³ |
| two reduction potentials | more positive value is reduced | E⦵_cell = E⦵_cathode-E⦵_anode |
| overall redox equation | reverse one half-equation and equalise electrons | never multiply E⦵ |
| feasibility claim | inspect the sign of E⦵_cell | say “as written under standard conditions”; rate is separate |
| Gibbs-energy data | find n from the balanced reaction | use J mol⁻¹ in Δ G⦵ = -nFE⦵ |
| electrolysis current and time | Q = It, then n(e⁻) = Q/F | time in s; use the half-equation ratio |
| aqueous electrolyte | list water and solute competitors | use potential/concentration evidence given |
| industrial cell | identify both electrode reactions | link them to oxide growth or copper transfer |
F = Le connects the charge per mole of electrons, the Avogadro constant and the charge on one electron.
Common Exam Traps
- Calling data-booklet values oxidation potentials instead of reduction potentials.
- Multiplying E⦵ when a half-equation is scaled.
- Reversing the subtraction in E⦵_cell.
- Treating a positive standard cell potential as proof of a fast visible reaction.
- Ignoring a non-standard concentration or gas pressure.
- Choosing n from one unbalanced half-equation in Δ G⦵ = -nFE⦵.
- Mixing J mol⁻¹ and kJ mol⁻¹ or minutes and seconds.
- Predicting aqueous electrolysis as if water were absent.
- Swapping anode/cathode definitions when the electrode signs change.
- Describing anodising or copper purification without the electrode reactions.