Electrochemistry
Study H2 Chemistry electrode potentials, cells and electrolysis, balancing redox equations and relating cell voltage to reaction feasibility.
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Before you begin
This topic measures how strongly redox couples gain electrons, and uses those electrode potentials to predict which reactions are feasible, how much energy they release and what electrolysis produces. It starts with the standard hydrogen electrode, builds cells and overall equations from half-equations, links cell potentials to Gibbs energy, then applies the ideas to electrolysis, batteries and fuel cells.
Be comfortable with: Gibbs energy and feasibility, Le Chatelier’s principle, the mole, gas calculations, and assigning oxidation numbers.
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 electrolysis products in molten and aqueous electrolytes, use Faraday’s law to calculate yield, and explain anodising and copper purification.
- Batteries and fuel cellsUse electrode potentials to explain how batteries and fuel cells work.
Practise and check
Or choose
Topic reference
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 half-cell to measure | connect it to the SHE | 298 K, 1 bar, 1.00 mol dm⁻³; platinum if no solid metal |
| two electrode potentials | the more positive couple is reduced | E⦵_cell = E⦵_reduction-E⦵_oxidation |
| a reaction to test | keep its direction | positive means feasible under standard conditions |
| an overall equation to build | reverse one half-equation, match electrons | never multiply E⦵ |
| Δ G⦵ from a cell | Δ G⦵ = -nFE⦵_cell | n from the balanced equation; answer in J mol⁻¹ first |
| a combined half-equation | add the Δ G⦵ values, not the potentials | convert back with the new n |
| a concentration change | shift the equilibrium Ox + ne⁻ ⇌ Red | more oxidised form makes E more positive |
| an aqueous electrolyte | compare the ions with water | concentration decides chloride versus oxygen |
| a current and time | Q = It, then n(e⁻) = Q/F | time in seconds; ratio from the half-equation |
Core knowledge to remember
- Electrode potentials are reduction potentials measured against the standard hydrogen electrode, E⦵ = 0.00 V, at 298 K, 1 bar and 1.00 mol dm⁻³.
- A more positive E⦵ means the oxidised form is a stronger oxidising agent.
- E⦵_cell = E⦵_reduction-E⦵_oxidation; a positive value means feasible under standard conditions, but says nothing about rate.
- In a cell producing a current, the anode is negative; in electrolysis, the anode is positive. Oxidation is always at the anode.
- Δ G⦵ = -nFE⦵_cell and F = Le = 9.65 × 10⁴ C mol⁻¹.
- Increasing the concentration of the oxidised form of a couple makes its electrode potential more positive.
- In aqueous electrolysis with inert electrodes: reactive metals give hydrogen at the cathode; sulfate gives oxygen at the anode; dilute chloride gives mainly oxygen and concentrated chloride gives chlorine.
- New cells aim to be smaller, lighter and higher in voltage; the hydrogen–oxygen fuel cell gives 1.23 V and water.
Common mistakes
- Calling data-booklet values oxidation potentials.
- Multiplying E⦵ when a half-equation is scaled.
- Reversing the subtraction for E⦵_cell, or reversing a reaction you were asked to test.
- Treating a positive cell potential as proof that a reaction is fast.
- Taking n from an unbalanced half-equation in Δ G⦵ = -nFE⦵, or mixing J and kJ.
- Adding electrode potentials to find the potential of a combined half-equation.
- Using minutes in Q = It, or assuming one mole of electrons gives one mole of product.
- Predicting aqueous electrolysis as if water were absent.
- Describing anodising or copper purification without the electrode reactions.