Thermodynamics and Electrochemistry

Electrochemical Cells — AP CHEMISTRY Revision Notes & Topical Questions

5 min read 10 parts Lesson 2 of 15 Unit quiz

0%read

Lesson summary

Electrochemical Cells — AP CHEMISTRY notes: Describe voltaic and electrolytic cells, cell potential, free energy, the Nernst equation and Faraday's law.FeatureVoltaic (galvanic)ElectrolyticReactionSpontaneousNon-spontaneous, driven by an applied voltageAnodeOxidation, negativeOxi…

Salt bridge
Electron flow
Cell notation
Standard conditions
Calculation
Feasibility

Section 1 of 10Cell types

Feature	Voltaic (galvanic)	Electrolytic
Reaction	Spontaneous	Non-spontaneous, driven by an applied voltage
Anode	Oxidation, negative	Oxidation, positive
Cathode	Reduction, positive	Reduction, negative
Energy	Chemical to electrical	Electrical to chemical

Section 2 of 10Voltaic cell parts

1Half-cell

Half-cell a metal electrode in a solution of its ions

2Salt bridge

Salt bridge completes the circuit with an inert electrolyte such as KNO₃

3Electron flow

Electron flow from anode to cathode through the external circuit

4Cell notation

Zn(s) | Zn²⁺(aq) || Cu²⁺(aq) | Cu(s); anode on the left, cathode on the right

Section 3 of 10Standard cell potential

Standard conditions

1.00 M, 298 K and 100 kPa; the standard hydrogen electrode is 0.00 V

Calculation

E°_cell = E°_cathode − E°_anode
What the symbols mean: E_cell = standard cell potential (V) · E_right = electrode potential of the positive (reduction) half-cell · E_left = potential of the negative half-cell (V)
Cu²⁺/Cu (+0.34 V) with Zn²⁺/Zn (−0.76 V) gives +1.10 V.

Feasibility

Feasibility a reaction is favoured when E°_cell is positive

Section 4 of 10Cell potential and free energy

Equation

ΔG° = −nFE°_cell, with F = 96 485 C mol^-1

Signs

E° > 0 gives ΔG° < 0; E° < 0 gives ΔG° > 0

Electrolytic cells

Electrolytic cells an applied voltage greater than |E°| is needed;
- overpotential and resistance raise the practical value
Water electrolysis needs more than 1.23 V.

Section 5 of 10Nonstandard conditions

Nernst equation

E = E° − (RT/nF) ln Q, or E = E° − (0.0592/n) log Q at 298 K

Q < 1

E is larger than E°; Q > 1 — E is smaller than E°

Equilibrium

Q = K and E = 0, so the cell is dead

Concentration cell

Concentration cell same half-reaction at two concentrations;
- oxidation at the dilute side and reduction at the concentrated side
0.50 M and 1.00 M Zn²⁺: E = −(0.0592/2) log 0.5 = +0.0089 V.

Section 6 of 10Electrolysis and Faraday's law

Charge

Q = It, in coulombs
What the symbols mean: Q = charge (C) · I = current (A) · t = time (s)

Moles of electrons

Moles of electrons n = Q/F

Mass

Mass find moles of electrons, use the half-equation ratio, then multiply by molar mass
2.20 A for 15 min gives 1980 C and 0.25 g of Mg (2 mol e⁻ per Mg).

Section 7 of 10Electrolysis of molten salts

Cathode

Cathode positive ions are reduced (metal formed)

Anode

Anode negative ions are oxidised
- (non-metal formed)

Example

Example molten PbBr₂: Pb²⁺ + 2e⁻ → Pb and 2Br⁻ → Br₂ + 2e⁻

Section 8 of 10More from the specification

Faraday constant

The charge of one mole of electrons, 96 500 C mol^-1.

Reaction quotient

The ratio of product to reactant concentrations at any moment.

Section 9 of 10Key terms

Term	Meaning
Voltaic cell	Cell producing electrical energy from a spontaneous redox reaction.
Electrolytic cell	Cell using electrical energy to drive a non-spontaneous reaction.
Standard cell potential	E°_cell = E°_cathode − E°_anode.
Nernst equation	E = E° − (RT/nF) ln Q.
Faraday constant	96 485 C per mole of electrons.
Salt bridge	Ion-containing bridge completing the circuit.

Section 10 of 10Exam technique

Confusing the anode in the two cell types — oxidation is always at the anode, but its sign differs.

Doubling E° when doubling the equation — E° is intensive and does not change.

Using concentrations of solids in Q.

Forgetting to convert minutes to seconds in Q = It.

Last updated: 14 October 2026