Q: What does H2 Chemistry Notes: EXTENSION TOPICS, Topic 12 - Electrochemistry cover?
A: Connect electrochemical cells, standard electrode potentials, the Nernst equation, and electrolysis design for the 2026 H2 Chemistry syllabus.
Electrochemistry links thermodynamics to practical energy devices. This note summarises cell conventions, potential calculations, and electrolysis reasoning expected in 2026.
Cell notation: anode on left, cathode on right, double bar for salt bridge. Electrons flow from anode to cathode.
Feasibility criterion:Ecell∘>0 indicates spontaneous reaction under standard conditions.
3 Nernst Equation
For non-standard conditions:
E=E∘−nFRTlnQ
Where:
R=8.314J⋅mol−1⋅K−1
Temperature T in Kelvins K
n electrons transferred
F=96485C⋅mol−1
Q reaction quotient (products/reactants raised to stoichiometric coefficients).
At 298K, the base-10 version:
E=E∘−n0.0592log10Q
Use to compute potential when concentrations differ from standard or to find equilibrium constants (set E=0).
4 Electrochemical vs Electrolytic Cells
Feature
Electrochemical (voltaic)
Electrolytic
Energy conversion
Chemical → electrical
Electrical → chemical
Spontaneity
Spontaneous redox (\(E_\text{cell} > 0\))
Non-spontaneous; requires external power source.
Anode polarity
Negative
Positive
Cathode polarity
Positive
Negative
Remember: oxidation always occurs at anode, reduction at cathode, irrespective of cell type.
5 Electrolysis Considerations
5.1 Factors Affecting Product
Nature of electrodes: Inert or reactive (e.g. copper electrodes dissolve in copper electrorefining).
Concentration: High concentration of halide ions can favour halogen evolution over oxygen in aqueous electrolysis.
Overpotential: Real systems may require higher voltage; choose predictions aligned with data booklet and canonical outcomes.
5.2 Faraday’s Laws
Amount of substance produced is proportional to charge passed: Q=It.
Moles of electrons: ne=FQ.
Relate to stoichiometry of electrode reaction to find mass deposited or gas evolved.
6 Worked Example
Question:
Determine the concentration of CuX2+ in the cell Cu(s)∣CuX2+(aq)∣∣AgX+(aq)∣Ag(s) if the measured cell potential at 298K is 0.78V. Standard potentials: E∘(AgX+/Ag)=+0.80V, E∘(CuX2+/Cu)=+0.34V. Silver-ion concentration is 1.00mol⋅L−1.
Solution:
Identify electrodes: Silver has higher E∘; acts as cathode. Reaction:
Cu(s)+2AgX+(aq)CuX2+(aq)+2Ag(s)
Standard cell potential:Ecell∘=0.80V−0.34V=0.46V.
Apply Nernst (n=2):
0.78=0.46−20.0592log10([AgX+]2[CuX2+])
Rearranging:
log10(1.002[CuX2+])=0.02960.46−0.78=−10.8
[CuX2+]=10−10.8=1.6×10−11mol⋅L−1
Low copper-ion concentration raises potential above standard value.
7 Electrolysis Example
Electrolysis of aqueous NaCl with inert electrodes:
Electrode
Reaction
Considerations
Cathode (reduction)
\(\ce{2H2O + 2e^- -> H2 + 2OH^-}\)
Water reduced instead of \(\ce{Na^+}\) due to higher reduction potential.
Anode (oxidation)
\(\ce{2Cl^- -> Cl2 + 2e^-}\)
High chloride concentration favours \(\ce{Cl2}\) over \(\ce{O2}\).
Reversing sign incorrectly when computing cell potential (always cathode minus anode).
Forgetting to square concentrations when stoichiometric coefficient >1 in Nernst expression.
Assuming Na+ reduces to sodium metal in aqueous electrolysis (water reduces first).
Ignoring electrode material (graphite vs platinum vs copper) when predicting products.
9 Quick Drills
Calculate Ecell∘ for FeX3++eX−FeX2+ (+0.77V) coupled with CeX4++eX−CeX3+ (+1.70V) and predict spontaneity.
A current of 2.50A runs through molten AlX2OX3
Sketch a labelled diagram for a galvanic cell using Zn/ZnX2+ and Cu/CuX2+
Electrochemistry problems tie together thermodynamics, stoichiometry, and practical chemistry. Keep data booklet standard potentials close, practise Nernst calculations, and explore more problems at https://eclatinstitute.sg/blog/h2-chemistry-notes.
