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Q: What does H2 Chemistry Notes: Topic 12 - Electrochemistry cover? A: Connect electrochemical cells, standard electrode potentials, concentration effects (Nernst as optional enrichment), 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.
Status: SEAB H2 Chemistry (9476, first exam 2026) syllabus and Chemistry Data Booklet last checked 2026-01-13.
Quick revision box
What this topic tests: Cell conventions, E° calculations, redox direction, and electrolysis design.
Top mistakes to avoid: Sign errors in E°cell; inconsistent oxidation-state accounting; weak electrolysis product justification.
20-minute sprint plan: 5 min redox + sign recap; 10 min electrochemical calculations; 5 min electrolysis product reasoning.
1 Standard Electrode Potentials
Electrode reactions are written as reductions. E∘ values measure tendency to be reduced relative to the standard hydrogen electrode (SHE).
Half-equation
E∘/V
CuX2++2eX−Cu
+0.34V
ZnX2++2eX−Zn
−0.76V
2HX++2eX−HX2
0.00V
More positive E∘ indicates stronger oxidising agent (read reaction forward).
Use the standard potential values in the SEAB Data Booklet at 298K when answering exam questions.
2 Cell Potentials
For cell Zn(s)∣ZnX2+(aq)∣∣CuX2+(aq)∣Cu(s):
Ecell∘=Ecathode∘−Eanode∘=0.34V−(−0.76V)=1.10V
Cell notation: anode on left, cathode on right, double bar for salt bridge. Electrons flow from anode to cathode.
Feasibility criterion:Ecell∘>0 indicates a spontaneous reaction under standard conditions.
3 Nernst Equation (optional enrichment)
SEAB’s syllabus expects you to predict qualitatively how electrode potential changes with concentration. The Nernst equation below is a useful enrichment tool when you want to practise that concentration effect quantitatively.
For non-standard conditions:
E=E∘−nFRTlnQ
Where:
R=8.31J⋅K−1⋅mol−1
Temperature T in Kelvins K
n electrons transferred
F=9.65×104C⋅mol−1
Q reaction quotient (products/reactants raised to stoichiometric coefficients).
At 298K, you may see the base-10 version:
E=E∘−nF2.303RTlog10Q≈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 (Ecell>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([AgX+]2[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)
2HX2O+2eX−HX2+2OHX−
Water reduced instead of NaX+ due to higher reduction potential.
Water: cathode reactant in aqueous electrolysis half-equation balancing.
Chlorine: diatomic anode product when chloride oxidation is favoured.
Representing these species explicitly helps with half-equation balancing and with identifying which gaseous products to test at each electrode.
8 Common Misconceptions
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
Using SEAB Data Booklet values, calculate Ecell∘ for FeX3++eX−FeX2+ (+0.77V) coupled with IX2+2eX−2IX− (+0.54V). Predict spontaneity and write the overall balanced redox equation.
A current of 2.50A runs through molten AlX2OX3
Sketch a labelled diagram for a galvanic cell using Zn/ZnX2+ and Cu/CuX2+
