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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.
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
2HX++2eX−HX2
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+
Common exam mistakes
Mistake: Subtracting electrode potentials in the wrong order - Ecell∘=Ecathode∘−Eanode∘; reversing the subtraction gives the wrong sign and an incorrect spontaneity conclusion.
Mistake: Flipping the sign of E∘ for the oxidation half-equation and then subtracting it again - only use standard reduction potentials directly in the formula; do not flip the anode E∘ before applying the cell formula.
Mistake: Predicting that NaX+ or KX+ is reduced at the cathode in aqueous electrolysis - water is preferentially reduced to give HX2
Mistake: Ignoring the effect of electrode material in electrolysis - reactive (non-inert) anodes (e.g. copper) dissolve instead of releasing oxygen; always state whether electrodes are inert or reactive.
Mistake: Forgetting to square or otherwise adjust concentrations in the Nernst expression when stoichiometric coefficients are greater than one - [AgX+]2 appears when two electrons are transferred and the silver half-equation involves 2AgX+
Mistake: Confusing anode polarity between electrochemical and electrolytic cells - the anode is negative in a voltaic cell (spontaneous oxidation) but positive in an electrolytic cell (forced oxidation); oxidation at the anode holds in both cases.
Frequently asked questions
How do I tell which electrode is the anode and which is the cathode? Oxidation always occurs at the anode; reduction always occurs at the cathode - this rule holds for both voltaic and electrolytic cells. In a voltaic cell, identify the half-reaction with the lower (more negative) E∘ as the anode. In an electrolytic cell, the anode is connected to the positive terminal of the external power supply.
When is ClX2 produced at the anode instead of OX2 during aqueous electrolysis? Chlorine is preferentially discharged when the chloride ion concentration is high (e.g. concentrated NaCl solution). At low chloride concentrations, or in dilute solutions, water is oxidised to give oxygen instead. Overpotential considerations also favour chlorine under concentrated conditions.
Does a positive Ecell∘ always mean the reaction will proceed? A positive Ecell∘ indicates spontaneity under standard conditions (298K, 1mol⋅L−1 concentrations, 1bar pressure). Under non-standard conditions, use the Nernst equation to recalculate the cell potential; the sign may change at very different concentrations.
Struggling with Electrochemistry? Our H2 Chemistry tuition programme covers this topic with structured practice, Paper 4 practical drills, and worked exam solutions.
