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Q: What does H2 Chemistry notes: Electrochemistry (9476) cover? A: Master standard electrode potentials, galvanic and electrolytic cells, cell potential calculations, electrolysis product prediction, and Faraday's laws for the 2026 H2 Chemistry syllabus.
Electrochemistry connects redox chemistry to real-world energy devices and industrial processes. This topic is one of the more calculation-heavy sections of H2 Chemistry and regularly appears in Papers 2 and 3 as structured questions involving cell potential computation, electrolysis reasoning, or data-booklet lookups.
Status: SEAB H2 Chemistry (9476, first exam 2026) syllabus and Chemistry Data Booklet last checked 2026-03-23. [1][2]
Quick revision box
What this topic tests: Standard electrode potentials, cell potential calculations, galvanic vs electrolytic cell identification, electrolysis product prediction, and Faraday's law computations.
Before tackling electrochemistry, ensure you can assign oxidation states reliably and identify which species is oxidised or reduced. Electrochemistry builds directly on the redox concepts from earlier topics.
Key principles to recall:
Oxidation is loss of electrons; reduction is gain of electrons (OIL RIG).
An oxidising agent is itself reduced; a reducing agent is itself oxidised.
Balance half-equations for both atoms and charge before combining.
2 Standard Electrode Potentials
2.1 The standard hydrogen electrode (SHE)
All electrode potentials are measured relative to the SHE, which is assigned E∘=0.00V by convention. The SHE consists of a platinum electrode in contact with HX2 gas at 1bar and HX+ ions at 1.00moldm−3, at 298K.
2.2 Interpreting the electrochemical series
Half-equations in the Data Booklet are written as reductions. A more positive E∘ means a stronger tendency to be reduced, so the species on the left is a stronger oxidising agent.
Half-equation
E∘/V
FX2+2eX−2FX−
+2.87
AgX++eX−Ag
CuX2++2eX−Cu
2HX++2eX−HX2
ZnX2++2eX−Zn
LiX++eX−Li
2.3 Standard conditions
Standard electrode potentials apply at 298K, with all ions at 1.00moldm−3 and all gases at 1bar. Deviations from these conditions shift the measured potential. The SEAB syllabus expects you to predict qualitatively how concentration changes affect electrode potential.
3 Electrochemical (Galvanic) Cells
3.1 How they work
A galvanic cell converts chemical energy into electrical energy through a spontaneous redox reaction. Two half-cells are connected by a salt bridge that maintains electrical neutrality and an external wire through which electrons flow.
Anode: Oxidation occurs here. In a galvanic cell, the anode carries a negative polarity.
Cathode: Reduction occurs here. In a galvanic cell, the cathode carries a positive polarity.
Salt bridge: Allows migration of ions to complete the circuit and prevent charge build-up.
3.2 Cell notation
The conventional way to represent a galvanic cell:
Zn(s)∣ZnX2+(aq)∣∣CuX2+(aq)∣Cu(s)
Read left to right: anode | anode solution || cathode solution | cathode. The double vertical line represents the salt bridge.
3.3 Calculating cell potential
Ecell∘=Ecathode∘−Eanode∘
For the Daniell cell:
Ecell∘=+0.34V−(−0.76V)=+1.10V
A positive Ecell∘ indicates the reaction is thermodynamically feasible under standard conditions.
Important:E∘ values are intensive properties. Do not multiply them by stoichiometric coefficients when calculating cell potentials.
4 Electrolytic Cells
4.1 How they differ from galvanic cells
An electrolytic cell uses an external power source to drive a non-spontaneous redox reaction. The polarity of the electrodes is reversed compared to a galvanic cell.
Feature
Galvanic cell
Electrolytic cell
Energy conversion
Chemical to electrical
Electrical to chemical
Spontaneity
Spontaneous (Ecell>0)
Non-spontaneous; external voltage required
Anode polarity
Negative
Positive
Cathode polarity
Positive
Negative
The one constant: oxidation always occurs at the anode, and reduction always occurs at the cathode.
4.2 Predicting electrolysis products
Three factors determine which species are discharged at each electrode:
Position in the electrochemical series: Species with more positive E∘ values are preferentially reduced at the cathode. At the anode, species with more negative E∘ (for the reverse reaction) are preferentially oxidised.
Concentration of ions: A high concentration of a particular ion can shift the product. For example, concentrated NaCl(aq) favours ClX2 at the anode rather than OX2.
Nature of electrodes: Inert electrodes (graphite, platinum) do not participate. Reactive electrodes (e.g. copper) may dissolve at the anode.
4.3 Worked example: Electrolysis of dilute sulfuric acid
With inert electrodes:
Cathode:2HX+(aq)+2eX−HX2(g) - hydrogen ions are more easily reduced than water.
Anode:2HX2O(l)OX2(g)+4HX+(aq)+4eX−
Volume ratio of HX2:OX2 collected is 2:1.
4.4 Worked example: Electrolysis of concentrated NaCl(aq)
Cathode:2HX2O(l)+2eX−HX2(g)+2OHX−(aq) - water is reduced because NaX+ has a very negative E∘.
Anode:2ClX−(aq)ClX2(g)+2eX−
The solution becomes alkaline near the cathode. This is the basis of the chlor-alkali industry.
5 Faraday's Laws
5.1 Core relationships
The amount of substance deposited or liberated at an electrode is proportional to the charge passed:
Q=I×t
where Q is charge in coulombs (C), I is current in amperes (A), and t is time in seconds (s).
The number of moles of electrons transferred:
ne=FQ
where the Faraday constant F=96485Cmol−1.
5.2 Calculation procedure
Calculate the total charge: Q=It.
Find moles of electrons: ne=Q/F.
Use the stoichiometry of the half-equation to find moles of substance produced.
Convert to mass or volume as required.
5.3 Worked example
A current of 2.50A is passed through CuSOX4(aq) for 30.0min using inert electrodes. Calculate the mass of copper deposited.
Q=2.50×(30.0×60)=4500C
ne=4500/96485=0.04664mol
From CuX2++2eX−Cu
Mass =0.02332×63.5=1.48g
6 Concentration Effects on Electrode Potential
The SEAB syllabus requires you to predict qualitatively how changes in ion concentration shift electrode potentials and cell potentials. When the concentration of the oxidised form increases (or the reduced form decreases), the reduction potential becomes more positive, favouring reduction.
Enrichment - the Nernst equation:
For quantitative treatment (beyond syllabus requirements, but useful for deeper understanding):
E=E∘−nFRTlnQ
At 298K, this simplifies to:
E=E∘−n0.0592log10Q
This equation is not required for exam calculations, but understanding the direction of the shift is examinable. [1]
7 Linking Electrochemistry to Other Topics
Chemical Energetics (Topic 7):ΔG∘=−nFEcell∘ connects cell potential to Gibbs free energy. A positive Ecell∘ corresponds to a negative ΔG∘, confirming spontaneity.
Chemical Equilibria (Topic 9): At equilibrium, Ecell=0 and Q=K. This relationship links the equilibrium constant to the standard cell potential.
Transition Elements (Topic 13): Many electrochemical series entries involve transition metal ions; variable oxidation states are central to electrochemistry.
8 Common Exam Pitfalls
Sign errors: Always apply Ecell∘=Ecathode∘−Eanode∘. Never change the sign of E∘ values from the Data Booklet when using this formula.
Multiplying E∘ by coefficients: Electrode potentials are not multiplied by stoichiometric coefficients. They are intensive properties.
Confusing anode polarity: The anode is negative in a galvanic cell but positive in an electrolytic cell. Oxidation occurs at the anode in both.
Predicting Na metal from aqueous electrolysis:NaX+ has a very negative E∘; water is always reduced first in aqueous solution.
Ignoring electrode material: Reactive electrodes (e.g. copper anode in CuSOX4 electrolysis) dissolve, altering the products.
Forgetting units in Faraday calculations: Keep t in seconds, I in amperes, and state the Faraday constant with units.
9 Quick Retrieval Check
What is the standard hydrogen electrode, and why is it assigned E∘=0.00V?
In the cell Mg(s)∣MgX2+(aq)∣∣FeX2+(aq)∣Fe(s), which electrode is the cathode? Calculate Ecell∘.
Explain why concentrated NaCl(aq) produces ClX2 at the anode, but dilute NaCl(aq)
A current of 1.50A flows for 20.0min through molten NaCl. Calculate the mass of sodium deposited.
Where can I find the full H2 Chemistry Notes series? Start at the H2 Chemistry Notes hub, then follow the topic sequence from Atomic Structure through to Electrochemistry.
Is the Nernst equation examinable? The SEAB 9476 syllabus expects qualitative predictions of how concentration changes affect electrode potential. The Nernst equation itself is enrichment material, not required for exam calculations. [1]
How does electrochemistry link to the practical paper? Paper 4 may include electrochemical cell setups, voltmeter readings, or electrolysis measurements. Familiarity with cell diagrams and data recording is helpful.
