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Better Equipped Practical Teaching Guides
An enhanced A Level Chemistry practical guide to constructing electrochemical cells, measuring signed cell potentials, interpreting polarity and electron flow, calculating standard cell potentials, writing cell notation, troubleshooting results and evaluating uncertainty.
Teacher note: This resource is a practical support guide to accompany laboratory equipment. Teachers should adapt procedures and risk assessments to suit curriculum requirements, examination-board specifications and local laboratory policies.
Students construct a zinc-copper electrochemical cell from two half-cells. A zinc electrode is placed in zinc sulfate solution and a copper electrode is placed in copper(II) sulfate solution. The solutions are connected by an inert salt bridge and the electrodes are connected to a high-resistance digital voltmeter.
Students record the signed cell potential, identify the direction of electron flow, compare the experimental result with a value calculated from standard electrode potentials, write the associated half-equations and conventional cell notation, and evaluate why an experimental value may differ from a tabulated standard value.
The core method can be extended by comparing several metal/metal-ion half-cells, using an inert electrode with an ion/ion redox couple, or investigating how ion concentration affects cell potential.
A half-cell contains a redox couple in contact with an electron conductor. In a common metal/metal-ion half-cell, a metal electrode is immersed in a solution containing ions of that metal. Oxidation and reduction can occur in opposite directions at the metal-solution interface, producing an interfacial potential difference.
An isolated electrode potential cannot be measured absolutely. Only the potential difference between two half-cells can be measured, so individual standard electrode potentials are defined relative to a reference electrode.
Oxidation occurs. Electrons are released into the external circuit. In a spontaneous Zn/Cu cell, zinc is the negative anode.
Reduction occurs. Electrons arrive from the external circuit and are accepted by the oxidising species. In a spontaneous Zn/Cu cell, copper is the positive cathode.
Ions move through the bridge to complete the ionic circuit and prevent rapid charge build-up in either half-cell. Electrons travel through the external wire, not through the salt bridge.
Memory aid: AN OX, RED CAT - oxidation at the ANode; REDuction at the CAThode.
Anode: Zn(s) → Zn2+(aq) + 2e-
Cathode: Cu2+(aq) + 2e- → Cu(s)
Overall: Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)
Electron flow: zinc electrode → copper electrode.
E°cell = E°cathode - E°anode
Zn2+/Zn: E° = -0.76 V
Cu2+/Cu: E° = +0.34 V
E°cell = +0.34 - (-0.76) = +1.10 V
The standard hydrogen electrode is assigned E° = 0.00 V and provides the reference against which standard electrode potentials are defined.
At A Level, standard conditions are normally treated as 298 K, gases at 100 kPa and aqueous ions at 1.00 mol dm-3. More strictly, standard electrode potentials refer to species at unit activity.
Electrode-potential tables are written as reduction equilibria. The more positive E° value represents the greater thermodynamic tendency for the oxidised form to gain electrons and undergo reduction.
For a spontaneous standard cell, the half-reaction with the more positive E° remains as reduction at the cathode; the other half-reaction is reversed to give oxidation at the anode.
Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s)
Important limitation: a positive E°cell indicates thermodynamic feasibility under standard conditions. It does not guarantee that the reaction will be rapid; activation energy, passivation and other kinetic effects may prevent an observable reaction.
| Item | Suggested specification / quantity | Purpose or note |
|---|---|---|
| Digital voltmeter | High input resistance; 0-2 V or auto-range; preferably 0.001 V resolution | Measures the cell potential under near-open-circuit conditions while drawing very little current. |
| Connecting leads | Two, with insulated crocodile clips | Connect each electrode to the meter; keep clips above solution level. |
| Beakers | 2 x 50 cm3 or 100 cm3 | One for each half-cell. |
| Measuring cylinder or pipettes | Suitable for measuring or transferring 25.0 cm3 portions | Transfers equal solution volumes. |
| Zinc strip | One clean strip, approximately 10-15 mm wide | Zinc electrode. |
| Copper strip | One clean strip, approximately 10-15 mm wide | Copper electrode. |
| Abrasive paper | Fine emery paper or sandpaper | Removes oxide and contamination immediately before use. |
| Filter-paper salt bridge | Approximately 12 cm x 1-2 cm | Soak thoroughly in an inert electrolyte. |
| Thermometer | Resolution 0.5-1 °C | Records solution temperature. |
| Deionised water and tissues | Small wash bottle and lint-free tissue | Rinses and blots electrodes after cleaning. |
| Waste container | Labelled for metal-ion solutions | Prevents inappropriate sink disposal. |
Note: 0.40 mol dm-3 is a robust classroom concentration, but it is not the 1.00 mol dm-3 concentration commonly used in the examination-board treatment of standard conditions.
Compatibility warning: do not use a chloride-containing salt bridge with Ag+(aq), because chloride ions form insoluble AgCl and change the half-cell composition. The bridge electrolyte must not react significantly with the electrode materials or half-cell solutions.
| Component | Correct position or connection | Why it matters |
|---|---|---|
| Zinc half-cell | Zinc strip immersed in ZnSO4(aq) | Matches the Zn2+/Zn redox couple. |
| Copper half-cell | Copper strip immersed in CuSO4(aq) | Matches the Cu2+/Cu redox couple. |
| Salt bridge | Both wet ends immersed; central section bridges the beaker rims | Completes the ionic circuit without directly mixing the bulk solutions. |
| Black meter lead | Connected to zinc for the expected spontaneous direction | Zinc is the lower-potential, negative electrode. |
| Red meter lead | Connected to copper | The meter should show a positive value close to 1.1 V under suitable conditions. |
| Crocodile clips | Attached to clean, dry metal above the liquid | Avoids contamination and corrosion of the clips. |
| Electrode immersion | Similar depth in each solution | Improves reproducibility and stabilisation; open-circuit EMF is not set by electrode area. |
Do not simply "make the number positive". A digital voltmeter displays V(red) - V(black). A negative reading usually means the red lead is connected to the lower-potential electrode. Record the sign and connections before changing the leads.
Zinc atoms lose electrons and enter solution as hydrated Zn2+ ions:
Zn(s) → Zn2+(aq) + 2e-
Hydrated Cu2+ ions gain electrons and become copper atoms:
Cu2+(aq) + 2e- → Cu(s)
Anions migrate towards the zinc compartment as oxidation increases positive charge there. Cations migrate towards the copper compartment as Cu2+ is removed by reduction. This helps maintain approximate electrical neutrality.
Why visible changes may be small: a high-resistance voltmeter draws only a tiny current, so very little zinc dissolves or copper is deposited. A low-resistance load would permit more current and introduce stronger effects from internal resistance, polarisation and changing concentrations.
| Trial | ZnSO4 / mol dm-3 | CuSO4 / mol dm-3 | Temp. / °C | Red lead | Black lead | Stable Ecell / V |
|---|---|---|---|---|---|---|
| 1 | ||||||
| 2 | ||||||
| 3 | ||||||
| Mean | ||||||
| Range |
| Trial | E at 10 s / V | E at 20 s / V | E at 30 s / V | E at 40 s / V | Chosen stable value / V |
|---|---|---|---|---|---|
| 1 | |||||
| 2 | |||||
| 3 |
| Cell notation | Red lead electrode | Black lead electrode | Signed Ecell / V | Anode | Cathode | Electron direction |
|---|---|---|---|---|---|---|
Recording standard: record the sign and unit every time. A bare number such as "1.08" is incomplete. A result such as "+1.08 V, red lead on Cu, black lead on Zn, 21.0 °C" is interpretable.
For +1.08 V, +1.09 V and +1.08 V:
Mean Ecell = (1.08 + 1.09 + 1.08) / 3 = 1.083... V ≈ +1.08 V to the resolution of a 0.01 V meter.
Cu2+/Cu = +0.34 V
Zn2+/Zn = -0.76 V
E°cell = +0.34 - (-0.76) = +1.10 V
For Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s):
E°cell = E°right - E°left
+1.10 = +0.34 - E°(Zn2+/Zn), giving E°(Zn2+/Zn) = -0.76 V.
Measured mean = +1.08 V; theoretical standard value = +1.10 V.
Absolute difference = |1.08 - 1.10| = 0.02 V
Percentage difference = (0.02 / 1.10) x 100 = 1.8%
Meter resolution is the smallest displayed increment; it is not automatically the total measurement uncertainty. Use the manufacturer's accuracy specification where available, report the spread of independent repeats and discuss systematic effects separately.
Precision warning: only use the degree symbol when values genuinely refer to standard conditions. A potential derived from non-standard classroom measurements is an experimental or apparent potential, not automatically a standard electrode potential.
Exam rule: do not multiply an electrode potential by a stoichiometric coefficient. E° is an intensive potential. Balance electrons in the half-equations, but use the tabulated E° values unchanged.
| Observation | Interpretation |
|---|---|
| Positive reading with red lead on Cu and black lead on Zn | Copper is at the higher potential and acts as the positive cathode; zinc acts as the negative anode. |
| Negative reading with red lead on Zn and black lead on Cu | The same cell can be operating, but the meter leads are reversed relative to the expected polarity. |
| Stable value close to +1.10 V for Zn/Cu | Broadly consistent with the predicted redox direction and standard-potential difference, subject to actual conditions. |
| Value lower than expected | Possible non-standard concentrations, oxide films, a poorly wetted or nearly broken ionic connection, liquid-junction potential, contamination, meter loading, temperature difference or side reactions. |
| Value changes continuously | Surface reactions, changing contact, a drying bridge, solution mixing, polarisation or unstable redox composition may be occurring. |
| Near-zero reading | Possible open circuit, dry bridge, poor lead contact, identical half-cells, depleted meter battery or incorrect meter setting. |
| Unexpected precipitate | The bridge electrolyte or transferred contamination has reacted with a half-cell species; the result is invalid. |
A measured value close to a tabulated result supports the proposed cell reaction, but it does not prove that all species are at standard activity, that the bridge contributes no junction potential, or that the electrode surfaces are ideal. The practical measures the potential difference of the complete cell, not the absolute potential of a single electrode.
| Observation | What students should see | Interpretation |
|---|---|---|
| Zinc-copper voltage | A positive reading commonly near 1.0-1.1 V under suitable classroom conditions with red on Cu and black on Zn. | Correct polarity and a working ionic/electronic circuit. |
| Copper solution | Blue solution remains visibly blue during a short high-resistance measurement. | Very little Cu2+ is consumed because current is tiny. |
| Zinc solution | Usually colourless with no obvious visual change. | Zn2+ is colourless and only a tiny amount forms during open-circuit measurement. |
| Electrode mass | No obvious mass change in a brief measurement. | The high-resistance meter minimises current and chemical conversion. |
| Salt bridge | Wet throughout with both ends immersed. | The ionic circuit is complete. |
| Gas formation | None expected in the intended Zn/Cu cell. | Gas suggests contamination, an unintended reaction or severe non-standard conditions. |
The following data are illustrative and should not be presented as guaranteed experimental outcomes.
| Trial | Temperature / °C | Stable Ecell / V |
|---|---|---|
| 1 | 20.5 | +1.08 |
| 2 | 20.6 | +1.09 |
| 3 | 20.5 | +1.08 |
Evaluation: the result is close to the theoretical standard value, but the core solutions are 0.40 mol dm-3 rather than 1.00 mol dm-3. Because the zinc-ion and copper-ion concentrations are nominally equal, the ideal concentration ratio is close to 1, but activities, temperature, surface condition and junction effects remain. Describe the result as consistent with theory, not as a direct measurement of two perfect standard half-cells.
| Problem | Likely cause | Corrective action |
|---|---|---|
| No reading | Open lead, dead meter, wrong mode, dry bridge or no contact | Check the meter on a known source; inspect leads; re-wet or replace the bridge. |
| Negative reading | Red and black leads reversed relative to expected polarity | Record the connections and sign; use the result to identify the higher-potential electrode. |
| Reading is much too low | Dirty electrode, nearly dry or poorly contacting bridge, incorrect concentration, unintended short circuit, unsuitable meter input resistance or side reaction | Re-clean surfaces, replace the bridge, verify concentrations and rebuild the cell. |
| Reading drifts | Bridge drying, solutions mixing, unstable surface, temperature change or polarisation | Use a fresh bridge, stabilise temperature and define a reading time or stability criterion. |
| Different groups get very different values | Inconsistent cleaning, concentrations, meters, temperatures or bridge preparation | Standardise preparation, cross-check meters and use a common reference cell. |
| Precipitate at bridge end | Bridge ion has reacted, for example Cl- with Ag+ | Discard the result, replace solutions and use a compatible bridge electrolyte. |
| Iron half-cell gives poor repeatability | Oxide/rust, Fe2+ oxidation by air or contamination | Use freshly prepared solution and freshly polished iron; minimise exposure and delays. |
| Magnesium result is unreliable | Passivation and reaction with water interfere with the intended equilibrium | Use as an evaluation example rather than a precision reference unless the method has been validated. |
"The salt bridge carries electrons."
False. It carries ions; electrons move through the external circuit.
"The more reactive metal must be the positive electrode."
In the Zn/Cu voltaic cell, the metal more readily oxidised is the negative anode.
"Doubling electrode area doubles Ecell."
Open-circuit potential is not proportional to electrode area, although area can affect stabilisation and current under load.
"Cathodes are always positive."
Cathodes are always sites of reduction, but their sign depends on whether a cell is voltaic or electrolytic.
"A positive E°cell means the reaction must be fast."
It indicates thermodynamic feasibility under stated conditions, not reaction rate.
"E° values are measured for single electrodes on their own."
They are defined and measured relative to a reference electrode.
"Room temperature automatically means standard conditions."
Standard treatment also requires specified pressure and ionic activity/concentration conditions.
| Mistake | Consequence |
|---|---|
| Using a dry or barely wet bridge | Zero, low or unstable voltage. |
| Putting the wrong metal in a solution | The intended half-cell is not constructed; contamination may occur. |
| Touching cleaned electrode surfaces | Grease and salts alter the interface and repeatability. |
| Allowing crocodile clips to touch solution | Contamination, corrosion and competing electrode materials. |
| Recording magnitude but not sign or lead connections | Polarity cannot be interpreted. |
| Using an ammeter or shorting the electrodes | Large current, rapid polarisation and chemical change. |
| Reusing a contaminated bridge | Half-cell compositions change and results drift. |
| Adding E° values without considering direction | Incorrect E°cell. |
| Multiplying E° by the number of electrons | Incorrect because potential is not an extensive quantity. |
| Calling the cell reading "the zinc electrode potential" | A complete cell potential has been measured, not an isolated electrode potential. |
| Source of variation or bias | Effect on result | Improvement |
|---|---|---|
| Meter resolution and calibration | Limits discrimination and may introduce systematic offset. | Use a calibrated high-input-resistance meter with 0.001 V resolution for small effects. |
| Electrode oxide, grease or corrosion | Slow or unstable interfacial response; poor repeatability. | Polish consistently immediately before each trial; rinse, blot and avoid touching. |
| Concentration error | Changes electrode potentials and the overall reaction quotient. | Prepare solutions with volumetric glassware; label clearly; use dedicated pipettes. |
| Temperature variation | Changes equilibrium potentials and may alter stabilisation behaviour. | Allow thermal equilibration; use a thermostated water bath for high-precision work. |
| Salt bridge drying or inconsistent saturation | Increases internal resistance and may cause slow stabilisation, drift or an abnormally low reading. | Use fresh bridges of identical dimensions and a fixed soaking time. |
| Bridge electrolyte reacts with half-cell | Changes composition or forms a precipitate. | Choose an inert electrolyte; avoid chloride with Ag+. |
| Liquid-junction potential | Adds an unmeasured potential at the interface. | Use a concentrated inert bridge electrolyte with similar ion mobilities and acknowledge residual bias. |
| Cross-contamination | Changes concentrations and may introduce competing redox couples. | Use a fresh bridge and clean apparatus for every pair; separate droppers and electrodes. |
| Reading taken too early | Captures transient response rather than stable open-circuit potential. | Use a defined stability criterion and record time-series readings. |
| Non-independent repeats | Underestimates variation. | Reconstruct the apparatus rather than merely rereading the same cell. |
| Passivation or side reactions | Measured potential may not represent the intended couple. | Select suitable metals, use fresh surfaces and discuss kinetics separately from thermodynamics. |
The largest discrepancy is not always random meter uncertainty. Departure from standard conditions, non-ideal activities, surface films and liquid-junction potentials can produce systematic differences that repeats will not remove.
Research question: How does copper(II)-ion concentration affect the cell potential of a zinc-copper electrochemical cell?
Independent variable: CuSO4 concentration, for example 0.100, 0.0100 and 0.00100 mol dm-3, prepared by accurate serial dilution.
Dependent variable: stable open-circuit Ecell, preferably measured to 0.001 V.
Expected trend: lowering [Cu2+] increases the reaction quotient [Zn2+]/[Cu2+] and should reduce Ecell. The relationship is logarithmic rather than linear in concentration.
Alternative extension: construct an electrochemical series using several pairwise cell potentials such as Zn2+/Zn, Fe2+/Fe, Cu2+/Cu and, where approved, Ag+/Ag. Use one half-cell as a reference, retain the signs and derive a consistent order of reduction potentials.
Strict electrochemical thermodynamics uses activities. At higher ionic strength, activity coefficients mean that a 1.00 mol dm-3 solution does not behave ideally.
The boundary between bridge electrolyte and half-cell solution contributes a small potential because ions diffuse at different rates.
A single electrode potential is not directly measurable. The standard hydrogen electrode defines zero; practical laboratories often use secondary reference electrodes.
EMF is measured with negligible current. Under load, terminal potential is reduced by internal resistance, polarisation and concentration changes.
A positive E°cell does not overcome a large activation barrier or passivating film. Aluminium and magnesium can illustrate this limitation.
E° is not multiplied by stoichiometric coefficients. By contrast, ΔG° and charge transferred scale with the amount of reaction.
ΔG° = -nFE°cell
Combining ΔG° = -RT ln K with ΔG° = -nFE°cell gives ln K = nFE°cell/RT.
E changes with temperature according to reaction entropy; there is no universal "temperature increases voltage" rule. Inert Pt or graphite provides an electron-transfer surface for solution redox couples, and closely grouped repeats can still share systematic bias.
E°(Fe3+/Fe2+) = +0.77 V and E°(Cu2+/Cu) = +0.34 V, so Fe3+/Fe2+ is reduced and copper is oxidised.
Cathode: 2Fe3+ + 2e- → 2Fe2+
Anode: Cu → Cu2+ + 2e-
Overall: 2Fe3+ + Cu → 2Fe2+ + Cu2+
E°cell = +0.77 - (+0.34) = +0.43 V
Cell notation: Cu(s) | Cu2+(aq) || Fe3+(aq), Fe2+(aq) | Pt(s)
| Stage | Activity | Time / min |
|---|---|---|
| 1 | Context, redox recap and learning objectives | 5 |
| 2 | Safety briefing and apparatus demonstration | 7 |
| 3 | Prepare solutions, electrodes and salt bridge | 10 |
| 4 | Construct core Zn/Cu cell and record stability readings | 10 |
| 5 | Independent repeats or additional half-cell pairs | 10-15 |
| 6 | Process data and calculate E°cell | 10 |
| 7 | Interpret polarity, equations and cell notation | 8 |
| 8 | Evaluation, exam questions and plenary | 5-10 |
Observe whether a student:
Overall risk: low to moderate after suitable controls. Schools should complete their own current risk assessment using concentration-specific safety information and local laboratory procedures.
| Hazard | Main risk | Control measures |
|---|---|---|
| Zinc sulfate solution | Harmful if swallowed; eye/skin exposure; environmental harm depending on concentration | Wear eye protection; use small volumes; clean spills promptly; avoid ingestion/contact; collect as metal-ion waste. |
| Copper(II) sulfate solution | Eye/skin exposure and environmental harm depending on concentration | Wear eye protection; minimise volume; prevent release; collect in labelled waste. |
| Saturated potassium nitrate | Oxidising; can intensify combustion, particularly if allowed to dry on paper | Keep away from flames, heat and reducing/combustible materials; use small quantities; keep soaked paper wet until approved disposal. |
| Iron(II) sulfate solution, optional | Harmful at higher concentration; stains and spills | Eye protection; small quantities; avoid contact; collect as metal-ion waste. |
| Silver nitrate, optional | Oxidising, eye/skin damage at higher concentration, staining and environmental harm | Use only an approved dilute concentration; teacher control where appropriate; avoid chloride; keep silver waste separate. |
| Metal strips and abrasive paper | Sharp edges, splinters, metal dust and contaminated hands | Handle by edges; avoid vigorous abrasion; wash hands; do not touch the face. |
| Glassware | Breakage and cuts | Keep beakers away from bench edges; clear breakages with proper equipment. |
| Electrical connections | Very low shock risk, but short circuits can heat wires and alter chemistry | Use only low-voltage measuring equipment; avoid direct short circuits; inspect insulation. |
Disposal: collect metal-ion solutions and contaminated bridges in labelled waste according to local policy. Do not pour them into sinks unless the institution's current procedure explicitly permits it. Keep silver-containing waste separate.
It completes the internal ionic circuit and allows ions to migrate so charge does not rapidly build up in either half-cell.
From the zinc anode to the copper cathode through the external circuit.
Zn(s) → Zn2+(aq) + 2e-; Cu2+(aq) + 2e- → Cu(s); overall Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s).
E°cell = +0.34 - (-0.76) = +1.10 V.
It draws very little current, so the reading is close to the near-open-circuit cell potential and half-cell compositions change only slightly.
Zinc is at lower potential than copper, so V(red) - V(black) is negative. The lead polarity is reversed relative to the conventional positive reading.
To remove oxide, corrosion and grease and produce a more reproducible metal-solution interface.
The cell may not be at 298 K or standard activity; surfaces may be imperfect; junction potentials, contamination, passivation and meter limitations can contribute.
Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s).
Ecell decreases because the reaction quotient [Zn2+]/[Cu2+] increases, making the forward reaction less thermodynamically favourable.
Potential is an intensive property. Scaling the amount of reaction changes charge and Gibbs energy proportionally, but not the potential difference.
The reaction as written is thermodynamically feasible under standard conditions. It may still be kinetically slow or inhibited.
| Plenary question | Expected answer |
|---|---|
| What is oxidised in a zinc-copper electrochemical cell? | Zinc is oxidised. |
| In which direction do electrons flow? | From the zinc anode to the copper cathode through the external circuit. |
| What is the purpose of the salt bridge? | It completes the circuit and allows ions to move between half-cells, maintaining electrical neutrality. |
| How is the standard cell potential calculated? | E°cell = E°cathode - E°anode. |
| Why must E° values not be multiplied when balancing half-equations? | Electrode potential is an intensive property, so it does not change when a half-equation is multiplied. |
| Why is little visible chemical change usually observed? | The high-resistance voltmeter allows only a very small current to flow, so very little zinc dissolves or copper is deposited. |
A high-resistance voltmeter draws only a tiny current, so very little zinc dissolves or copper is deposited during the measurement.
Both ends must contact the half-cell solutions to complete the internal ionic circuit.
No. It normally means the red lead is connected to the lower-potential electrode. Record the sign and lead connections.
The classroom cell may not be at 298 K, unit activity or other standard conditions, and it also includes surface and liquid-junction effects.
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This guide was written and reviewed by Better Equipped's technical team and further reviewed by former A Level Science Teachers. Our technical team draw on experience supplying practical science equipment to schools, colleges, laboratories and science departments throughout the UK. They include ex-school laboratory technicians and are here to support schools, colleges and laboratories. If you have feedback on this guide, we'd love to here it so please contact us. Don't forget we will be regularly updating our guides and resources on our Better-Resources Hub. If you would like us to cover a particular subject matter in these guides or have some top tips you'd like to share then again we'd love to hear from you.
Last reviewed and updated: August 2026
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