In this investigation, you will measure the electromotive force (EMF) of a simple electrochemical cell as you systematically vary the concentration of one electrolyte. By applying the Nernst equation, you can derive a linear relationship between EMF and the logarithm of concentration, allowing you to calculate fundamental thermodynamic quantities. This is one of the most analytically powerful IB Chemistry HL Internal Assessment topics available.
This practical is suitable for IB Diploma Chemistry HL.
Background Theory
A simple electrochemical cell consists of two half-cells connected by a salt bridge. For a Cu/Zn cell:
Zn(s) | Zn²⁺(aq) || Cu²⁺(aq) | Cu(s)
The standard cell potential E° = +1.10 V. The actual cell potential at non-standard concentrations is given by the Nernst equation:
E = E° − (RT/nF) ln Q
Where R = 8.314 J mol⁻¹ K⁻¹, T = temperature (K), n = moles of electrons transferred, F = 96485 C mol⁻¹, and Q = reaction quotient. At 25 °C, this simplifies to:
E = E° − (0.0257/n) ln Q
If the Zn²⁺ concentration is kept constant and only [Cu²⁺] is varied, a plot of E vs ln[Cu²⁺] should be linear with gradient = 0.0257/n = 0.01285 V (for n = 2). This gradient allows experimental determination of F/R — a remarkable result from a school laboratory.
Variables
- Independent variable (IV): Concentration of Cu²⁺(aq) (mol dm⁻³) — e.g. 0.001, 0.005, 0.010, 0.050, 0.100, 0.500, 1.000 mol dm⁻³
- Dependent variable (DV): Cell EMF (V) measured by a high-resistance voltmeter or multimeter
- Controlled variables (CV): Concentration of ZnSO₄ solution (kept at 1.0 mol dm⁻³), temperature (25 °C), same electrodes cleaned before each measurement, same salt bridge composition
Equipment
- Copper strip electrode and zinc strip electrode
- CuSO₄ solutions at 0.001, 0.005, 0.010, 0.050, 0.100, 0.500, 1.000 mol dm⁻³ (prepared by serial dilution)
- ZnSO₄ solution (1.0 mol dm⁻³)
- Salt bridge (filter paper soaked in saturated KNO₃ solution)
- High-resistance digital voltmeter or multimeter
- Sandpaper to clean electrodes
- 100 cm³ beakers (one per concentration)
- Water bath at 25 °C (optional but recommended)
Safety
Safety
⚠️ Copper sulfate is harmful if ingested and an irritant to skin and eyes — wear gloves and eye protection. Zinc sulfate is also an irritant. Dispose of all solutions into the appropriate waste disposal bottles provided.
Method
- Prepare CuSO₄ solutions by serial dilution from a 1.000 mol dm⁻³ stock. Verify concentrations using the dilution equation C₁V₁ = C₂V₂.
- Pour 50 cm³ of 1.0 mol dm⁻³ ZnSO₄ into a beaker. Place the cleaned zinc electrode into this solution.
- Pour 50 cm³ of your first CuSO₄ concentration into a second beaker. Place the cleaned copper electrode into this solution.
- Connect the two beakers with a salt bridge. Connect the voltmeter with the positive terminal to the copper electrode.
- Record the EMF once it has stabilised (typically 30–60 seconds). Record to 3 decimal places if possible.
- Remove the salt bridge, rinse and dry the electrodes with sandpaper, and repeat for each CuSO₄ concentration.
- Take at least three readings per concentration and calculate the mean EMF.
Results Table
| [Cu²⁺] (mol dm⁻³) | ln[Cu²⁺] | EMF 1 (V) | EMF 2 (V) | EMF 3 (V) | Mean EMF (V) |
|---|---|---|---|---|---|
| 0.001 | −6.91 | ||||
| 0.005 | −5.30 | ||||
| 0.010 | −4.61 | ||||
| 0.050 | −2.99 | ||||
| 0.100 | −2.30 | ||||
| 0.500 | −0.69 | ||||
| 1.000 | 0.00 |
Analysis
1. Plot mean EMF (y-axis, V) against ln[Cu²⁺] (x-axis). The Nernst equation predicts a straight line.
2. Calculate the gradient of the best-fit line. Theory predicts: gradient = RT/nF = 0.01285 V at 25 °C.
3. Rearrange to find the Faraday constant: F = RT / (n × gradient)
4. Compare your calculated F to the accepted value of 96485 C mol⁻¹. Calculate the percentage error and use worst-case lines to determine the uncertainty in your gradient and propagate this to δF.
5. Read off the y-intercept of your graph. Theory predicts this equals E° − (0.01285) ln[Zn²⁺]. Use your Zn²⁺ concentration to calculate the theoretical intercept and compare.
Discussion Points
- Why does the EMF decrease as [Cu²⁺] decreases? Relate your answer to Le Chatelier’s principle and the Nernst equation.
- Why must the voltmeter have a very high internal resistance?
- Why is it important to clean the electrodes with sandpaper between readings?
- What is the purpose of the salt bridge and why does KNO₃ work well for this?
- Why does the graph not pass exactly through E° = 1.10 V at [Cu²⁺] = 1.0 mol dm⁻³?
Guidance
IA Guidance
This IA is exceptional because it allows experimental determination of the Faraday constant — a fundamental physical constant — from a simple school electrochemistry setup. To score highly:
- Research Design: Justify your concentration range using the Nernst equation — show that at very low [Cu²⁺] the EMF change per decade is predictable. Explain why a log scale is used for concentration.
- Data Analysis: Include error bars based on the spread of your three readings. Use worst-case gradient lines to propagate uncertainty to δF. Compare your F value to the accepted value and comment on whether it falls within your uncertainty range.
- Conclusion: State your derived value of F with uncertainty. Discuss whether any systematic errors (electrode surface oxidation, temperature variation, salt bridge junction potential) could explain a consistent offset.
- Evaluation: Suggest using a reference electrode (standard hydrogen electrode or calomel electrode) to eliminate junction potential errors, and a temperature-controlled bath to ensure the Nernst equation applies exactly.
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