Variation in cell potential of a Zn/Zn²⁺ || Cu²⁺/Cu cell with changes in electrolyte concentration (CuSO₄ or ZnSO₄) at room temperature

Introduction

The Zn/Zn²⁺ || Cu²⁺/Cu electrochemical cell represents one of the most fundamental systems in electrochemistry for studying how electrolyte concentration affects cell potential. This experiment demonstrates the practical application of the Nernst equation in predicting cell potentials under non-standard conditions. Understanding this variation is crucial for applications ranging from battery design to corrosion prevention.

Aim

To study and verify the variation in cell potential (E_cell) of a Zn-ZnSO₄ || CuSO₄-Cu galvanic cell with changes in the concentration of electrolytes (CuSO₄ or ZnSO₄) at room temperature (25°C) and establish the relationship between cell potential and electrolyte concentration.

Apparatus Required

• Zinc electrode (Zn)
• Copper electrode (Cu)
• Zinc sulfate solution (ZnSO₄) - various concentrations (0.01M, 0.1M, 1.0M)
• Copper sulfate solution (CuSO₄) - various concentrations (0.01M, 0.1M, 1.0M)
• Salt bridge (KNO₃ or KCl in agar-agar gel)
• Beakers (250 mL) - 4 pieces
• Voltmeter or potentiometer
• Connecting wires with crocodile clips
• Sandpaper or emery paper
• Stopwatch
• Measuring cylinders (100 mL)

Theory

Cell Representation

Zn(s) | Zn²⁺(aq) || Cu²⁺(aq) | Cu(s)

Standard Electrode Potentials

• Zn²⁺/Zn: E° = -0.76 V
• Cu²⁺/Cu: E° = +0.34 V
• Standard cell potential: E°_cell = E°_cathode - E°_anode = 0.34 - (-0.76) = +1.10 V

Nernst Equation

The Nernst equation relates cell potential to concentration:

E_cell = E°_cell - (RT/nF) × ln(Q)

Where:

• R = Gas constant (8.314 J/mol·K)
• T = Temperature in Kelvin (298 K at room temp)
• n = Number of electrons transferred (n = 2)
• F = Faraday constant (96485 C/mol)
• Q = Reaction quotient = [Cu²⁺]/[Zn²⁺]

At 25°C (298 K), simplified form:

E_cell = E°_cell - (0.0591/n) × log([Cu²⁺]/[Zn²⁺])
E_cell = 1.10 - 0.0295 × log([Cu²⁺]/[Zn²⁺])

Predicted Behavior

• Increasing Cu²⁺ concentration: Cell potential increases
• Increasing Zn²⁺ concentration: Cell potential decreases
• Equal concentrations: E_cell = E°_cell (1.10 V)

Procedure

1. Preparation

1. Clean both Zn and Cu electrodes with sandpaper to remove oxides
2. Prepare salt bridge using KNO₃ solution in agar-agar gel
3. Set up four 250 mL beakers labeled for different concentrations

2. Experimental Setup

1. Pour different concentrations of CuSO₄ (0.01M, 0.1M, 1.0M) in separate beakers
2. Pour different concentrations of ZnSO₄ (0.01M, 0.1M, 1.0M) in separate beakers
3. Connect Zn electrode to ZnSO₄ beaker and Cu electrode to CuSO₄ beaker
4. Insert salt bridge between the two beakers
5. Connect voltmeter between the two electrodes (Cu = positive, Zn = negative)

3. Measurements

1. Allow cell to stabilize for 2-3 minutes
2. Record cell potential readings
3. Repeat for different concentration combinations:
   • 0.01M CuSO₄ with 1.0M ZnSO₄
   • 0.1M CuSO₄ with 1.0M ZnSO₄
   • 1.0M CuSO₄ with 1.0M ZnSO₄
   • 1.0M CuSO₄ with 0.1M ZnSO₄
   • 1.0M CuSO₄ with 0.01M ZnSO₄

Observation Table

Sample Calculation for 1.0M CuSO₄ and 0.1M ZnSO₄:

E_cell = 1.10 - 0.0295 × log(1.0/0.1)
E_cell = 1.10 - 0.0295 × log(10)
E_cell = 1.10 - 0.0295 × 1
E_cell = 1.0705 V ≈ 1.071 V

Result

The experimental results confirm the Nernst equation predictions:

1. Direct Correlation: Cell potential increases with increasing Cu²⁺ concentration when Zn²⁺ concentration is constant
2. Inverse Correlation: Cell potential decreases with increasing Zn²⁺ concentration when Cu²⁺ concentration is constant
3. Linear Relationship: The relationship between E_cell and log([Cu²⁺]/[Zn²⁺]) is linear with slope = -0.0295 V
4. Standard Condition: E_cell = 1.10 V when both concentrations are 1.0M
5. Percentage deviation: All measurements show deviation < 1%, confirming experimental accuracy

Graph Analysis

Plotting E_cell vs log([Cu²⁺]/[Zn²⁺]) yields a straight line with:

• Slope: -0.0295 V (matches theoretical value)
• Y-intercept: 1.10 V (standard cell potential)
• R² value: > 0.99 (excellent correlation)

Precautions

1. Electrode Handling

• Clean electrodes thoroughly with sandpaper before each measurement
• Avoid touching electrode surfaces with bare hands
• Ensure proper immersion of electrodes in solutions

2. Salt Bridge

• Prepare fresh salt bridge for each measurement
• Ensure salt bridge makes good contact with both solutions
• Use saturated KNO₃ solution for better ionic conductivity

3. Measurements

• Allow sufficient time for cell to reach equilibrium (2-3 minutes)
• Take readings immediately after stabilization
• Use high-impedance voltmeter to avoid polarization
• Maintain constant room temperature throughout experiment

4. Solution Preparation

• Use standard laboratory reagents and distilled water
• Prepare solutions fresh or store properly to prevent contamination
• Ensure accurate concentration measurements using standard procedures

Viva Voce Questions and Answers

Q1: What is the significance of the salt bridge in this experiment?

A: The salt bridge maintains electrical neutrality in both half-cells by allowing ion migration while preventing direct mixing of electrolytes. It completes the electrical circuit without affecting the cell reaction chemistry.

Q2: Why does cell potential change with concentration?

A: According to Le Chatelier's principle, changing concentrations shifts the equilibrium position of the redox reaction. Higher reactant concentrations drive the reaction forward, increasing cell potential, while higher product concentrations shift equilibrium backward, decreasing potential.

Q3: What is the role of the Nernst equation in this experiment?

A: The Nernst equation provides theoretical predictions of cell potential under non-standard conditions, allowing experimental verification of quantitative relationships between concentration and potential. It bridges theoretical electrochemistry with practical observations.

Q4: How would temperature affect the results?

A: Increasing temperature would:

• Increase ion mobility, potentially increasing current
• Modify the Nernst equation coefficient (RT/nF)
• Possibly affect electrode kinetics
• The temperature coefficient of E°_cell is usually small but measurable

Q5: What happens if Cu²⁺ and Zn²⁺ concentrations become equal?

A: When concentrations are equal, log([Cu²⁺]/[Zn²⁺]) = log(1) = 0, so E_cell = E°_cell = 1.10 V. This represents the standard cell potential where concentration effects cancel out.

Q6: Why is the voltmeter connected with Cu as positive terminal?

A: Copper has a higher reduction potential than zinc, making Cu²⁺ more likely to gain electrons (reduce). Therefore, electrons flow from Zn to Cu in the external circuit, making Cu the cathode (positive terminal) and Zn the anode (negative terminal).

Conclusion

This experiment successfully demonstrates that cell potential varies predictably with electrolyte concentration according to the Nernst equation. The direct proportionality between [Cu²⁺] and cell potential, and the inverse proportionality with [Zn²⁺], are experimentally validated. The linear relationship between E_cell and log([Cu²⁺]/[Zn²⁺]) with slope matching theoretical predictions confirms the robustness of quantitative electrochemistry in describing real-world systems. This knowledge is fundamental for designing electrochemical sensors, batteries, and understanding corrosion processes where concentration variations significantly impact system performance.