8.1.7 Making Simple Cells

Objective:

To measure the electromotive force (EMF) produced by different combinations of metal electrodes and investigate the effect of changing conditions (such as concentration) on the EMF of an electrochemical cell.

This experiment explores how standard electrode potentials influence the voltage generated between two half-cells, as well as how deviations from standard conditions affect EMF.

Apparatus & Chemicals:

• Metal electrodes (e.g., $Zn$, $Cu$, $Ni$, $Mg$, $Fe$)
• 1.0 mol dm$⁻³$ solutions of metal ions (e.g., $Zn²⁺$, $Cu²⁺$, $Fe²⁺$, $Ni²⁺$)
• Potassium nitrate ($KNO₃$) or similar solution for the salt bridge
• High-resistance voltmeter
• Filter paper or salt bridge tube
• Beakers
• Sandpaper (for cleaning electrodes)
• Measuring cylinder
• Distilled water

Key Procedure Steps:

1. Preparation of Metal Electrodes:

• Clean the metal electrodes thoroughly using sandpaper to remove any oxide layer that may have formed, ensuring good electrical contact.
• Rinse the electrodes under cold running water to remove any remaining debris.

2. Assembly of Electrochemical Cells:

• Set up a simple electrochemical cell by placing one metal electrode into a beaker containing a 1.0 mol dm$⁻³$ solution of its ions (e.g., a $Zn$ electrode into $Zn²⁺$solution).
• In a second beaker, place the other metal electrode in its corresponding ion solution (e.g., $Cu$ electrode in $Cu²⁺$solution).

3. Salt Bridge Setup:

• Prepare a salt bridge by soaking a strip of filter paper in saturated potassium nitrate solution (or use a pre-prepared salt bridge).
• Ensure both ends of the salt bridge are immersed in the two beakers, connecting the solutions of the two half-cells.

4. Measuring the EMF:

• Connect the two metal electrodes to a high-resistance voltmeter, ensuring the correct polarity (the positive terminal of the voltmeter should be connected to the electrode with the higher reduction potential).
• Record the measured EMF from the voltmeter.

5. Repeat for Different Electrode Combinations:

• Repeat the experiment using different combinations of metal electrodes (e.g., $Zn$ and $Fe$, $Mg$ and $Ni$).
• Compare the measured EMF values with the calculated EMF values from standard electrode potentials in data tables.

Investigation of Concentration Effect:

Once the EMF of the different electrode combinations has been measured, the same setup can be used to explore how concentration affects the EMF of a cell.

6. Dilution of Solutions:

• After measuring the EMF for a standard 1.0 mol dm⁻³ solution, dilute one of the solutions by a factor of 10.
• For example, take 25 cm³ of the original solution and dilute it to 250 cm³ with distilled water.

7. Measurement of EMF:

• Measure the EMF again using the diluted solution and record the new value.
• Repeat the dilution process for a total of 5 dilutions, each time diluting by a factor of 10 and measuring the corresponding EMF.

Analysis:

• The measured EMF values for the electrode combinations are usually lower than the calculated values based on standard electrode potentials.
• This is because non-standard conditions (such as temperature variations, concentration changes, or electrode surface irregularities) affect the EMF of the cell.
• In the concentration investigation, plot a graph of EMF against the dilution factor.
• The graph should show that the EMF decreases as the concentration of the ions in the solution decreases, per the Nernst equation:

$$E = E^\circ - \frac{0.059}{n} \log \left( \frac{[\text{reduced species}]}{[\text{oxidised species}]} \right)$$

• where $n$ is the number of electrons involved in the half-reaction.
• Lower concentrations result in fewer ions available to react, thereby reducing the EMF.

Specimen Results:

Measured EMF Specimen Results Table:

Electrode Combination	Measured EMF / V	Calculated EMF / V
$Zn$/$Zn²⁺$and $Cu²⁺$/$Cu$	1.04	1.10
$Mg$/$Mg²⁺$and $Ni²⁺$/$Ni$	1.50	2.12
$Zn$/$Zn²⁺$and $Fe²⁺$/$Fe$	0.25	0.32
$Ni²⁺$/$Ni$ and $Cu²⁺$/$Cu$	0.42	0.59

Measured values are generally lower than the calculated EMF from standard electrode potentials due to non-standard conditions.

Concentration and EMF Specimen Results Table:

Concentration (mol dm$⁻³$)	EMF / V
1.0	1.4
0.1	1.2
0.01	0.8
0.001	0.6
0.0001	0.4

Graph of EMF Against Dilution Factor:

• The graph shows that as the concentration decreases, the EMF also decreases.
• The relationship is approximately proportional to the logarithm of the dilution factor, consistent with the Nernst equation.

Practical Tips:

• Ensure the metal electrodes are thoroughly cleaned before immersion in the electrolyte solutions to avoid contamination that can affect EMF values.
• Use metals with a large difference in electrode potentials for the most meaningful results. Combinations with close electrode potentials may result in very small or negligible EMF.
• The temperature of the solutions should be controlled as variations can affect EMF measurements.

Further Investigations:

• Temperature Effects: Investigate how changing the temperature affects the EMF of the cell. Higher temperatures typically increase reaction rates but can also affect the position of equilibrium in redox reactions.
• Electrode Spacing: Changing the spacing between the electrodes can influence the internal resistance of the cell, impacting the EMF.
• Salt Bridge: Test different types of salt bridges (e.g., different salts or lengths of bridges) to observe their effects on the stability and magnitude of the EMF.

Conclusion:

This experiment demonstrates the relationship between electrode potentials, ion concentration, and the resulting EMF of electrochemical cells. By varying the conditions such as electrode combinations and concentrations, students can gain a deeper understanding of redox chemistry and the factors influencing cell potential.