ELECTROCHEMICAL CELLS 3-21-05 Purpose: In this experiment, several different half-cells will be prepared and connected to find the voltages generated. Also, the concentration will be change in one of the solutions to see how this affects the cell potential. Thirdly, the electrical potential of a cell containing silver and silver chloride will be measured. Lastly, a cell containing copper (II) and ammonia will be constructed. The potential and the Nernst equation will be used to calculate the formation constant of the (II) complex ion. This experiment uses a micro scale technique.
Procedure: 1. Collect materials and put on goggles for eye protection. 2. Prepare a test cell to measure the voltage of the copper and zinc half-cells. 3. Put approximately 2 mL 1.
0 M zinc nitrate solution in one of the center wells of a 24-well plate. 4. Put approximately 2 mL of 1. 0 M copper (II) nitrate in an adjacent well. Polish small strips of zinc and copper metal, and place the metal in the appropriate well containing the solution of the ions of that metal.
5. Take a small strip of filter paper that has been soaked in potassium nitration solution, and drape it across the wells so that one end dips in the solution in each well. This will act as a salt bridge. 6. Use a voltmeter to measure the potential difference between the two half-cells. 7.
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Connect the meter so the voltage reading is positive. 8. Use the most sensitive scale which is practical. Make note as to which electrode is the anode and which is the cathode. 9. When the voltmeter reads a positive voltage, the electrode connected to the positive terminal is the cathode and is undergoing reduction.
The oxidation is occurring at the electrode connected to the negative terminal, the anode. 10. Prepare half-cells in other wells of the 24-well plates. 11. First make a diagram of the order of the solutions in the wells so the different solutions will not be confused. 12.
Pour some 1. 0 M solution of each of the ions into different wells. 13. Polish the metals with sandpaper or steel wool so that they are shiny, and insert them into the wells which contains the ion of the same metal.
14. Use fresh strips of filter paper soaked in 1. 0 M potassium nitrate as salt bridges. 15.
The zinc electrode will be designated as the standard electrode. Measure the potential difference between the zinc electrode and each of the other electrodes. 16. Record the data. 17. Measure the potential difference between at least six combinations of the various electrodes.
18. Use table of electrode potentials to predict the voltage and which half-cell will be the anode and which the cathode. 19. Compare the predicted and the measured potentials. 20. Record the data.
21. Dilute the 1. 0 M copper (II) nitrate so it is 0. 0010 M. 22.
Count 18 drops of distilled water into a small test tube and add 2 drops of the 1. 0 M copper (II) solution. 23. Mix well by pouring back and forth from one test tube to another. This solution is now 0. 10 M 24.
Repeat this dilution process two more times, preparing a solution which is 0. 0010 M. 25. Pour some of 0. 0010 M copper (II) nitrate solution into one of the wells in the well plate. 26.
Add a piece of polished copper wire and measure the voltage against the standard zinc electrode. 27. Record the data. 28. Pour 10 mL of 1. 0 M Na Cl solution into a beaker.
29. Add one drop of 1. 0 M silver nitrate to the Na Cl solution, and stir well. 30. Pour some of the solution into one of the wells in the well plate and add a silver metal electrode. Measure the potential difference versus this half-cell and the zinc half-cell.
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31. Find the volume of one drops of 1. 0 M copper (II) nitrate solution by counting the number of drops in 1 mL. 32.
Put 10 mL of 6. 0 M ammonia in a beaker and add one drop of 1. 0 M copper (II) nitrate solution. Stir well. 33. Pour some of the solution into one of the wells in the well plate and add a copper metal electrode.
34. Measure the potential difference versus the zinc electrode. Record data. 35. Empty the solutions and solids into the sink.
36. Wash the materials and dry them well. Data: Voltage of each half-cell versus the zinc electrode Voltage Anode Cathode Zn versus Ag Zn versus Cu Zn versus Fe Zn versus Mg Zn versus Pb Reduction Equations for Each Ion Arranged in Decreasing Order of Potential Reduction equation Electrode Potentials using Zinc as the Standard, EZn Accepted Electrode Potentials using Hydrogen as Standard, Eo EZn – Eo Predicted and Measured Cell Potentials Anode Cathode Equation for the cell reaction Predicted potential from experimental data Measured Potential.
