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Electrochemical Cells PRE-LAB DISCUSSION In many redox reactions, there is a complete transfer of electrons from the substance being oxidized to the substance being reduced. If the electrons can be made to travel through an external conductor during this transfer, an electric current will be established in the conductor. This can be accomplished using an arrangement like the one shown in the figure. In this arrangement, the two half-reactions— oxidation and reduction—are carried out in separate vessels, called half-cells. The two half-cells are connected externally by metal wire attached to the two electrodes. In order to have a complete electrical circuit, ions -must be free to flow from one half-cell to the other. This is made possible by connecting the solutions in the two half-cells with a salt bridge. The complete system is called an electrochemical cell, or simply a chemical cell. electrode halfcells In this experiment, you will observe several electrochemical cells, using different combinations of metal electrodes. In each of these cells, the electrode consisting of the more active metal will be oxidized. This will be the oxidation half-cell. Electrons will flow from this electrode through the wire conductor to the reduction half-cell. There, the less active metal electrode will be built up due to reduction of ions of that metal. The relative activities of the various metals can be determined by checking their positions on the table of standard electrode potentials. A voltmeter will be used to detect the presence of an electric current through the conductor. This lab will aid in the understanding of redox reactions and electrochemical cells. PURPOSE: Set up and test the voltage of several different electrochemical cells. EQUIPMENT Glass cup porous cup D.C. voltmeter MATERIALS 0.5 M solutions of: Cu(NO3)2 Zn (NO3)2 Fe (NO3)2 safety glasses copper wire, insulated steel wool or emery paper (3) 100mL Beakers metal strips: Cu, Zn, Fe PROCEDURE: Copper Zinc 1. Using steel wool or emery paper, clean the strips of copper, zinc, and iron. 2. Obtain 50 mL of , Cu(NO3)2 Zn (NO3)2, and Fe (NO3)2 CuSO4 solution in the beakers. solution 2+ 2+ Part I Cu/Cu and Zn/Zn cell 3. Half fill a porous cup with Porous 0.5 M Cu(NO3)2 solution. Place cup a clean copper strip in the cup as shown in Figure. Zn SO4 solution 4. One third fill the glass with 0.5 M Zn (NO3)2 solution and place a clean zinc strip in the glass. 5. Using alligator clips, connect the wire leads to the metal strips as illustrated in the Figure. 6. Place the porous cup with the 0.5 M Cu(NO3)2 solution into the glass as shown. 7. Immediately touch the ends of the wire leads to the voltmeter terminals. If the voltmeter needle is deflected in the wrong direction, reverse the leads on the voltmeter. 8. Read and record the voltage immediately. Disconnect the leads from the metal strips. Part II Cu/Cu2+ and Fe/Fe2+ cell 9. Pour the 0.5 M Zn (NO3)2 solution from the glass back into its beaker. Wipe the zinc metal strip dry and rinse the glass with water. 10. One third fill the glass with 0.5 M Fe (NO3)2 solution and place a clean iron strip in the glass. 11. Repeat steps 5 thru 8. Part III Fe/Fe2+ and Zn/Zn2+ cell 12. Pour the 0.5 M Cu(NO3)2 solution from the porous cup back into its beaker. Wipe the copper metal strip dry and rinse the porous cup with water. 13. One third fill the porous cup with 0.5 M Zn (NO3)2 solution and place a clean zinc strip in the porous cup. 14. Repeat steps 5 thru 8. Part IV 15. Clean all beakers, glass and the porous cup with water. Rinse the metal strips with water and wipe them dry. When finished ask your teacher if you should return the solutions to their original bottles. Be extremely careful to return the solutions to the correct containers. OBSERVATIONS AND DATA Cell Voltage Cu/Cu2+ and Zn/Zn2+ volts Cu/Cu2+ and Fe/Fe2+ volts Fe/Fe2+ and Zn/Zn2+ volts CALCULATIONS Using the table of standard electrode potentials, write the oxidation half-reaction and the reduction halfreaction. Calculate the theoretical voltage for each electrochemical cell observed in this experiment. Cell Voltage (1) Cu/Cu2+ and Zn/Zn2+ (2) Cu/Cu2+ and Fe/Fe2+ (3) Fe/Fe2+ and Zn/Zn2+ CONCLUSIONS AND QUESTIONS 1. For each cell studied in this experiment, show: a. the overall molecular redox reaction. b. the oxidizing and reducing agents. 2. Identify the electrode where oxidation takes place and the electrode where reduction takes place for each cell studied in this experiment and indicate the direction of the flow of electrons. Identify the anode and the cathode. 3. Describe three conditions under which the voltmeter reading will be 0. 4. Discuss an electrochemical cell in which the half-cells are Ag/Ag+ and Cu/Cu2+. Write the oxidation and reduction half-reactions and the overall molecular redox reaction. Name the oxidizing and reducing agents and calculate the theoretical net electrode potential. 5. For each cell in the experiment calculate your percent deviation (error).