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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).