EXPERIMENT 23
ELECTROCHEMISTRY: VOLTAIC CELLS
INTRODUCTION
This experiment deals with cells in which spontaneous oxidation-reduction reactions can be used to produce
electrical energy. The reactants in the oxidation-reduction reaction are separated physically, so there cannot be a
direct electron transfer from one reactant to the other as there would be if the reactants were in physical contact.
Instead, the two reactants are joined by a wire through which electrons travel from one reactant to the other. This
flow of electrons through a wire constitutes an electric current. You should recognize one important difference
between voltaic and electrolytic cells: in the voltaic cell, electrons flow unassisted from one reactant to the other
because the reaction is thermodynamically spontaneous; in an electrolytic cell, a non-spontaneous reaction is forced
to occur by the application of an outside power supply (a battery) which “pumps” electrons through the electrolytic
cell.
FIGURE 23-1 shows a voltaic cell in which Cu2+ is reduced by Zn metal; the cell reaction is
Zn(s) + Cu2+(aq) → Cu(s) + Zn2+(aq)
(23-1)
Figure 23-1. A Cu-Zn voltaic cell.
This particular cell consists of:




a Zn strip immersed in 0.5 M ZnSO4 solution
a Cu strip immersed in 0.5 M CuSO4 solution
a salt bridge constructed by filling a U-shaped tube with 0.5 M K2SO4, plugging the ends of the
tube with cotton, inverting the tube, and immersing one arm of the tube in each beaker
a conducting wire through which electrons flow from the Zn electrode to the Cu electrode.
A voltmeter is inserted in the circuit in FIGURE 23-1 to measure the cell voltage. The value of the voltage (the
electromotive force or emf), can be regarded as a quantitative measure of the tendency for electrons to flow in the
cell, that is, the tendency for the electrons to be transferred from one reactant (Zn) to the other (Cu 2+).
EXPERIMENT 23
23-1
The overall reaction in a voltaic cell can be divided into two half-cell reactions, an oxidation half and a
reduction half, each one occurring at a different electrode:
Oxidation:
Reduction:
Overall Reaction:
Zn(s)  Zn2+(aq) + 2 e (at the anode)
Cu2+(aq) + 2 e  Cu(s) (at the cathode)
Zn(s) + Cu2+(aq) → Cu(s) + Zn2+(aq)
(23-2)
(23-3)
(23-4)
Thinking about the total reaction as the sum of two separate half-reactions, each one occurring in a half-cell,
suggests an easy way to construct other voltaic cells. We could simply replace one half-cell, say, Zn/Zn2+, with a
series of different half-cells.
The concept that voltaic cells consist of two half-cells also suggests that the measured cell voltage is the sum of
contributions from both half-cells. In mathematical language:
Etotal = Eoxidation + Ereduction
(23-5)
In this experiment you will construct several voltaic cells, measure their voltages, and then investigate the effect on
cell voltage of changing solution conditions.
TECHNIQUE:
Measuring the Cell Voltage
The instrument employed for measuring emf is called a voltmeter, and your MeasureNet workstation can serve
as a high-quality voltmeter. Measurement with this type of meter causes only an extremely small electrical current
to flow, so large salt bridges, as in FIGURE 23-1, are not needed to achieve accurate results; a piece of filter paper
saturated with 0.1 M K2SO4 is sufficient. When the clips from the voltage test leads are connected to the electrodes,
the voltage generated by the cell is displayed. If the connections are reversed, the magnitude of the voltage will be
the same, but the sign will be opposite. When you begin the experiment you must first determine how to connect
the voltmeter to the Cu-Zn cell in order to get a positive voltage reading. For this cell, you know that the Zn
electrode is the anode and the Cu electrode is the cathode. Determine and record which wire of the voltmeter must
be connected to the anode and which to the cathode to obtain a positive voltage reading. For the remaining cells,
you can use this information to determine which electrode is the cathode and which electrode is the anode, and,
consequently, in which direction electrons flow through the external cell circuit.
EQUIPMENT NEEDED
crucibles (2)
dropping bottle
filter paper strips
sandpaper
MeasureNet
voltage probe
EXPERIMENT 23
23-2
CHEMICALS NEEDED
Cu strip
Fe strip
Mg strip
Pb strip
Sn strip
Zn strip
0.1M CuSO4; copper sulfate
0.1M FeSO4; iron(II) sulfate
0.1M MgSO4; magnesium sulfate
0.1M Pb(NO3)2; lead nitrate
0.1M SnCl2; tin(II) chloride
0.1M ZnSO4; zinc sulfate
dil. HCl; hydrochloric acid
conc. NH3(aq); aqueous ammonia
0.1M K2SO4, potassium sulfate
0.1M Na2S; sodium sulfide
PROCEDURE
A. Cell Voltages
Press MAIN MENU on the workstation, enter the experiment number, then select VOLTAGE from the list of
measurement types, select VOLT v TIME, and press DISPLAY to monitor the voltage—the initial reading should
be 0.000V.
Pour ~10 mL of 0.1 M CuSO4 into one crucible and ~10 mL of any of the other solutions into another. Fill a
medicine dropper bottle with 0.1M K2SO4. Obtain the metal strips corresponding to the solutions you selected and
~10 strips of filter paper.
Clean the metal strips with sandpaper and rinse with distilled water. This cleans off any coating of metal oxide
on the surface so that the pure metal will be in contact with the solution. Fold each strip over the side of the beaker
containing the corresponding solution (Cu in CuSO4, etc) so that the strips are partially submerged in the solutions.
Attach the alligator clips of the voltage probe to the metal strips—the red clip should be attached to the copper strip.
Fold a strip of filter paper in half, open it, and place one end in each of the two half-cell solutions. Place a few
drops of K2SO4 solution on the filter paper so that it is saturated, and monitor the voltage on the MeasureNet
display.
Pour the solution that is not CuSO4 into the Liquid Inorganic Waste container (save the CuSO4 solution—
you’ll still need it) and rinse the crucible with distilled water. Return the used metal strip to the paper towel in front
of the container and obtain a different solution and metal strip. Construct a new half-cell, connect it to the
Cu/CuSO4 half-cell, and measure the voltage as above. Repeat this for all of the remaining metals and metal
solutions. For each cell identify the anode and the cathode. All solutions should be poured into the Liquid
Inorganic Waste container. Replace the metal strips on the paper towels in front of the proper containers after
rinsing them thoroughly with distilled water—do not put them back into the container.
B. Concentration Effects on Cell Voltage
All the solutions that you have used so far in this experiment have been 0.1 M. Next you will determine (at
least qualitatively) the effect on cell voltage of changing the salt solution concentrations.
Design experiments that will reveal the effects of varying the solution concentrations. Before you come to the
laboratory, your experiments should be thought out and written down in your laboratory notebook. As you design
your experiments, keep the following points in mind:
1. Use a Zn-Cu cell so that the results of different students can be easily compared.
EXPERIMENT 23
23-3
2. If your time is limited, the effect may be investigated for only the cathode (Cu/Cu 2+ compartment solution
instead of for both solutions.
3. Vary the concentration slightly from 0.1 M to see what effect this has. Then try larger variations such as 0.01
M, 0.001 M, and so on. In order to reduce the metal concentration to a very low value, add an equal volume of
0.1 M Na2S to the 0.001M solution of the metal salt. This will precipitate the metal sulfide and reduce the metal
ion concentration to less than 1030 M.
4. Try to determine whether the cell voltage varies in some regular way with concentration. For example, is the
variation linear? That is, does a tenfold decrease in concentration lower the voltage to one tenth of its former
value? And so on.
C. Complexation Effects on Cell Voltage
The blue color of the 0.1 M CuSO4 solution is actually due to the presence of a complex ion of copper,
[Cu(H2O)6]2+. Design and carry out an experiment that will show what effect (if any) is produced by having the copper
ions in the form of the ammonia complex [Cu(NH3)4]2+ instead of [Cu(H2O)6]2+. The hydrated ion can be converted to
the ammoniated ion by adding concentrated NH3 solution to a CuSO4 solution until an intense deep blue colored
solution results. Use this half cell to measure the voltages of 3 of the half cells from Part A.
For both of the above experiments that you design for yourself, your report should describe the experimental
procedures you use (this should be presented in the PROCEDURE section of the report)
RESULTS
1. Record the cell voltage data on the Chem21 REPORT SHEET.
2. Provide data tables summarizing your results for the concentration and complexation experiments.
3. For each cell for which you measured voltage, write the anode half-reaction and the cathode half-reaction. In
each case this involves determining which electrode is the anode and which electrode is the cathode.
4. For each cell add the two half-reactions to obtain the overall spontaneous cell reaction.
5. The potential (emf) of each cell is the sum of the contributions from each half-cell, but the two half-cell
contributions are impossible to measure separately. However, if we define that the Cu2+/Cu couple has Eo =
0.000 volts, we can assign the entire measured cell potential (emf) to the other couple. In order for the reaction
determined by experiment to be spontaneous to have a positive voltage, the entire measured positive voltage
must be assigned to the half-reaction occurring at the non-Cu electrode. Write the half-reaction that occurs at
each of the non-Cu electrodes and the voltage assigned.
6. The half-reactions to which you just assigned voltages are oxidations. In order to compare voltages on an
“equal” basis, it is customary to write all reactions as reductions. Do this for each of the half-reactions you used
and assign to each reduction half-reaction a voltage with the proper algebraic sign. This value is Eassigned.
7. Arrange the list of metals in order of their decreasing ability to donate electrons to Cu 2+; that is, list the metal
with greatest tendency to donate electrons first.
8. From your textbook or from a handbook obtain accepted values of standard half-cell potentials for the couples
used in this experiment. Next to the appropriate half-reaction written as a reduction, record this tabulated value
of Eo. For each half-reaction, record the difference between this value and the value you assigned for E in this
experiment; that is, record Eo - Eassigned.
EXPERIMENT 23
23-4
Experiment 23
REPORT SHEET
Name: _______________________________________
Date:__________
A. CELL VOLTAGES
Cell
Measured Voltage
I.
Zn/Zn2+  Cu2+/Cu
______________________
II.
Fe/Fe2+  Cu2+/Cu
______________________
III.
Mg/Mg2+  Cu2+/Cu
______________________
IV.
Pb/Pb2+  Cu2+/Cu
______________________
V.
Sn/Sn2+  Cu2+/Cu
______________________
B. CONCENTRATION EFFECTS ON CELL VOLTAGE
EXPERIMENT 23
23-5
C. COMPLEXATION EFFECTS ON CELL VOLTAGE
Half-reactions
I.
Anode
Cathode
Zn  Zn2+ + 2 e
Cu2+ + 2 e  Cu
II.
III.
IV.
V.
Overall Cell Reactions
I.
Zn + Cu2+  Cu + Zn2+
II.
III.
IV.
V.
EXPERIMENT 23
23-6
EXPERIMENT 23
REPORT SHEET (CONT.)
Name: _______________________________________
Reaction at Non-Cu Half-cell
I.
Date:__________
Assigned Voltage

Zn  Zn + 2 e
2+
II.
III.
IV.
V.
Reduction Half-Reaction
I.
2+
Zn
Assigned Voltage

+ 2 e  Zn
II.
III.
IV.
V.
List the metals in order of decreasing tendency to reduce Cu2+
EXPERIMENT 23
23-7
Tabulated E
Eassigned
Difference
Cu2+ + 2 e  Cu
______________
______________
______________
Zn2+ + 2 e  Zn
______________
______________
______________
Fe2+ + 2 e  Fe
______________
______________
______________
Mg2+ + 2 e  Mg
______________
______________
______________
Pb2+ + 2 e  Pb
______________
______________
______________
Sn2+ + 2 e  Sn
______________
______________
______________
Half-Cell Reaction
EXPERIMENT 23
23-8
Notes for Experiment 23
Helpful Hints:

Concentration effects: you should have at least three dilution data points (dilutions
should at least a factor of 10), including cell in which you add Na2S to the CuSO4
solution. Be sure to describe how you prepared your diluted solutions in the
procedure section on the Chem21 report sheet.

Complexation effects: you should have at least three data points. Add NH3(aq) to
0.1M CuSO4 1 mL at a time with stirring until a dark blue solution is obtained.
Ammonia is quite pungent! Do not leave the ammonia container uncapped!

.Rinse off your metal strips thoroughly, then place them on the paper towels
provided. Do not place wet electrodes back in the container!

.All solutions should be poured in containers marked Inorganic Waste.
5/10
EXPERIMENT 23
23-9