Experiment : Electrochemistry – Galvanic Cells
Chemistry 163
Introduction
Chemical reactions involving the transfer of electrons are called oxidiation-reduction
reactions or redox reactions. The broad field dedicated to the study of these reactions
is known as electrochemistry. A redox reaction always consists of an oxidation reaction
and a reduction reaction, hence the name. The simplest and most thorough way to
examine a redox reaction is to look at each of these individually. This method is known
as the method of half-reactions. One half-reaction describes the oxidation while a
second half-reaction describes the reduction. A solid piece of zinc dissolving into a
solution of zinc ions, Zn2+, is an example of an oxidation.
Zn (s)  Zn2+(aq) + 2eThe oxidation number of solid zinc is zero and the oxidation number of the aqueous zinc
ion is two. That means that during the reaction the oxidation number goes up, defining
an oxidation. As always, the chemical equation representing this process must be both
mass and charge balanced. In this oxidation there is one zinc atom on both sides of the
equation. The charge is also balanced because the 2+ charge on the zinc is balanced
by two electrons, 2e-. A reduction reaction is defined as a decrease in oxidation
number during the course of the reaction. An example reduction is the formation of
solid copper from copper ions in solution.
Cu2+ (aq) + 2e-  Cu (s)
In this reaction the oxidation number of the aqueous copper is two which decreases to
zero for the solid copper. There is one copper atom on each side of the equation and
the 2+ charge on the copper is balanced by two electrons, 2e-, only this time on the
reactant side of the equation. The redox reaction results when these two half-reactions
are combined, being careful to maintain both a mass and charge balance.
Zn (s) + Cu2+ (aq)  Zn2+ (aq) + Cu (s)
The electrons are not illustrated since they are present on both sides of the equation
and cancel out. In the redox reaction, the solid zinc is referred to as the reducing agent
since it reduces the copper. Likewise, the aqueous copper is called the oxidizing agent
since it oxidizes the zinc.
For comparison, all half-reactions can be expressed as reductions and the tendency of
undergoing this reduction is called the reduction potential, E. At standard conditions, 25
C and 1.0 M concentrations of all aqueous solutions, this reduction potential is defined
as the standard reduction potential, E. Standard reduction potentials have been
measured for numerous half-reactions and are listed in lengthy tables, a short list is
provided in the back of your textbook.
A galvanic cell or voltaic cell is a specific type of redox reaction that produces an
electric current spontaneously. In order for the flow of electrons generated by an
oxidation-reduction reaction to be useful, the current must be passed through an
electric load, such as a light bulb or motor. It is the design of the galvanic cell that
allows this to occur. In a galvanic cell, aqueous solutions of the oxidation and
reduction half-reaction are placed in separate compartments called half-cells. The
metal of each half-reaction is called an electrode, and is placed in the solution and
connected to an external wire. The electrode at which oxidation occurs is called the
anode and the electrode at which reduction occurs is called the cathode. The two
half-cells are connected by a salt-bridge that allows movement of ions from one halfcell to the other. When the two electrodes are connected to an electric load the circuit
is completed, and the oxidation-reduction reaction takes place.
Figure 1
Electric current results when electrons move from the anode (-) to the cathode (+). The
cell potential, Ecell, is then the combination of the half-cell reduction potentials given by:
Ecell = Ecathode - Eanode
At standard conditions the standard cell potential, Ecell, is based upon the standard
reduction potentials.
Ecell = Ecathode - Eanode
For the redox reaction above, standard reduction potentials are –0.763 V for zinc and
+0.337 V for copper. Zinc undergoes the oxidation and is therefore, the anode while the
copper reduction takes place at the cathode.
Ecell = +0.337 V – (-0.763 V) = +1.100 V
The positive voltage for the cell potential indicates that at standard conditions the
reaction is spontaneous, and that the reaction would produce an electric current when
set up as a galvanic cell.
When conditions are not at standard conditions, the Nernst equation can be used as a
correction. In the Nernst equation R is the universal gas constant with a value of
8.3145 J/Kmol, T is the temperature in K, n is the number of moles of electrons
transferred during the redox reaction, Q is the reaction quotient for the cell
(products/reactants), and F is the Faraday constant with a known value of 96,500
J/vmol.
 RT 
Ecell  E 0 cell - 
ln Q 
 nF 
For our example, Zn (s) + Cu2+ (aq)  Zn2+ (aq) + Cu (s):
Q
[Zn 2 ]
[Cu 2 ]
The Nernst equation then becomes:
2
 RT   [Zn ]  
Ecell  E cell - 
 ln
2
 nF   [Cu ]  
0
Procedure
Part 1: Galvanic Cells
1. Obtain the following items:
 copper strip and 5 mL of 0.1M copper(II)sulfate
 iron strip and 5 mL of 0.1M iron(II)sulfate
 lead strip and 5 mL of lead(II)nitrate
 zinc strip and 5 mL of zinc(II)sulfate
 10mL of potassium nitrate
2. Saturate 7 filter paper strips in potassium nitrate to be used as salt bridges.
3. Construct a galvanic cell with copper and zinc using 0.1 M CuSO4 as the aqueous
Cu2+ and 0.1 M ZnSO4 as the aqueous Zn2+. Measure and Record the voltage produced
by this galvanic cell.
4. Repeat this procedure for galvanic cells of Fe-Zn, Pb-Zn, Cu-Pb, and Fe-Pb. Use a
freshly prepared filter paper strip for the salt bridge of each cell.
Part 2: Concentration Cells
1. Prepare three dilute solutions of a 1.0 M CuSO4 stock solution by serial dilution.
2. Transfer 1.0 mL of the previous solution into a 100.00 mL volumetric flask and
dilute to the mark with de-ionized water. Be sure to rinse the pipet with the
solution to be transferred prior to use. Your solutions should then be:
Solution A – 1 M CuSO4 (stock solution).
Solution B – 0.01 M CuSO4
Solution C – 0.0001 M CuSO4
Solution D – 0.000001 M CuSO4
3. Prepare a half-cell of copper with the copper solution D.
4. Remove the zinc electrode from the zinc half-cell used previously and clean and
re-polish it.
5. Use 0.1 M ZnSO4 for the aqueous zinc. Connect the two half-cells with a freshly
prepared salt bridge. Measure and record the voltage.
6. Repeat the measurement of Ecell for the remaining three copper solutions of
known concentration, starting with the most dilute solution. Record the
measured values in your notebook.
Prelab: Write your procedure and any needed data tables into your notebook, and
complete the following.
1. Draw a picture representing a galvanic cell made from of Cu+2/Cu, Fe+2/Fe. Give a
detailed description of all the parts of the Galvanic cell.
2. Look up the standard reduction potential for Cu+2/Cu, Fe+2/Fe, Pb+2/Pb, and Zn+2/Zn.
Be prepared to use these to calculate theoretical Ecell values.
Post Lab: (In addition to reporting your data results, conclusions, learning, problems,
and errors, complete the following.
Part 1:
Using the method of half reactions write balanced redox reactions for the metal
combinations used in each galvanic cell. Display both your measured and
theoretical voltage next to each equation ( a data table may be a nice way to
present this data).
Part 2:
Using the Nernst equation calculate the cell potential, Ecell, of each galvanic cell
constructed (Ecell is the same for every calculation). Display both the measured
values and theoretical values next to each other for comparison.