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Experiment: Electrochemistry
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
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concentrations of all aqueous solutions, this reduction potential is defined as the standard
reduction potential, E. Standard reduction potentials have been measured for numerous halfreactions 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 half-cell 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.
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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):
[Zn 2 ]
Q
[Cu 2 ]
The Nernst equation then becomes:
2
 RT   [Zn ]  
Ecell  E 0 cell - 
 ln 
2
 nF   [Cu ]  
One of the oldest and most important applications of electrochemistry is to the storage and
conversion of energy. You already know that a galvanic cell converts chemical energy to work;
similarly, an electrolytic cell converts electrical work into chemical free energy. Devices that
carry out these conversions are called batteries. In ordinary batteries the chemical components
are contained within the device itself. If the reactants are supplied from an external source as
they are consumed, the device is called a fuel cell. The term battery derives from the older use
of this word to describe physical attack or "beating"; Benjamin Franklin first applied the term to
the electrical shocks that could be produced by an array of charged glass plates.
For portable and transportation applications especially, a battery or fuel cell should store (and be
able to deliver) the maximum amount of energy at the desired rate (power level) from a device
that has the smallest possible weight and volume. The following parameters are commonly used
to express these attributes:

Storage capacity or charge density (coulombs/liter or coulombs/kg)

Energy density (J/kg or watt-hour/lb)

Power density (watts/kg)

Voltage efficiency (ratio of output voltage to E°)

Lifetime: shelf-life (resistance to self-discharge) or charge/recharge cycles
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Physical limitations of battery performance
The most important of these are:

Effective surface area of the electrode. A 1-cm2 sheet of polished metal presents far less
active surface than does one that contains numerous surface projections or pores. All
useful batteries and fuel cells employ highly porous electrodes.

Current density of electrode surface. Expressed in amperes m–2, this is essentially a
measure of the catalytic ability of the electrode, that is, its ability to reduce the activation
energy of the electron transfer process.

Rate at which electroactive components can be delivered to or depart from the active
electrode surface. These processes are controlled by thermal diffusion and are inhibited
by the very narrow pores that are needed to produce the large active surface area.

Side reactions and irreversible processes. The products of the discharge reaction may
tend to react with the charge-storing components. Thermal diffusion can also cause selfdischarge, limiting the shelf life of the battery. Recharging of some storage batteries may
lead to formation of less active modifications of solid phases, thus reducing the number
of charge/discharge cycles possible.
Clearly, these are all primarily kinetic and mechanistic factors, which require a great deal of
experimentation to understand and optimize.
Procedure
Part 1: Galvanic Cells
1. Obtain the following items:
 copper strip and 1-2 mL of 0.1M copper(II)sulfate
 iron strip and 1-2 mL of 0.1M iron(II)sulfate
 lead strip and 1-2 mL of lead(II)nitrate
 zinc strip and 1-2 mL of zinc(II)sulfate
 10mL of potassium nitrate solution
2. Saturate 7 filter paper strips in potassium nitrate solution 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
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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, 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.
Part 3.
Suppose you are stranded on a boat somewhere in the ocean and you have to power up your GPS
(global positioning system) unit in order to find your way home. You have no replacement
batteries. How much voltage can you generate from items commonly found on your boat? Is it
enough to power you GPS?
Prelab Questions: (To be submitted at the beginning of lab)
1. 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.
2. Bring items to class that might be useful in constructing a battery that are also commonly
found on a boat.
Post Lab
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.
Part 3:
Write a discussion explaining to a non-scientist how a battery works. Use background
information (i.e. galvanic cells, your research on a selected battery, etc…) as well as
results and pictures from your self-constructed battery. This does not need to be technical
writing! Have fun with it!
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