Voltaic Cell Online Lab
Purpose: To understand how a voltaic cell works and what factors affect the cell potential.
Background: An electrochemical cell is produced when an oxidation reaction and a reduction
reaction occur spontaneously, and their resulting electron transfer between the two processes
occur through and external wire. The oxidation and reduction reactions are physically separated
from each other and are called half-cell reactions. A half-cell can be prepared from almost and
metal in contact with a solution of its ions. Since each element has its own electron
configuration, each element develops a different electrical potential, and different combinations
of oxidation and reduction half-cells result in different voltages for the completed
electrochemical reaction.
The standard reduction potential is the voltage that a half-cell, under standard conditions (1M,
1atm, and 25 oC), develops when it is combined with the standard hydrogen electrode, that is
arbitrarily assigned a potential of zero volts. A chart of reduction half-cell reactions, arranged in
order of decreasing standard reduction potential, shows the relative ease of reduction of each
substance listed. The more positive the reduction potential, the easier the reduction. A
spontaneous cell (a battery) can be constructed if two half-cells are connected internally using a
salt bridge, and externally using a metallic connector. In an electrochemical cell, the reaction
listed in the standard reduction potential chart with the more positive voltage occurs as a
reduction, and the reaction listed with the less positive voltage reverses and occurs as an
oxidation reaction. The cell voltage can be found by adding the voltages listed in the table, with
the value of the voltage for the oxidation reaction becoming the negative of its reduction reaction
voltage.
As an example, consider a cell made up of copper and aluminum half-cells.
Cu2+(aq) + 2e- Cu(s)
Eo= +0.34 V
Al3+(aq) + 3e-  Al(s)
Eo= -1.66 V
The copper reaction has the more positive potential and remains a reduction reaction. The
aluminum react with the less positive (more negative) potential is reversed and becomes an
oxidation reaction. Its potential is now an oxidation potential:
Al(s)  Al3+(aq) + 3e-
Eo= +1.66 V
The reduction potential and the oxidation potential are added to find the cell voltage:
3Cu2+(aq) + 2Al(s)  3Cu(s) + 2Al3+(aq)
Eocell = Eoreduction + Eooxidation
Eocell = 0.34 V + 1.66 V = 2.00 V
A positive value of Eocell indicates the oxidation-reduction reaction, as written, is spontaneous.
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A cell representation such as the following: Zn(s) | Zn2+(1.0M) || Cu2+(0.0010M) | Cu(s) means
that a cell is constructed of zinc metal dipping into a 1.0 M solution of Zn2+. The symbol “|”
refers to a phase boundary. The symbol “||” indicates a salt bridge is present between the zinc ion
solution and the copper ion solution. The anode is on the left (where oxidation occurs) and the
cathode is on the right (where reduction occurs).
The table of standard potentials assumes that all ion concentrations are 1.0 M, gas pressures are 1
atm, and temperature is 25 oC. Calculations for potentials under nonstandard conditions can be
made using the Nernst Equation:
E = Eo – (RT/nF)lnQ
Where E = the measured cell potential, Eo = the standard cell potential, R is the gas constant, T is
the temperature (K), n = the number of moles of electrons transferred as shown by oxidationreduction equation, and F is the Faraday constant (9.65 x 104 C/mol). Q is the reaction quotient:
the actual concentration of products and reactants substituted into the equilibrium constant
expression.
Using base 10 or common logarithms the expression can be written as:
E = Eo – (2.303RT/nF)logQ
Substituting for the constants 2.303, R, F, and using the temperature 298K the expression can be
simplified to:
E = Eo – (0.0592/n)logQ
A measurement of the cell potential, E, under non-standard conditions, can be used to calculate
the value of Q, which can then be used to determine unknown concentrations of ions actually
present in a solution.
Pre-Lab Questions:
1. Electrons flow from (anode/cathode) to (anode/cathode).
2. The theoretical output voltage is determined by the (sum/difference) of the two metals
and will be a (positive/negative number).
3. A reaction is (spontaneous/nonspontaneous) if the cell potential is positive.
4. What is the purpose of using the Nernst Equation?
5. Write the half reactions for Copper and Zinc in an electrochemical reaction. What is
being oxidized? What is being reduced?
6. Write the combined reaction and balance the equation.
7. What would the standard cell potential be in the reaction?
8. Solve to find the nonstandard cell potential for the reaction if the zinc ion solution is 0.05
M and the copper ion solution is 0.20 M. Temperature is 25 oC.
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Procedure: Use the online lab to find the cell potentials of three different cells. Once done at
standard state, change the concentrations of the anode and cathode to see the effects. Fill in the
data charts.
Click on or type in the following link to access the online lab.
http://www.kentchemistry.com/moviesfiles/Units/Redox/VoltaicCellEMF.htm
Data:
Table 1:
[Zinc ion solution] (M)
1.0
1.0
1.0
1.0
1.0
2.0
0.1
0.01
0.001
2.0
0.001
[Copper ion solution] (M)
1.0
2.0
0.1
0.01
0.001
1.0
1.0
1.0
1.0
0.001
2.0
Cell potential (E)
Table 2:
Metal/solution (1M)
Zinc
Zinc
Copper
Metal/solution (1 M)
Copper
Silver
Silver
Cell potential (E)
Table 3:
[Zinc ion solution(cathode)]
(M)
1.0
1.0
1.0
1.0
1.0
[Zinc ion solution(anode)] (M) Cell potential (E)
1.0
2.0
0.1
0.01
0.001
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Post-Lab Questions:
1. Use the Nernst Equation to verify three different cell potentials in Table 1.
2. Write the half reactions for each of the reactions in Table 2. For each, which is the
anode? Which is the cathode?
3. Write the combined reaction equation for each reaction in Table 2 and calculate cell
potential at standard state.
4. What concentrations of Zinc and Copper solution will produce the most voltage? Explain.
5. Explain how a Zinc/Zinc voltaic cell can have an electrical output as shown in Table 3.
6. What is the purpose of a salt bridge in a voltaic cell?
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