1. The document describes an experiment to test the validity of the Nernst equation by measuring the voltage of electrochemical cells containing varying concentrations of zinc or magnesium ions. 2. Results show cell voltage decreases with decreasing log of the concentration of zinc ions, following the linear relationship predicted by the Nernst equation. 3. A magnesium-copper cell is also constructed, producing enough voltage to power a small LED, demonstrating a spontaneous redox reaction and energy conversion.

1. Page 2
Physical Chemistry
2.5 Electrochemical Cells
Introduction
An electrochemical cell is a device capable of either generating electrical energy from chemical reactions
(termed Galvanic cells) or facilitating chemical reactions through the introduction of electrical energy
(termed Electrolytic cells). Below is an example of a Galvanic cell based on the following half-cell chemical
reactions:
Cu2+
(aq) + 2e- → Cu(s) Eº = +0.34 V
Zn2+
(aq) + 2e-→ Zn(s) Eº = -0.76 V
[Note: by convention standard reduction potentials are written as reduction reactions despite whether the
reaction occurs at the anode (oxidation) or the cathode (reduction)].
The theoretical cell voltage under standard conditions (Eºcell) can be calculated using the formula:
anode
o
cathode
o
cell
o
EEE 
Therefore, for the Zn-Cu cell:
Eºcell = (+0.34) – (–0.76) = +1.10 V
This formula is adequate for calculating theoretical cell voltages of cells operating under standard conditions
(i.e. 1 M reactants/products at 298K) but the concentrations (or, more correctly activities) of reactants and
products may not always be present at 1 M concentrations. In general, variation of the reactants and products
concentrations of a cell reaction will have a marked effect on the cell voltage, Ecell.
Consider a general, reversible redox reaction:
aA + bB → cC + dD
where, A, B etc are charged or uncharged species and a, b…. are the number of moles of the species in the
reaction. From thermodynamics, the potential, E, for this cell process is approximated by the Nernst
equation:












 ba
dc
cell
o
cell
BA
DC
nF
RT
EE
][][
][][
ln
Where Eº is the standard electrode potential or standard cell voltage, R is the gas constant, T is the absolute
temperature, F is Faraday’s constant (96,485 Coulombs mol–1) and n is the number of moles of electrons
participating in the balanced redox reaction (see P. Kavanagh lecture notes).
For the Galvanic cell: Zn + Cu2+ → Zn2+ + Cu the Nernst equation takes the form:












 

][
][
ln
2
10.1 2
2
Cu
Zn
F
RT
Ecell
and where the concentration of elemental zinc and copper is taken as unity.

2. Page 3
Physical Chemistry
In this experiment the validity of the Nernst equation will be tested. In addition, a Mg-Cu cell will be
constructed and used to power a 3mm red LED.
Mg2+
(aq) + 2e- → Mg(s) Eº = -2.37 V
Ohm’s law (V = iR) will be applied to calculate the current (i) and power (P = V.i) produced by the Mg-Cu
cell.
Experimental
1. Working in pairs, assemble the cell as follows: half-fill one compartment with 1 M solution of
zinc(II)sulphate and dip the zinc electrode into this solution. Half-fill the other compartment with 1 M
solution of copper(II)sulphate and dip the copper electrode into this solution. A salt bridge, a saturated
solution of potassium nitrate in agar, provides an electrical (ionic conduction) path between the two
compartments.
2. Connect a high-impedance voltmeter (multimeter) across the copper and zinc electrodes to measure the
cell voltage, Ecell. Record this value in the table.
3. By serial dilution, prepare a series of 10–1, 10-2, 10-3, and 10-4 M Zn2+ solutions from the 1 M Zn2+
(zinc(II)sulphate) stock solution using the 100 ml volumetric flasks provided (i.e. use 1 M solution to
prepare the 10-1 M standard, the 10-1 M solution to prepare the 10-2 M standard etc.).
4. Repeat step 1 for each of the Zn2+ solutions and measure the cell voltage (i.e. [Cu2+] is constant whilst
varying the [Zn2+] concentration). Record your values.
5. Repeat step 1 using 1 M solution of magnesium(II)sulphate in the anode compartment. Place the
magnesium electrode strip into the solution and record the cell voltage.
6. Connect the LED circuit (with 5 Ω resistor) across the cell and note the cell voltage.
Results and Discussion
Tabulate your data in table and copy into report.
Anode solution [Zn2+] Cell voltage (experimental) Cell voltage (theoretical)
1 M
10-1 M
10-2 M
10-3 M
10-4 M
Anode solution [Mg2+]
1 M
Plot a single graph of cell voltage, E in volts, (along the y-axis) versus ln{[Zn2+]/[Cu2+]} along the x-axis
(see example); determine the best straight-line and evaluate the slope and the intercept. Compare your results
with those predicted by the Nernst equation.
That is, at 298 K the slope should be equal to −(RT/nF) which is numerically equal to –(0.025693/n) V and,
the intercept, Eºcell = +1.1 V. (Note: n represents the number of moles of electrons from balanced redox
reaction i.e. 1 for 1 e-, 2 for 2 e-, etc..)
Questions
1. Write the balanced redox reaction between the copper and zinc species in the electrochemical cell. What
is the direction of spontaneous change for the redox reaction, justifying your answer with reference to
G for the process (hint: relate G to Eºcell)?
2. Using the Nernst equation, calculate the theoretical cell voltage (Ecell) for each of the assembled cells and
compare to experimental values. Explain any variations observed.
3. Compare cell voltages for the Zn-Cu and Mg-Cu cells. Explain any variation observed (hint: use the
‘electrochemical series’ to justify your answer).

3. Page 4
Physical Chemistry
4. With the aid of Ohms law, calculate the current (amps) and power (watts) produced by the Mg-Cu cell.
Compare the Mg-Cu cell voltage before and after connection of the LED circuit. Explain any variation
observed.