Chemistry 212 Lab, Fall 2002
Electrochemical Cells: Determination of Reduction Potentials for a Series of
Metal/Metal Ion Systems, Verification of Nernst Equation, and
Determination of Formation Constant of Cu(NH3)42+Aqua Complex
Purpose
There are three parts in today’s lab experiments. In part A, the purpose is to construct
simple electrochemical cells, and measure cell potentials. Through these measurements, it will be
possible to calculate the standard cell potential and confirm the observed "activity series" for some
of the common metals. In part B, you will verify the Nernst equation by measuring the cell
potential as a function of ion concentration in one of the half-cells. In part C, the measured cell
potential, after addition of NH3 as the complexing agent, in the Cu/Cu2+ half-cell will be used to
calculate the formation constant of a Cu(NH3)42+ aqua complex.
Introduction
You will need to rely heavily on the explanations of voltaic cells, concentration cells, and
the Nernst equation that are provided in your lecture text. There is simply not enough space to
cover all the necessary information here in the lab manual. This text is designed to help you to
build the cells, take the experimental measurements, and suggest ways to interpret and quantify the
information you will obtain while doing the experiment. For your information, an abbreviated
table of standard reduction potentials is included as part of this text.
Design of the Experiment: Standard and Non-Standard Reduction Potentials For
Metal/Metal Ion Half-Cells
Typical electrochemical cells are comprised of two half-cells, linked by a salt bridge
(allowing transport of ions in both directions), with a voltmeter completing the circuit to measure
the difference in voltage between the two halves. One half-cell has a reduction taking place (the
cathode) while the other half-cell has an oxidation taking place (the anode). The cell potential,
Ecell, as measured by a voltmeter is a measure of the free energy change, G = -nFEcell, that occurs
in the overall oxidation/reduction reaction.
Electrons will move in response to this potential difference, with the flow going toward
the cathode from the anode. Thus, the sign of the potential difference (as measured by a voltmeter)
is sufficient for identifying which half-cell has the oxidation reaction and which has the
reduction reaction.
Standard half-cell reaction potentials are measured using the standard hydrogen
electrode (SHE) as a reference point. The standard hydrogen electrode is assigned a value for
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potential of 0.00 V, and the half-cell reaction for the standard hydrogen electrode is the reduction
of 1.00 M aqueous acid to gaseous hydrogen at 1 atm pressure on a platinum electrode:
+
2H (aq, 1 M) + 2e  H2 (g, 1 atm)
All other half-cell reactions can be measured in tandem with the standard hydrogen electrode,
which gives rise to the table of standard reduction potentials, an abbreviated version of which
appears as part of this text.
Measurements of standard half-cell potentials are supposed to be carried out at standard
conditions i.e., at solution concentrations of 1.0 M and at 25°C. It is not always practical (or
possible) to carry out electrochemical reactions at these conditions, as is the case in today's
experiment. For all reversible reactions the cell potentials measured at non-standard conditions
(Ecell) can be related to potentials at standard conditions (E°cell) by the Nernst equation:
Ecell = Eocell - (0.0591/ n) log Q (Q = equilibrium quotient)
where, Eocell = Eooxidation + Eoreduction
Also, it is not practical to set up a standard hydrogen electrode for each of you to do your
measurements. Instead, you will use the potential of the Ag+/Ag reduction half-reaction as your
standard (+0.80 V). Using this value, and assuming it to be the cathode half-reaction, will allow
you to determine half-cell potentials for the other systems you will study (half-cells of Cu, Sn, Pb,
and Zn) whose general reduction half-reactions are of the form:
M2+ + 2e-  M
As an example of today's measurements, let us look at the half-cell reactions of Al/Al3+
and Sn/Sn2+ at standard conditions. The reduction half-reactions and potentials are as follows
(from table of standard reduction potentials):
Sn2+ + 2e  Sn
Al3+ + 3e  Al
Eo = 0.14 V
Eo = 1.66 V
To construct a complete electrochemical cell with a positive overall potential, it will be
necessary to consider the aluminum half-cell not as a reduction, but as an oxidation. This requires
writing the half-cell reaction shown above in reverse. It also requires changing the cell potential
magnitude and sign to equal, but opposite, that of the reduction half-reaction. The new halfreactions, and the overall reaction are shown below.
Oxidation (anode):
Reduction (cathode):
Overall Reaction:

(Al  Al3+ + 3e ) x 2
Eo = +1.66 V
(Sn2+ + 2e  Sn) x 3
Eo = 0.14 V
2Al + 3Sn2+  2Al3+ + 3Sn Ecell° = +1.52 V
What you are being asked to do today is in a sense related to this, with the silver half-reaction as
the standard potential. Taking this example further, you'd first have to correct for non-standard
conditions. In today's experiment, solution concentrations are 0.1 M (temperature effects will not
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be considered) and not the 1.0 M of "standard conditions". Thus the Nernst equation that would
be needed to "correct" the cell potential of the above example to non-standard conditions is:
Ecell = Eocell - (0.0591 / n) log ([0.1]2 / [0.1]3)
Where, n = 6 (the overall number of electrons required to balance the reactions)
Ecell = 1.52 - (0.0099)
= 1.51 V  This is what you should observe on the voltmeter
The above principle can be used to measure the standard potential of an unknown redox couple if
the cell potential is measured and if one half of the couple has known standard reduction potential
(in this case aluminum). Continuing the above example, if you know the "standard" reduction
potential for the aluminum half cell (-1.66 V from Table), you then measure the overall cell
potential with the voltmeter, and solve for the half-cell potential for tin. It plays out like this.
Assume the following to be true: (Oxidation and reduction are abbreviated as ox and red)
Standard reference half-cell potential: Eoox = +1.66 V (for the Al3+/Al half-cell as anode)
Measured overall cell potential: Ecell = +1.50 V
Solution concentrations: [Sn2+] = 0.01 M and [Al3+] =0.01 M
+1.50 = Eocell - (0.0591 / 6) log ([0.01]2 / [0.01]3)
Eocell = + 1.52 V = Eoox + Eored

Eored =  0.14 V (Non standard reduction potential for Sn2+ + 2e  Sn)
So…the actual objective is to determine half-cell potentials for Sn, Pb, Cu and Zn at non+
standard conditions, using the Ag /Ag half-reaction as your standard reference and
examining the all-possible combinations of half-cells.
Notes on Experimental Procedure and Precautions





There are 15 sets of pre-assembled and labeled weighing boat setup's located on the various
benches. Obtain one of these for use by your group of two students.
There are also several sets of solutions placed around the room on the benches. These are
used by those students occupying a given bench space, so please use what you need and leave
the sets where you found them so that others may also get what they need. No one should
have to cross the lab to find a reagent they need.
Treat the pH meters (being used here as voltmeters) with care, and return them to the cabinets
when you are finished.
Waste bottles for silver and lead salts generated during this experiment are in the hoods.
Please carefully follow the waste disposal directions.
You must not spill silver solution on your skin, clothes or on the bench; it may produce a
semi-permanent mark where spilled. Promptly clean and wash the spilled area with plenty of
water.
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Experimental Procedure
A. Measurement of Non-Standard Reduction Potentials
Place 5.0 mL of the following solutions into the labeled weighing boats. IT IS IMPORTANT TO
PUT THESE SOLUTIONS INTO THE CORRECT PLACE, OR YOUR RESULTS WILL BE
UTTERLY USELESS.
0.10 M KNO3 into the boat labeled "Salt Bridge"
0.10 M AgNO3 into the boat labeled "Ag"
0.10 M Zn(NO3)2 into the boat labeled "Zn"
0.10 M Cu(NO3)2 into the boat labeled "Cu-1"
0.01 M Cu(NO3)2 into the boat labeled "Cu-2"
0.001 M Cu(NO3)2 into the boat labeled "Cu-3"
Salt bridge: filter
paper soaked in
KNO3
Cu
Cell: Plastic
weighing cup
PH/ mV Meter
Ag
Schematic diagram of the electrochemistry cell setup.
To construct "salt bridges", dip one end of a piece of filter paper into the boat containing KNO 3
solution, and the other into a boat containing a metal ion solution. A total of 5 such pieces of filter
paper must start from the KNO3 boat, and end up in each of the 5 different solutions. Do not let the
salt bridges dry. Keep a medicine dropper handy to put a few drops of KNO3 on the filter paper
whenever necessary.
Obtain pieces of Cu, Ag, and Zn metals to use as electrodes. Place each metal strip into its
corresponding metal ion solution (copper metal into the copper nitrate solution, etc), bending each
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strip as needed to allow a portion to stick above the solution surface for potential measurements.
It is important to handle these strips as little as possible, and never with bare fingers.
Attach electrical leads to the pH meter by carefully inserting the connector and giving a slight
twist. Making sure the meter is set to "mV", turn on the meter and connect the two leads (alligator
clips) to one another. A voltage reading of zero should result. If not, use the calibration knob to
set the voltage to zero. Digital meters usually autozero by itself.
Attach one lead to the end of the strip of zinc metal, the other lead to the strip of silver metal and
make sure that the metal strips remain dipped in solution. If the voltage is negative, reverse the
leads to obtain a positive voltage. Measure the instantaneous cell voltage and record two more
readings immediately after that. Take the average of the three as the measured cell potential.
At this point, the silver electrode is acting as the cathode (i.e., reduction is occurring at the silver
metal). Leaving the connection to the silver electrode intact, move the other lead to the other
metals in any sequence you wish and take reading. Obtain voltage values for all possible
combinations of two metals.
Convert your readings to volts (if you have been taking them in millivolts) and use the Nernst
+
equation to find half-cell potentials for the Zn2+/Zn and Cu2+/Cu half-cells. Remember, Ag /Ag
half-cell as the reference, assuming it to be the cathode, with a potential of Eored = + 0.80 V
(+0.7993 V to be exact).
Table 1. Determination of Reduction Potentials.
Cell #
1
2
5
Electrode 1
(Cathode)
Reduction
Ag
Ag
Zn
Electrode 2
(Anode)
Oxidation
Zn
Cu
Cu
Ecell (V)
Measured
Eocell, (V)
Nernst
Eocell(V)
Standard
Explanations
 First, identify the anode (Ox) and cathode (Red) and place them in proper columns. Placement
shown in table may not be correct.
 EoCell Nernst is obtained from the Nernst equation, (EoCell Nernst = Ecell, Measured + (0.0591/n) log
Q). This is the non standard cell potential. You must write a balanced redox reaction to find n.
 EoCell Standard is the standard cell potential = Eoox + Eored. Standard potentials are listed in
Table 2 as obtained from table of standard reduction potentials in your textbook.
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Table 2. Calculated standard reduction potentials
Half Cell Reaction
Eored (V)
Non standard
Ag+ + e-  Ag

Cu2+ + 2e  Cu

Zn2+ + 2e  Zn
Eored (V)
Table value
+0.7993
+0.339
Percent error
na
-0.762
Note: Knowing the Eocell, Nernst from Table 1 and given the EoAg/Ag+ one can calculate the nonstandard reduction potential.
B. Verification of Nernst equation for Cu/Cu2+ half-cell by changing [Cu2+]
Measure cell potentials for all Cu/Cu2+ (0.1, 0.01, and 0.001 M) // Ag+ (0.10 M)/Ag cells. You
should have three measurements for three different copper concentrations (0.1, 0.01, and 0.001M).
For the overall cell reaction: 2Ag+ + Cu (s)  Cu2+ + 2Ag (s), according to Nernst equation:
Ecell = Eocell  (0.0591/2) log ( [Cu2+]/ [Ag+]2)
a plot of Ecell vs. log ( [Cu2+]/ [Ag+]2) should be a straight line with a slope = -0.0295 V (per tenfold
change in concentration) and an intercept, Eocell. By knowing standard reduction Eo for Ag/Ag+ and
the Eocell, you can calculate the standard reduction potential for Cu/Cu2+. Compare the calculated
value with that of the literature.
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C. Determination of Formation Constant, Kf, of Cu(NH3)42+ Aqua Complex
In this experiment you will use the cell: Ag/Ag+ (0.10 M) // Cu2+ (0.10 M)/Cu for the above
measurement. Measure the cell potential as you have done earlier to ensure the cell is working.
With a graduated cylinder, add 5.0 mL of 6.0 M NH3 into the copper solution to form a deep blue
complex. Stir the solution with a glass rod to mix well and measure the new cell potential. The free
[Cu2+] in the cell after complexation can be calculated from the cell potential by using the Nernst
equation.
Calculation of Kf
The formation of the complex is given by the reaction
Cu2+ (aq) + 4 NH3(aq)  Cu(NH3)42+ (aq)
Kf = [Cu(NH3)42+ (aq)] / ([Cu2+ (aq)] x [NH3(aq)]4)
Here, we have added excess NH3 so that initial [Cu2+] = [Cu(NH3)42+ (aq)] at equilibrium. The free
[NH3(aq)] can be obtained from mass balance and the free [Cu2+ (aq)] can be obtained from Nernst
equation as follows:
Initial mols (M x V)
Equilibrium mols
Equilibrium conc., M
Cu2+ (aq)
NH3(aq)
5.0 mL, 0.1 M
5.0 mL, 6.0 M
0.10 x 0.005 = 0.0005
6.0 x 0.005 = 0.03
To be calculated
0.03 - 4 x 0.0005 = 0.028
Use Nernst and Cell Pot 0.028/0.01 = 2.8
Total volume = 10.0 mL or 0.01 L
Assuming the following cell reactions:
Oxidation
: Cu  Cu2+ + 2e
Reduction
: 2Ag+ + 2 e  2Ag
Cu(NH3)4 2+ (aq)
0
0.0005
0.0005/ 0.01 = 0.05
Eo ox =  0.34 V
Eored = + 0.80 V
The measured cell potential is, Ecell = Eocell  (0.0591/2) log ([Cu2+]/[Ag+]2).
Given, Eocell = Eoox + Eored and [Ag+] = 0.10 M, one can solve for [Cu2+]. Use this value of [Cu2+]
along with other equilibrium concentrations in the equilibrium expression to calculate the K f.
Write your Kf value on the chalkboard for the class to copy. Compare your result with that of the
literature Kf = 6.6 x 1012.
Waste Disposal and Clean Up
Place two drops of concentrated HCl (or H3PO4) into the Pb and Ag solutions, and stir
briefly. Allow the solid formed to settle, and decant the solution into the drain. Place the
solid slurries of PbCl2 (or Pb3(PO4)2) and AgCl (or Ag3PO4 ) into their appropriate
containers in the hood. All other solutions may be rinsed down the drain. Return the metal
strips to the containers
near the balances
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