Purpose:
The purpose of Part 1 of this laboratory is to construct a table listing the
reduction potentials of a series of metal ions in order of ease of reduction. The
series of half-cells is constructed by placing a piece of metal into a 1.0 M
solution of its ions for each metal in the series. The metals are Cu, Fe, Pb, Mg,
Ag, and Zn. The half-cells are connected by a salt bridge constructed of a strip
of filter paper soaked in a solution of KNO3. The zinc half-cell is used as the
reference standard and assigned an E◦ of 0 volts, and all reduction potentials
are measured with respect to the zinc electrode. In Part 2, the Nernst equation
is applied to the voltage measurement of a cell with nonstandard copper ion
concentration. A solution of 0.0010 M Cu2 + is prepared, and the voltage of
the cell: Zn(s) | Zn2+ (1.0 M) || Cu2 + (0.0010 M) | Cu(s) is measured. The
measured voltage is compared to that calculated from the Nernst equation. In
Part 3, the solubility product constant of AgCl is determined from the Nernst
equation and the voltage of a cell in which the zinc half-cell is connected to a
solution containing Ag+ ions in a 1.0 M solution of NaCl.
Background:
An electrochemical cell is produced when a redox reaction occurs. The
resulting electron transfer between the reactions runs through an external wire.
The oxidation and reduction reactions are physically separated from each
other, so they are called half-cell reactions. A half-cell can be prepared with
almost any metal in contact with a solution of its ions. Each element’s own
unique electron configuration means that each element has a different electrical
potential, so different combinations of oxidation and reduction half-cells result
in different voltages for the completed electrochemical cell.
The standard reduction potential is the voltage that a half-cell, under
standard conditions (1 M, 1 atm, 25°C), develops when it is combined with the
standard hydrogen electrode that is arbitrarily assigned a potential of zero
volts. A positive E◦cell means that the reaction in that particular cell is
spontaneous.
A standard reduction potential chart is arranged in order of decreasing
standard reduction potential, showing the relative ease of reduction of each
substance listed. 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 is rewritten reversed 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 being the negative of its reduction reaction voltage.
Calculations of nonstandard potentials can be made using the Nernst
Equation: E=E◦−(RT/nF)ln(Q), where E is the measure cell potential, E◦ is the
standard cell potential, n is the # of moles of e- transferred as shown by the
redox reaction, R is the gas constant (8.314 J/mol·K), T is the temperature in
kelvins, n is the number of moles of electrons transferred as shown by the
redox reaction, and F is Faraday’s constant (9.65×104 C/mol).
At STP, the Nernst equation can be simplified to: E=E◦−(0.0592/n)logQ
Materials:
 Cu(NO3)2, 1.0 M, 3 mL
 Copper foil
 Fe(NO3)3, 1.0 M, 2 mL
 Iron nail, Fe
 Pb(NO3)2, 1.0 M, 2 mL
 Lead foil, Pb
 Mg(NO3)2,1.0 M, 2 mL
 Magnesium ribbon
 AgNO3, 1.0 M, 2 mL
 Silver foil
 Zn(NO3)2, 1.0 M, 2 mL
 Zinc strip
 NaCl, 1.0 M, 10 mL
 KNO3, 1.0 M, 10 mL
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Reaction plate, 24-well
Beral-type pipets,
graduated, 10mL
Graduated cylinder, 10mL
and 50mL
Filter paper
Wires and alligator clips
Voltmeter
Beaker, 50-m L
Wash bottle filled with
distilled water
Steel wool
CBL
TI Graphing Calculator
Safety Precautions:
Silver nitrate solution is toxic if ingested and irritating to body tissue. It
also stains skin and clothing. Lead nitrate solution is toxic if ingested or
inhaled; irritating to eyes, skin, and mucous membranes. Zinc nitrate solution
is slightly toxic if ingested; it is corrosive to body tissue. Cupric nitrate
solution is slightly toxic of ingested and irritating to skin, eyes, and mucous
membranes. Ferric nitrate solution is corrosive to body tissue. Magnesium
nitrate solution is a body tissue irritant. Wear chemical splash goggles and
chemical - resistant gloves and apron. Wash hands thoroughly with soap and
water before leaving the laboratory.
Procedure:
Part 1: Determine Reduction Potentials
1) Prepare a test cell to measure the voltage of the copper and zinc halfcells. Using a graduated Beral-type pipet, put approximately 2ml of 1.0
M ZnNO3 solution in one of the center wells of the 24-well plate. In
the new pipet, put approx. 2ml of 1.0M CuNO3 in an adjacent well.
2) Polish small strips of zinc and copper metal with steel wool and place
each metal in the appropriate well containing the solution of its ions.
3) Cut a sheet of filter paper into strips, and soak them in KNO3 solution.
The strips will act as salt bridges, so drape a fresh one across each well
for each measurement so that one end dips in the solution in each well.
4) Use a voltmeter the measure the potential difference between the two
half-cells. Connect the negative terminal of the voltmeter to the zinc
electrode. Make note as to which electrode is the anode and which is
the cathode. When the voltmeter reads a positive voltage, the electrode
connected to the positive terminal is the cathode and is undergoing
reduction, while oxidation is occurring at the electrode connected to the
negative terminal, the anode.
5) Prepare half-cells in other wells. Make a diagram of the the solutions in
each cell of the well plate.
6) Use a new pipet to put approximately 2mL of each solution into their
designated wells.
7) Polish the metals with steel wool so that they are shiny and insert them
into the well that contains the ion of the same metal.
8) Use fresh strips of filter paper soaked in 1.0M potassium nitrate as salt
bridges. The electrodes to be tested are:
Ag | A g+ Mg | Mg2 +0
Cu | Cu2 + Pb | Pb2 +
Fe | Fe3 + Zn | Zn2 +
9) Designate the zinc electrode as the standard electrode. Measure the
potential difference between the zinc electrode and each of the other
electrodes. Note which terminal is the anode and which is the cathode
in each cell. Record the data in Part 1 Data Table.
10) Measure the potential difference between at least six combinations for
the various electrodes. Record your data, including the equation for each
cell reaction.
Part II: Change Ion Concentration
11) Dilute the 1.0M CuNO3 to 0.0010M by counting 18 drops of distilled
water into a small test tube, and add 2 drops of the 1.0M CuNO3
solution.
12) Mix well by pouring the solution from one test tube to another several
times.
13) Repeat this dilution process two more times.
14) Pour approximately 0.2mL of this 0.0010M CuNO3 2 solution into one
of the wells in the well plate. Add a piece of polished copper foil,
measure the voltage against the standard zinc electrode, and record the
data.
Part III: Solubility Product of AgCl
15) Pour 10mL of 1.0M NaCl solution into a 50-ml beaker.
16) Add one drop of 1.0M AgNO3 to the NaCl solution and stir well.
17) Pour some of the solution into one of the wells in the well plate and add
a silver metal electrode. Measure the potential difference between this
half-cell and the zinc half-cell. Record the voltage in the Part 3 Data
Table.
Observations:
In Part 1, when the galvanic cells were created by making half-cells of
metal sample and a solution of its ions, no qualitative change was apparent. In
Part 2, when the CuNO3 was diluted by adding water, the color gradually
became a lighter blue. In Part 3, when AgNO3 was reacted with NaCl solution,
almost all of the silver ions combined with chloride ions to precipitate solid
AgCl.
Post Lab Questions:
1. What is electrode potential?
Electrode potential is the voltage generated when a half-cell is
connected to a standard half-cell, usually the hydrogen electrode, which
is assigned a potential of zero volts.
2. Did the ranking of reduction equations agree with that in a published
chart of E◦?
Yes; the ranking of reduction equations was the same as the
ranking in a published chart of E◦ values.
3. How should the values found using the zinc electrode as a standard
compare with those in the E◦ table that are based on the standard
hydrogen electrode? Did they?
The values found using the zinc electrode as a standard should
be greater than the values in the E◦ table based on the standard
hydrogen electrode by approximately 0.76V. This is because the zinc
electrode was used as the standard and assigned a voltage of 0.
However, when the hydrogen electrode is used as the standard, the
voltage for the zinc electrode is not zero – it is -0.76V.
4. What factors can cause a difference between experimental and reported
values?
Contamination of solutions would have produced a different redox
reaction than what was intended, which would have yielded a different voltage
with respect to the zinc electrode. This would have severely impacted
subsequent calculations concerning the predicted and measured cell potentials,
which are used in other parts of the lab. Concentration contaminations could
impact the cell reaction because a different amount of electrons would have
been transferred via the salt bridge, creating a different voltage. Another error
that was a factor was the voltmeter. Because the voltmeter showed a constantly
fluctuating reading, there was no clear moment in which the measurement
should have been taken. Therefore, the measurement for each cell varied with
respect to when each reading was recorded. Thus, human judgment caused the
reading to be variable. This unquestionably caused errors in the experimentally
determined voltage, which would have also impacted future calculations. Also,
since only one set of trials was used, there was no basis for comparison, so the
result could have varied widely.
5. What does a negative value for a standard potential indicate?
A negative value for a standard potential indicates that the cell is not
galvanic, meaning that the oxidation reaction is more likely to occur than the
reduction.
6. How did the change in concentration of the copper ions in Part 2 affect
the cell potential? Is this change in agreement (quantitatively) with that
which would be predicted by LeChatelier’s Principle? Did the
calculated and measure values agree?
The electrode potential depended on the concentration of substances
and the temperature (according to the Nernst equation). The concentration of
copper was reduced from 1.0M to 0.001M. For the 1M solution, the reduction
potential was +0.948V, whereas the 0.001M solution had a potential of
+0.87V. This is in agreement with LeChatelier’s Principle because the amount
of volts generated by the reaction was reduced when the concentration of
copper was reduced. The calculated and measured potential values differed by
approximately 0.078 volts.
7. Explain how the AgCl solubility product was determined.
The AgCl solubility was determined using the Nernst equation in
conjunction with the equation for Ksp. The unknown in this case, was the
concentration of Ag+. Therefore, with the standard and nonstandard measure of
potential known through the experiment, the concentration of Ag+ was found.
Because the amount of Cl- was known to be approximately 1, the Ksp was also
equal to the concentration of Ag+.
Conclusion:
The calculated voltages of Ag, Cu, Mg, Fe, and Pb using Zinc as the
standard were +1.359, +0.948, +1.261, +0.733, and +0.459, respectively.
These values were found by making galvanic cells consisting of half-cells of
each metal and a solution of its ions.
The experimentally measured reduction potential of the nonstandard
copper cell was +0.87V, while the theoretical potential is +1.1V. The Nernst
equation made this calculation possible by providing a mathematical
relationship in which Ecell = E◦ - the product of the gas law constant and the
temperature of the solution in Kelvins divided by the product of the number of
moles of electrons involved in the reaction and Faraday’s constant times the
natural log of the reaction quotient.
The solubility product constant (Ksp) value for AgCl was determined to
be 6.039×10−9. The actual value is 1.8 x 10-10, which equates to a percent error
of 32.55%. The voltage for the cell reaction was experimentally determined to
be 0.87V. Once again, the Nernst equation allowed the calculation of the
theoretical voltage of AgCl so that it could be compared to the voltage
measured in lab to calculate the percent error.