Electrochemistry – Determination of the Standard
Reduction Potential and Solubility Product Ksp
Written by Dr. Dansby-Sparks and Lauren Landers
Adapted from:
Dillard, C. R.; Kammeyer, P. H. J. Chem. Educ. 1963, 40, 363.
Ahn, M.; Reuland, D.; Chadd, K. J. Chem. Educ. 1992, 69, 74.
OBJECTIVES
• Build an electrode of known potential to use as a reference
potential
• Determine the Standard Reduction Potential, E°, for a half-cell
involving the reduction of Cu2+
• Become familiar with the use of the Nernst Equation for calculations
involving cell potentials
• Determine the Solubility Product, Ksp, for the dissolution of
insoluble Cu2S by determining the concentration of Cu+.
INTRODUCTION
Electrochemistry is a branch of chemistry dealing with the interconversion of
electrical and chemical energy. Therefore, electrochemistry deals with oxidation
and reduction reactions, or redox reactions where electrons are transferred.
Many real-world applications use electrochemical methods because of their ease
of use and low costs. For example, diabetics who routinely check their blood
sugar levels do so though enzyme mediated oxidation of glucose on disposable
electrodes. The oxidation produces an electrical signal that is used to determine
the concentration of glucose in the blood.
The individual oxidation and reduction reactions of an overall redox process are
often separated into two half-reactions. The anode is where the oxidation
reaction takes place and where there is a loss of one or more electrons. The
cathode half-cell is where the reduction reaction takes place and where the
electrons produced from the anode are consumed. An electrochemical cell is
said to be a galvanic cell (sometimes called voltaic cell) if the reaction proceeds
spontaneously. If an external electrical force is needed to drive the reaction in
the non-spontaneous direction, the cell is termed an electrolytic cell.
For example, think about rechargeable batteries. When they are needed to
power a device, the electrons from the anode of the battery pass through the
device and terminate at the cathode. Once the battery is consumed, you must
supply an external electrical current from an outlet charger to drive the
spontaneous reaction in the reverse direction in order to use the batteries
electrical potential energy again.
The two half-cells are chemically separated so that their chemical contents do
not mix. However, the half-cells are electrically connected through a conducting
wire and salt bridge. The wire allows for the transfer of electrons generated by
the anode half-cell to flow in the direction towards the cathode where electrons
are consumed. The salt bridge allows for net charge neutrality by transfer of
ionic charges to allow the system to remain charge balanced.
Consider the Zn anode in the diagram above. In this electrochemical cell, the
oxidation reaction occurs at the anode. Solid Zn metal from the electrode, with
an oxidation state of zero, is oxidized to form an aqueous Zn2+ species.
Therefore, the oxidation half reaction is:
Zn (s) à Zn2+ (aq) + 2e-
E° = -0.76 V
The two electrons produced by this reaction are then transferred through a wire,
pass through a voltage detection device (LabQuest in today’s experiment) and to
the cathode. In the process, the anode half reaction gains a 2+ charged zinc ion.
In order to remain charge neutral and avoid a build up of excess charge,
negatively charged anions from the salt bridge move toward the anode. If there
were no salt bridge, the excess charge produced at the anode would halt the
redox reaction. At the cathode half reaction, the electrons produced at the anode
are used to reduce AgCl according to the following reaction:
AgCl (s) + e- à Ag (s) + Cl- (aq) E° = 0.222 V
The cathode reduction reaction on the other hand produces a negatively charged
chloride anion that must be neutralized by the positive cation component of the
salt bridge.
Sometimes it is easier to draw the short-hand notation of galvanic cells as
follows. The anode reaction is always written on the left and cathode reaction on
the right. Using this convention, electrons flow from left to right. Species are
separated by phase boundaries with solid lines and the two half cells are
separated by a double line representing the salt bridge.
2 In short-hand notation, the electrode material is always listed on the outside. In
the example used in this introduction, both the anode and cathode electrodes
can be made from the solid reduced form of the species involved in the redox
reaction. However, this is not always the case and in some instances a nonreactive conducting material such as carbon or platinum is used to transfer
electrons to the redox active species in the half-cell but is not itself oxidized or
reduced.
Individual potentials for half-cells cannot be measured. The only way to measure
potential is to take the potential difference between two half-cells.
The symbol E° represents the standard reduction potential when all aqueous
solutes are at 1M concentrations and all gases at 1 atm. Since it is impossible to
measure the potential of a single half-reaction, a common zero point is needed.
The Standard Hydrogen Electrode, SHE (or more commonly the Normal
Hydrogen Electrode, NHE) was long ago chosen to be the standard zero
potential to which all other reductions potentials are referenced.
While the SHE was chosen as the standard, its use is difficult and more
commonly a silver/silver chloride reference electrode (Ag/AgCl) is often used as
the reference electrode. However, the standard reduction potential is always
referenced to SHE. Therefore it is always important to state what reference
electrode the potential is measured against.
The cell potential of a galvanic cell is the differential voltage of the individual
half-cell reduction potentials for the cathode and anode:
𝐸!"## = 𝐸!"#!!"# − 𝐸!"#$%
To calculate the E°cell of the galvanic cell used in the introduction, the standard
reduction potentials as referenced to SHE would be found in the table of
standard reduction potentials:
3 Electrochemical reactions are rarely at standard states of 1M concentrations and
1 atm for gases. Therefore, the Nernst equation allows us to relate actual
conditions of the cell in the non-standard state, E, to the standard reduction
potential, E°:
𝐸 = 𝐸∘ −
0.0592
log 𝑄
𝑛
Where n represents the number of electrons transferred and Q is the reaction
quotient, or actual concentrations and pressures of products over reactants.
The Nernst equation can be used for both the balanced redox reaction using Ecell
and the balanced redox reaction or for a half reaction using Ered and the balanced
half reaction.
Therefore, if we pull the standard reduction potentials from a table of values, and
measure the actual potential, Ecell, then the Nernst equation can be used to
determine the concentration information or visa versa.
For this lab you will construct a half-cell containing electrode made of Ag/AgCl
that will serve as a reference electrode to which a differential potential can be
measured. The composition of the reference electrode should be stable and will
allow for the accurate measure of potential difference between the cell for which
we will know the concentration and standard reduction potential and the unknown
half-cell of interest. The two cells will be connected directly by submerging the
entire miniature reference half cell into the half-cell of interest as can be seen in
figures one and two. The two cells will still be chemically separated though a
salt-bridge in the end of the reference electrode.
In addition, you will use the reference electrode you build to determine the
concentration of Cu+ ions in a saturated solution of the very insoluble Cu2S
compound. From this information you will calculate the solubility constant, Ksp,
of the compound and compare it to the known value.
4 PROCEDURE
CAUTION: WEAR EYE PROTECTION!
Note: Create the salt bridge in the glass tube first so it can be setting as you
work!
Building a Ag/AgCl Reference Electrode:
1. Obtain a small piece of silver wire, a copper wire, a strip of parafilm
tape, and a small glass tube that has a small diameter opening on one
end of the tube.
2. Take a small (smaller than a pea) piece of loose cotton and use one
end of your copper wire to gently push the cotton plug to the sealed
end of the glass tube so that the cotton is no higher than 0.5 cm from
the tip of the tube.
3. To form the salt-bridge - take the glass tube with cotton plug and place
the constricted end into a beaker of hot liquid NaNO3 – agar solution
(prepared by the instructor) until the agar level is above that of the
cotton in the tube.
4. Remove from agar solution and gently wipe the excess agar from the
outside of the tube and set the tube aside for ~10 minutes until the
agar has solidly gelled.
5. Meanwhile, locate the premade mixture of H2O : conc HCl : conc HNO3
in a 1:1:1 ratio located in one of the hoods (Please use extreme
caution as strong acids can cause burns on contact with skin).
Place one end of the silver wire in the solution and move/stir with the
wire for approximately 30 seconds or until the end of the wire has a
thin black film of AgCl.
6. Remove the wire from the acid solution and rinse it well with water into
the rinse waste beaker located in the same hood.
7. When the agar has set, use a glass transfer pipette to carefully and
slowly fill the glass tube with a solution of saturated KCl and AgCl
leaving approximately 0.5 cm of air at the top of the tube. Inserting the
pipette too far or evacuating the solution with a large force could
disrupt your agar salt-bridge.
8. Place the oxidized end of the silver wire into the filled glass tube so
that the oxidized portion is submerged in the inner solution of the
electrode. Be sure to leave approximately 2 cm of wire extending out
of the top of the electrode for connecting to the LabQuest.
9. Seal the Ag wire in place with the strip of parafilm wax by stretching
the film tightly around the top of the tube and immobilizing the wire
according to Figure 1. The wire should not shift and the seal should be
tight.
5 Building the Copper Half-Cell:
10. Obtain a small measurement vial and fill it with approximately 20 mL of
a 0.1 M CuSO4 solution.
11. Place a rubber stopper with two holes on top of your measurement
vial. Place your reference electrode through the one of the holes in the
stopper.
12. Take the copper wire you used in step 2 and wind it around the end of
a pen or pencil to form a coiled section approximately 1 cm long.
13. Clean the coiled end of the copper wire electrode in the conc. nitric
acid (HNO3) bath found in the hood (Please use extreme caution as
strong acids can cause burns on contact with skin). The solution
should become light blue/green with use, and will produce brown NO2
gas. Do not inhale the gases. The cleaned electrode should be
shinny copper.
6 14. Place the coil of copper into the other hole of the stopper so that the
wire extends out the top of the stopper while ensuring that the coil is
submerged in the CuSO4 solution.
15. Using another piece of parafilm, wrap the top of the copper wire above
the stopper to immobilize the copper electrode.
Determination of the Standard Reduction Potential, E°, of Cu2+ from CuSO4
Solution:
16. Obtain a LabQuest data logger from the instructor or TA. Attach a
differential voltage probe to the LabQuest.
17. Connect the two alligator clip leads to each other and click on the
Sensors tab at the top of the screen, click on Zero, and select the
Differential Voltage Probe. The voltage reading should now read
approximately 0.000 V.
18. Attach the Black clip to the Ag/AgCl reference electrode lead wire and
the Red clip to the Cu electrode lead wire.
19. Select Mode from the upper right-hand side of the LabQuest. Confirm
that Time Based Mode is selected from the drop-down menu. Select
10 samples/s for a length of 20 seconds and select OK.
20. Select the Start button to record 20 seconds worth of differential
voltage for the electrochemical cell.
21. After collection of the data, select Analyze, Statistics, Potential, from
the top left menu. Record the Mean Voltage Difference for the cell as
Ecell1 on the data sheet.
7 Determination of Solubility Product, Ksp, of Cu2S:
22. Remove stopper and electrodes from the measurement vial, rinse the
electrodes and set aside.
23. Empty the CuSO4 solution into the waste container. Rinse the
measurement vial with water and fill with approximately 20 mL of
saturated Cu2S solution and return the rubber stopper and electrodes
to the vial. Reconnect the Black clip to the Ag/AgCl electrode and the
Red clip to the Cu wire.
24. Press the Start button again and select Store when prompted.
25. The Ecell for the new cell will be saved as “Run 2”. Repeat step 21 to
determine the average Ecell for the 20 sec measurement. Record this
as Ecell2.
26. Turn off LabQuest and click on Discard data.
27. Remove the stopper and electrodes and dispose of the solutions into
the appropriate waste containers. Rinse all glassware and remove
salt-bridge/cotton plug and put into the trash. Return the copper and
silver wires to the appropriate storage container.
Data Sheet:
Ecell1:
Ecell2:
Eref Ag/AgCl electrode:
0.210 V
Ered1:
Ered2:
E°, Cu2+:
E°red Cu+:
% error:
Ered2-E°:
0.520 V
[Cu+]:
Ksp:
Calculations:
1. The potential difference between the Ag/AgCl reference electrode half-cell
and the Cu2+/Cu and Cu+/Cu half-cells are recorded as Ecell from the data
sheet. Therefore, to calculate the reduction potential of each half-cell, Ered, in reference to the reference electrode potential, Eref, use the following
equation and record as Ered1 and Ered2 on the data sheet:
𝐸!"## = 𝐸!"# −𝐸!"#
therefore
𝐸!"# = 𝐸!"## + 𝐸!"#
8 2. The reduction potential for the half-cell can be used in the Nernst equation
to calculate the E°, the standard reduction potential for the Cu2+ reduction
to metallic copper:
°
𝐸!"#
= 𝐸!"#! +
0.0592
1
log !!
𝑛
[𝑀 ]
3. Determine your percent error based on the accepted E° for Cu2+:
𝑃𝑒𝑟𝑐𝑒𝑛𝑡 𝑒𝑟𝑟𝑜𝑟 = 𝐴𝑐𝑢𝑡𝑎𝑙 − 𝐸𝑥𝑝𝑒𝑟𝑖𝑚𝑒𝑛𝑡𝑎𝑙
𝑥 100%
𝐴𝑐𝑡𝑢𝑎𝑙
4. The rearranged Nernst equation is used to determine [Cu+]:
𝐶𝑢! = 10
°
!!"#! !!!"#
×
!
!.!"#$
5. Therefore, to determine Ksp for Cu2S, 𝐶𝑢! 𝑆 𝑠 ↔ 2𝐶𝑢! 𝑎𝑞 + 𝑆 !! (𝑎𝑞),
The formula is:
𝐾!" = 𝐶𝑢! ! 𝑆 !!
Since 1 mol of Cu2S produces 2 mol Cu+ and 1 mol S2-:
𝐾!" = 𝐶𝑢!
!
𝐶𝑢!
2
9 Pre-Laboratory Assignment 1. How does the cell you will set-up in lab differ from the typical
electrochemical cell where the two half-cells are physically separate from
each other.
2. Thinking about the Nernst Equation, write a half-cell reaction in which the
value of n would be 2. What does the value of 2 represent?
3. In a half-reaction with the following equation:
MnO2 + 4 H+ + 2 e- à Mn2+ + 2 H2O
E° = 1.23
The oxidation state of Mn in MnO2 is =
Mn is oxidized or reduced
4. Using the half-reaction given above. Determine the reduction potential if
the cell is composed of:
[MnO2] = 0.25 M
[Mn2+] = 4.3 x 10-4 M
pH = 2.45
10 Post-Laboratory Assignment 1. A galvanic cell has the following composition:
Pt(s) l Sn4+ (1M), Sn2+ (1M) ll Cu2+ (1M) l Cu(s)
a. From the shorthand notation, draw the galvanic cell. Label the anode
and cathode and indicate which species are in the aqueous phase and
which are solids (the electrodes). Under each beaker, write the
appropriate half-rxn.
b. Use the standard reduction potentials (Sn4+ = 0.15 V, and Cu2+ = 0.34
V) in the table to calculate the standard cell potential for the galvanic
cell.
c. Calculate Ecell for the electrochemical cell above if the copper half
reaction was replaced with the Ag/AgCl reference electrode used in
this lab (Ered = 0.210 V)
11 2. What would have happened if the two clips (the red and black connectors)
on the labquest had been reversed in the experiment? Explain.
3. Please indicate what you liked or disliked about this lab. Do you feel you
understand Chapter 17 material better after this lab? Honesty is
appreciated (and acceptable).
12