Do this before you come to lab (have at the top of your notebook page to begin the lab):

1. Safety: Discuss the safety hazards associated with lithium ion batteries (a good demonstration of how lithium ion batteries work is linked here: https://www.energy.gov/energysaver/articles/how-does-lithium-ion-battery-work).
   1. Why/how do lithium ion batteries cause fires?
   2. How would you put out a lithium ion battery fire?
2. Purpose: Write a few sentences about the purpose of this lab: What is the “big picture” viewpoint of the lab experiment? What question am I trying to answer? What data will be collected to answer the questions? What technique will I use to obtain these data?
3. Preparation: An example galvanic cell is shown below.
   1. Label A-E.
   2. What is the salt bridge solution we are using in this lab?

A. Introduction

This experiment involves making several electrochemical cells (batteries) and measuring their potentials (voltages) as discussed in the case study.

The response of electrochemical cells to changes in chemical concentration is an important application of electrochemistry. Electrochemistry can be used to measure the concentrations of chemical species. For example, in Topic 10 you used a pH meter to measure pH. pH meters use the principles of electrochemistry to quickly and reliably measure the [H⁺] concentration of a solution.

pH will be measured in this experiment using the reversible oxidation-reduction of hydroquinone to quinone, which is described in Chemical Foundations: pH Measurements. The reduction half-reaction for this system can be written as:
C₆H₄O₂ (quinone) + 2H⁺ + 2e^– C₆H₄(OH)₂ (hydroquinone)

B. Your laboratory Challenge

Part 1: Properties of electrochemical cells

Use the following instructions to construct the four voltaic cells below in a well plate:
a) Zn_(s) | ZnSO₄ (0.1 M) || CuSO₄ (0.1 M) | Cu_(s)
b) Mg_(s) | Mg(NO₃)₂ (0.1 M) || CuSO₄ (0.1 M) | Cu_(s)
c) Fe_(s) | FeSO₄ (0.1 M) || CuSO₄ (0.1 M) | Cu_(s)

Fill one well near the middle of the well plate with 0.1 M KNO₃ to use as the salt bridge solution. Fill one nearby well with the solution common to all cells (e.g., CuSO₄ (0.1M)) and place a copper electrode in the well. You should sand the surface of the metal electrodes (not the graphite) before using. Place the remaining solutions in other nearby wells and place the proper electrode in each. You may choose to use the template supplied in lab to organize your well plate. Twist a strip of paper towel and place one end of the paper towel into the salt solution and one into the copper half-cell. Repeat for the other half cell of interest.

To start the software, go to Desktop → 412 → Voltaic Cells. Logger Pro should open and in a few seconds, the program should automatically detect the voltmeter with a potential reading. Attach the alligator clips to each other. Then under the Experiment section of the toolbar, select Zero to zero the instrument.

Attach the alligator clips to the electrodes of the electrochemical cell. When measuring cell potential, the positive terminal (red) must be connected to the cathode as indicated in the cell notation shown above. (See Chemical Foundations: Representation of an Electrochemical Cell)

Record the measure cell potentials and the average in table 1. Measure each cell at least twice (disconnect and reconnect the wires). If your results are not consistent, determine why and make the necessary corrections. Several things may create problems. You may have to clean the metal surface with sandpaper, the alligator clip may be in the solution, or the salt bridge paper towels may not be fully wet with the solutions.

Table 1: Voltages Measured for Each Electrochemical Cell

Calculate the standard reduction potentials using the average measured cell voltage and assuming a value of +0.340 V for the Cu²⁺ +2e^– → Cu reduction potential. Record standard reduction potentials, both found experimentally and the value reported in literature, in table 2.

Table 2: Standard Reduction Potentials  for Different Half-Reactions 

Part 2: Determination of pH

Three buffered solutions of known pH will be provided, each containing a small amount of quinhydrone. In addition, a buffer of unknown pH will be assigned. Be sure to record the letter of your unknown.
Determine the pH of the unknown buffer by measuring the potential of the following series of concentration cells:

a) C(s, graphite) | pH 7.0 buffer, quinhydrone || pH 1.0 buffer, quinhydrone | C(s, graphite)
b) C(s, graphite) | pH 7.0 buffer, quinhydrone || pH 4.0 buffer, quinhydrone | C(s, graphite)
c) C(s, graphite) | pH 7.0 buffer, quinhydrone || pH 7.0 buffer, quinhydrone | C(s, graphite)
d) C(s, graphite) | pH 7.0 buffer, quinhydrone || buffer unknown, quinhydrone | C(s, graphite)

Set up the above voltaic cells in a similar fashion to the Part 1 electrochemical cells. Use 0.1 M KNO₃ for the salt bridge solution. Measure each electrochemical cell twice. Record the measure cell potentials and the average in table 3.

Table 3: Voltages Measured for Each Electrochemical Cell

Construct a calibration curve of cell voltage versus ∆pH by using the experimental potentials and known buffer pHs. Use the calibration curve to determine the pH of the buffer unknown and clearly record the pH in your notebook.

Table 4: Cell voltage vs ∆pH

C. Communicating Your Results

For a permanent record, both your data and report should be written in your notebook.

1. Safety
2. Purpose
3. Preparation
4. Complete table 1
5. Complete table 2
6. Complete table 3
7. Complete table 4
8. Complete cell voltage vs ∆pH graph with linear fit
9. What was the identity, and the pH, of the unknown?
10. Part 1: How do the literature values of the standard reduction potentials for the electrochemical cells in Part 1 compare with those that you obtained experimentally? Briefly discuss reasons for the agreement or disagreement of these values.
11. Part 2: Compare the experimental value of the pH to the actual values.
    1. What is the percent error associated with the quinhydrone method of pH measurement?
    2. How does the slope of the calibration curve compare to the slope that is predicted using the Nernst equation (E=0.0592 x ∆pH)? Offer reasons why your calibration curve may not match the theoretical slope.
12. Write a brief conclusion paragraph: comment on whether or not you achieved the goal of the experiment, respond to the questions posed in the introduction.  Support your conclusion using your data.

• References

  1. Sherman, B. C., Euler, W. B., and Forcé, R. R. Journal of Chemical Education, 1994, 71, pp. A94 – 96.
  2. Chang, R. Chemistry, 7th edition, 2002, New York: McGraw – Hill, Inc.
  3. http://www.gm.com/company/gmability/adv_tech/300_hybrids/index.html
  4. Silberberg, M. S. Chemistry: The Molecular Nature of Matter and Change, 4th edition, 2006, New York: McGraw-Hill, Inc.