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Experiment 21A
FV 04/20/
FARADAY’S LAW
MATERIALS: DPDT (double pole/double throw) switch, digital ammeter, J-shaped platinum electrode and holder, Hubbell plug with attached leads, carbon electrode, electrical lead, 10 mL graduated cylinder, 100 mL beaker, 400 mL beaker, buret, thermometer, starch-KI paper, starch solution, 0.20 M KI, 1 M H (^) 2SO (^) 4, 0.0200 M Na (^) 2S2O (^) 3.
PURPOSE: The purpose of this experiment is to determine values for the Faraday constant and Avogadro’s number.
LEARNING OBJECTIVES: By the end of this experiment, the student should be able to demonstrate the following proficiencies:
- Construct an electrolytic cell from a diagram.
- Determine the number of moles of products formed in a redox reaction from experimental data.
- Determine the total charge that has passed through an electrolytic cell.
- Calculate values for the Faraday constant and Avogadro’s number from experimental data.
DISCUSSION:
Forcing an electrical current through an electrolytic cell can cause a nonspontaneous chemical reaction to occur. For example, when direct current is passed through a solution of aqueous potassium iodide, KI, the following reactions occur at the electrodes:
anode: 2 I−^ (aq) → I 2 (aq) + 2 e−^ (oxidation)
cathode: 2 H (^) 2O (l) + 2 e−^ → H 2 (g) + 2 OH−^ (aq) (reduction)
overall redox: 2 I−^ (aq) + 2 H (^) 2O (l) → I 2 (aq) + H 2 (g) + 2 OH−^ (aq)
Electrons can be treated stoichiometrically like the other chemical species in these reactions. Thus, the number of moles of products formed is related to the number of moles of electrons that pass through the cell during the electrolysis. Iodine is formed at the anode in this electrolysis and dissolves in the solution upon stirring.^1 Hydrogen gas is formed at the cathode and will be collected in an inverted graduated cylinder by displacement of water. In this reaction, because the same number of electrons must pass through each electrode, the number of moles of iodide ion oxidized at the anode must equal the number of moles of water reduced at the cathode (i.e., in redox equations, electrons gained = electrons lost). Thus, equimolar quantities of hydrogen gas (H (^) 2) and molecular iodine (I (^) 2) will be produced by the electrolysis.
(^1) Iodine reacts with water according to the equilibrium below.
I 2 (aq) + H (^) 2O (l) ↔ HIO (aq) + H +^ (aq) + I−^ (aq)
For this reason, the pH of the solution must be adjusted to insure that I 2 is the predominant species in solution. For a more detailed explanation see: a) I. M. Kolthoff and R. Belcher, Volumetric Analysis, New York, Interscience (1957), Vol. 3, pp 214-215 and b) W. C. Bray and E. L. Connolly, J. Am. Chem. Soc. 33 (1911), 1485.
The current, or rate of flow of electricity, is measured in amperes, A. The ampere is the SI unit of current and corresponds to 1 coulomb of charge flowing for 1 second. Therefore, the total charge passing through the circuit, in coulombs, is equal to the product of the current in amperes and the time of current flow in seconds.
C = A × t (in sec)
In this experiment the following quantities will be obtained: (1) the total charge that has passed through the cell obtained from the average current and the elapsed time; (2) the number of moles of H 2 determined from the volume of gas collected; and (3) the number of moles of I 2 found by titration of the iodine with sodium thiosulfate (Na (^) 2S2O (^) 3). The number of moles of electrons that have passed through the cell can be obtained from the moles of hydrogen or the moles of iodine using stoichiometry. The total charge in coulombs that has passed through the cell can be obtained from the average current data. The value of the Faraday constant (F) can then be calculated from the total charge used in the electrolysis and the number of moles of electrons. Remember that 1 F is the electric charge, in coulombs, on 1 mole of electrons. Avogadro’s number can also be determined using the following procedure. The value of Avogadro’s number is the number of units in one mole. For the electrolysis, the required quantities for this calculation are: (1) the number of electrons and (2) the moles of electrons flowing through the electrolytic cell. As before, the moles of electrons can be determined stoichiometrically. The number of electrons can be determined from the total charge used, knowing that the charge on a single electron is 1.60 x 10 -19^ coulombs.
- Set the ammeter to the 100 mA setting. Simultaneously close the switch and record the initial starting time. Record the current, including units, at one-minute intervals in the DATA SECTION. Stir the solution frequently with the carbon electrode to dissolve the iodine which deposits on the electrode. Record your observations of the reaction in the cell in the DATA SECTION.
- When approximately 7 to 8 mL of hydrogen gas has been collected in the inverted graduated cylinder, open the switch and record the time. Note that the gas level must be below the liquid level in the beaker but still on the scale of the graduated cylinder.
Part B. Moles of H 2 formed.
- To equalize the pressure of the hydrogen gas inside the graduated cylinder with atmospheric pressure outside the beaker, raise or lower the graduated cylinder until the liquid levels inside the cylinder and in the beaker are the same. By carefully reading the scale on the inverted graduated cylinder, record the volume of hydrogen gas collected in the DATA SECTION.
- Record the temperature of the solution and the barometric pressure in the DATA SECTION. Obtain the vapor pressure of water at this temperature from an appropriate reference source, such as the CRC Handbook. Record this value in the DATA SECTION, being sure to include units.
Part C. Moles of I 2 formed.
- Remove the graduated cylinder and the J-electrode from the beaker. Rinse each into the beaker with a little distilled water. Do not remove the J-electrode from the wooden block.
- Add 10 mL of 1.0 M H (^) 2SO 4 to the solution in the beaker_._
- Fill a buret with standard 0.0200 M Na (^) 2S2O (^) 3. Record the initial buret reading in the DATA SECTION.
- Using the carbon electrode as a stirring rod, titrate the solution of I 2 in KI with the Na (^) 2S 2 O 3 until the color due to I 2 fades to a pale yellow. Then add two full droppers of starch solution to the solution being titrated. The solution will become dark blue due to the formation of a starch-iodine complex. If not, add more starch.
- Carefully continue the titration until the blue color just disappears. Record the final buret reading in the DATA SECTION. (Note that the total volume of Na (^) 2S2O 3 solution from the titration is needed, not just the volume after the addition of starch.)
- Do a second trial of the entire procedure if so directed by your instructor. Start completing the calculations on page E21A-6.
Clean-up:
- With the plug disconnected, disassemble the circuit.
- All aqueous solutions may be disposed in the sink. Clean all glassware items and return them to their proper locations.
Name Section
Partner Date
DATA SECTION Experiment 21A
Report all data with the proper number of significant figures and units.
Part A. Electrolysis and Quantity of Electrical Charge Used.
Trial 1
Initial Time _____________
Final Time _____________
Current
Average current (A) = ______________________
Total elapsed time (s) = _____________________
Observations of reaction: ________________________________________________________________________
Trial 2
Initial Time _____________
Final Time _____________
Current
Average current (A) = ______________________
Total elapsed time (s) = _____________________
DATA TREATMENT
Experiment 21A
Show your work for all calculations. Include the proper number of significant figures and units in your final answers. Try to complete these calculations before you leave lab.
Part A. Electrolysis and Quantity of Electrical Charge Used.
(A.1) Using the average current and the total elapsed time for the electrolysis, calculate the total electrical charge ( i.e ., number of coulombs) that passed through the cell during the electrolysis.
Trial 1: ___________________________ Trial 2: ___________________________
Part B. Moles of H 2 formed.
(B.1) From the volume of H 2 gas collected and the pressure of the dry H (^) 2, use the Ideal Gas Law to calculate the number of moles of hydrogen gas formed during the electrolysis.
Trial 1: ___________________________ Trial 2: ___________________________
Part C. Moles of I 2 formed.
(C.1) From your Na 2 S 2 O 3 titration data and the reaction stoichiometry, calculate the number of moles of iodine formed during the electrolysis.
titration reaction: I 2 (aq) + 2 Na (^) 2S2O 3 (aq) → 2 NaI (aq) + Na (^) 2S 4 O 6 (aq) ↑ ↑ electrolysis titrant product
Trial 1: ___________________________ Trial 2: ___________________________
Part D. Calculating the Value of the Faraday.
(D.1) From the moles of hydrogen gas formed at the cathode and the appropriate reaction stoichiometry on page E21A-1, calculate the moles of electrons that passed through the cell during the electrolysis.
Trial 1: ___________________________ Trial 2: ___________________________
(D.2) From the moles of iodine formed at the anode and the appropriate reaction stoichiometry on page E21A-1, calculate the moles of electrons that passed through the cell during the electrolysis.
Trial 1: ___________________________ Trial 2: ___________________________
(D.3) Using the values from (D.1) and (D.2), calculate the average number of moles of electrons that passed through the cell.
Trial 1: ___________________________ Trial 2: ___________________________
(D.4) Use the average moles of electrons that passed through the cell and the total charge that passed through the cell to calculate the Faraday, i.e., the number of coulombs per mole of electrons.
Trial 1: ___________________________ Trial 2: ___________________________
(D.5) Calculate the percent error between the value of the Faraday determined in this experiment and the accepted value. (The source for the accepted value must be properly referenced.)
Trial 1: ___________________________ Trial 2: ___________________________
Reference: ___________________________________
QUESTIONS
Experiment 21A
- Explain why the moles of electrons calculated in (D.1) and (D.2) should be approximately equal.
- The moles of electrons that passed through the cell during the electrolysis were calculated from two different measurements in (D.1) and (D.2). Which method of measurement might have more sources of experimental error? Explain. (In your explanation, consider the types of glassware and the techniques involved.)
- How would the calculated value of the Faraday be affected (larger or smaller) if the pressure of the hydrogen gas were not corrected for the presence of the water vapor? Support your answer (mathematically or otherwise).
Name ___________________________________________________ Section _________ Date ___________
Pre-Lab Exercise Experiment 21A
- In this experiment, you will be studying the electrolysis reaction occurring in an aqueous solution of potassium iodide: 2 I−^ (aq) + 2 H (^) 2O (l) → I 2 (aq) + H 2 (g) + 2 OH−^ (aq)
a. Using Table 18.1 in the textbook, determine the standard cell potential for this reaction.
b. Is this reaction spontaneous as written (under standard conditions)? YES NO
c. If the electrolysis process was carried out for 10 minutes at 120 mA, what was the total charge (in C) that passed through the cell?
- Hydrogen gas is collected during the electrolysis of aqueous potassium iodide. If 8.00 mL of gas was collected over water at standard temperature and pressure, how many moles of hydrogen gas were collected? The vapor pressure of water at this temperature is 23.8 mm Hg. Assume ideal gas behavior. (For a review of collecting a gas over water, see page 200 of your textbook.)
