CHM 152
updated May 2011
Lab 12: Calculating Faraday’s Constant and Avogadro’s Number with
Electrolysis.
Introduction
A simple experiment for measuring the values of Faraday’s constant and Avogadro’s number employs the
electrolysis of an acidic solution, reducing hydronium ion to hydrogen gas.
2H3O+(aq) + 2e- Æ H2(g) + 2H2O(l)
or
2H+(aq) + 2e- Æ H2(g)
We can measure the hydrogen gas that’s produced during the electrolysis by filling a buret with the acid
solution and placing it upside-down in the reaction beaker. Placing the cathode inside the buret will allow
us to collect the gas and measure its volume by the amount of solution it displaces.
To determine the pressure of hydrogen, we first measure the height of the solution in the buret in units of
millimeters (hsol). Assuming the solution’s density is approximate to pure water, we can convert “mmH2O”
to mmHg by multiplying the column height by the ratio of their densities
hsol (mmHg) = hsol (mmH2O) x
1.00 g / mL
= hsol (mmH2O) x 0.0735
13.6 g / mL
The pressure of hydrogen (PH2) can then calculated by the following equation
PH2 = Patm - PH2O - hsol
where Patm is the barometric pressure during the reaction and PH2O is the vapor pressure of water at the
reaction’s temperature. From there, the ideal gas law can be used to find the moles of hydrogen produced,
which in turned can be used to determine the moles of electrons that were transferred during the reaction.
By measuring the current applied during the electrolysis and time elapsed during the reaction, we can
calculate the charge applied in units of Coulombs. Faraday’s constant (F) is defined as the number of
coulombs per mole and can be calculated as
F=
number of coulombs applied (based on current applied)
moles of electrons transferred (based on H2 produced)
Likewise, Avogadro’s number (N) is defined as the number of units per mole. Using the known charge of a
single electron, 1.60 x 10-19 C, we can determine the number of electrons applied during the reaction and
calculate Avogadro’s number as
N=
number of electrons applied (based on current applied)
moles of electrons transferred (based on H2 produced)
Concepts to Review
•
•
•
•
Oxidation-reduction reactions
Electrolysis
Ideal gas law
Calculations with amperes
Procedure
1.
Obtain a power supply and two cables that have alligator clips at one end. Also obtain two small
pieces of insulated wire (approximately six inches, with a small amount of exposed wire at each
end).
2. Completely Fill a 50 mL buret with 1 M sulfuric acid (to the top, not 0 mL).
3. Add 150 mL of 1 M sulfuric acid to a medium-sized beaker (250-400 mL).
4. Take a piece of insulated wire and insert one end in the buret (about 1-2 inches). Place a finger
over the buret opening, holding the wire in place, then carefully invert the buret and set it into the
solution in the beaker. Use a ring stand and clamp to hold the buret in place.
[Note: the sulfuric acid you’re using is about the same concentration as household vinegar. Your
finger will be fine; just rinse it off after completing this step]
5. Clip the other end of the insulated wire to the cable attached to the negative end of your power
supply (keep the alligator clip out of the solution). This will serve as the cathode.
6. Take the other insulated wire and clip one end to the cable attached to the positive end of your
power supply. Place the other end of the wire in the beaker. This will serve as the anode.
7. Turn on the power supply and adjust the voltage until hydrogen bubbles form in solution. Allow
the reaction to continue until solution level in the buret is below the 50 mL line then turn off the
power.
8. Record the initial volume of the solution and turn on the power, noting the starting time of the
reaction as well as the initial current being applied. Record the current in 5 mL intervals.
9. Turn off the power once 25 mL of hydrogen has been collected, noting the final volume of the
solution.
10. Using a meter stick, measure the height of the solution remaining in the buret, starting from the
surface of the solution in the beaker (not the buret opening).
11. Repeat this experiment two more times for a total of three trials.
[Note: the exposed wire can dissolve in the acidic solution over time. After each trial, check your
wire and use a razor or scissors to expose more wire, if needed. Replace the solution in the beaker
if it’s turned blue.]
Waste disposal
Sulfuric acid solutions can be poured down the drain with running water.
Vapor pressure of water
T (ºC)
20
21
22
23
24
PH20 (mmHg)
17.5
18.7
19.8
21.1
22.4
T (ºC)
25
26
27
28
29
PH20 (mmHg)
23.8
25.2
26.7
28.3
30.0
Name __________________________________
Data
Trial 1
Trial 2
Trial 3
1. Starting time of reaction
________
________
________
2. Ending time of reaction
________
________
________
3. Reaction time (s)
________
________
________
4. Initial buret volume (mL)
________
________
________
5. Final buret volume (mL)
________
________
________
6. Initial current (A)
________
________
________
after 5 mL
________
________
________
after 10 mL
________
________
________
after 15 mL
________
________
________
after 20 mL
________
________
________
after 25 mL
________
________
________
7. Volume of H2 produced (mL)
________
________
________
8. Solution temperature (°C)
________
________
________
9. Solution temperature (K)
________
________
________
10. Column height (mm)
________
________
________
11. Column height (mmHg)
________
________
________
11. Barometric pressure (mmHg)
________
________
________
12. Vapor pressure of water
at reaction temperature (mmHg)
________
________
________
10. Pressure of H2 (mmHg)
________
________
________
11. Pressure of H2 (atm)
________
________
________
12. Moles of H2 produced
________
________
________
13. Moles of electrons transferred
(based on H2 produced)
________
________
________
14. Average current applied (A)
________
________
________
15. Coulombs applied
________
________
________
17. Experimental value of
Faraday’s constant
________
________
________
Average
____________
18. Actual value of Faraday’s const.
____________
19. Percent error
____________
20. Total electrons applied
(based on coulombs applied)
________
________
________
21. Experimental value of
Avogadro’s number
________
________
________
Average
____________
22. Actual value of Avogadro’s no.
____________
23. Percent error
____________
Percent Error =
Exp-Actual
Acutal
x 100
Show your work for each of the following calculations from Trial 1
a) Partial pressure of hydrogen
b) Moles of hydrogen produced
c) Moles of electrons transferred (based on H2 produced)
d) Coulombs applied
e) Total electrons applied (based on coulombs applied)
f) Faraday’s constant
g) Avogadro’s number
Name: _____________________________
Section: ________
Post-lab Questions
1. How would each of the follow errors affect your calculated value of Avogadro’s number (incorrectly
high, low, or no effect)? Explain your answers.
a) The vapor pressure of water wasn’t factored into your calculations.
b) Part of the cathode was sticking out of the buret opening
c) 1.5 M sulfuric acid was used instead of 1 M.
d) The column height was not converted to units of mmHg
2. When measuring Avogadro’s and Faraday’s numbers by this method, what assumption is being made
regarding the current being applied (other than the amperage remaining constant)?
3. This experiment could also be performed using a chloride solution instead of an acid.
2Cl-(aq) Æ Cl2(g) + 2eWhat modifications, if any, would need to be done to the procedure if this substitution was made?
Name: _____________________________
Section: ________
Pre-lab Questions
1. Define the following terms.
a) Coulombs
b) Amperes
2. Explain the importance of each the following for this experiment.
a) Using deionized water to prepare the sulfuric acid solutions.
b) Making sure the exposed part of the insulated wire is fully inserted in the buret in step 4.
3. How many electrons are transferred by a 0.433 A current running for 13 minutes?
4. How many moles of electrons are required to produce 43 mL of 752 mmHg hydrogen gas at 21 °C?
5. The solution will sometimes turn blue during electrolysis. Based on your previous lab experience, what
ion is probably causing this? Could the presence of this ion interfere with the reaction being studied? Why
or why not?