CHM130 Electrolytic Cell Lab

Purpose: The purpose of this experiment is to determine values for the Faraday constant. By the end of this experiment, the student should be able to demonstrate the following proficiencies: 1. Construct an electrolytic cell. 2. Determine the number of moles of products formed in a redox reaction from experimental data. 3. Determine the total charge that has passed through an electrolytic cell. 4. Calculate values for the Faraday constant from experimental data.

Introduction: Spontaneous oxidation-reduction reactions are accomplished through a voltaic or galvanic cell. Non-spontaneous reactions are accomplished through an electrolytic cell. Electric energy is often used to decompose compounds into their elemental forms. For example, liquid NaCl can be used to produce liquid sodium metal and chlorine gas. 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) → I2 (aq) + 2 e- (oxidation) - cathode: 2 H2O (l) + 2 e- → H2 (g) + 2 OH (aq) (reduction) - - overall redox: 2 I (aq) + 2 H2O (l) → I2 (aq) + H2 (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. Hydrogen gas is formed at the cathode and can be collected in an inverted buret 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 (H2) and molecular iodine (I2) will be produced by the electrolysis.

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)

The relationship between the amount of electricity and the amount of chemical change is given by the Faraday (F) where 1 F = 96,485 C/mol e-

In this experiment hydrogen gas will be produced and collected and the volume measured. When electrolysis occurs in an aqueous solution, the oxidation and reduction of water are always possible reactions. These reactions are easily identified by gas evolution and pH changes in the solution near the electrode.

Reduction (cathode, -) half-reaction for water: 2H2O(l) + 2e- → H2(g) + 2OH-(aq)

Oxidation (anode, +) half-reaction for water: 2H2O(l) → O2(g) + 4H+(aq) + 4e-

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Materials The following glassware and hardware are The following chemicals will be used: needed: 3 M H2SO4 50.0 ml Buret Blue pipet bulb 250 ml Beaker Thermometer DC power supply Ampere meter Insulated Copper wires Meter stick or ruler

Procedure: Please show all calculations on the data sheet. The spreadsheet should be down loaded to the laboratory computer and opened from its hard drive during the experiment.

PART 1: Hardware set-up 1. Ensure the power supply is not connected to the outlet.

New Power Supply:

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Old Power Supply:

2. Connect a wire form the Black output on the power supply to the Black side of the Ampere meter. 3. Connect a red copper wire to the Red side of the Ampere meter and be sure that the other exposed end of the wire is inserted into the buret, which is inverted in a 250 ml beaker. 4. Then connect a wire with an alligator clip from the Red output on the power supply and using the alligator clip connect a carbon electrode. The carbon electrode is also placed in the 250 ml beaker. 5. Use a clamp to support the buret in the position in the 250 ml beaker.

PART 2: Preparing the solution and filling the buret. 1. Add 100 ml of RO water to the beaker and then 50 ml of 3 M H2SO4, mix well with a glass stirring rod. 2. The buret can rest on the bottom of the beaker. 3. To fill the buret use the blue pipet bulb, open the stopcock and use the bulb to draw in the acidic solution. If not sufficiently filled close the stopcock and prepare the bulb to draw in more solution the reopen the stopcock and complete the filling. The solution level should not exceed the graduations of the buret. Record the initial buret reading on the data page.

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PART 3: Beginning the electrolysis. 1. You will need to use a stopwatch. Try this site: http://www.online-stopwatch.com/ 2. Plug the DC power supply in then turn it on. 3. When the voltage is started begin to time the process. The current will be read from the ampere meter. 4. The current must be kept constant. The voltage may require some adjustments. Attempt a setting of 0.55 amps or as high as you can manage for Trial 1. 5. When approximately 30 mL of hydrogen has been collect turn off the power supply. Record the time. 6. Record the final buret reading on the data page. 7. Measure the temperature of the solution. 8. Measure the height of the solution remaining in the buret (see Figure below) 9. Repeat the production of Hydrogen three times using different current readings (lower than Trial 1 and higher than Trial 1). Refilling the buret each time from the same solution in the beaker.

PART 4: Calculations 1. Determine the pressure of the hydrogen gas. 2. Atmospheric pressure will equal the pressure of hydrogen after adjusting for the vapor pressure of water and the pressure of the column of solution. You can convert the height of the solution in the buret measured in millimeters to mm Hg by dividing it by 13.6. You must use the table in Appendix B of our text to find the vapor pressure water based on the temperature of the aqueous solution. Convert the correct value to atmospheres. Figure 3. Partial Pressures in the inverted buret: Page 5 of 5 CHM130 Electrolytic Cell

3. Convert the volume of hydrogen to liter. 4. Using the ideal gas law determine the number of moles of hydrogen. R = 0.08206 L atm/mol K 5. Calculate the moles of electrons needed to produce that number of moles of hydrogen. 6. Determine the number of coulomb produced. C= amp x sec. 7. Determine the number of coulombs per mole of electrons.

Clean-up: 1. With the plug disconnected, disassemble the apparatus. 2. All aqueous solutions may be disposed in the sink. Clean all glassware items and return them to their proper locations.