Electrolytic Cells IntroductIon In an electrolytic cell, electrical energy from a direct current external source drives a non-spontaneous reaction. Common external sources include batteries, fixed output AC Power Adapters, car engine alternators, and variable power sources. You will use a variable power source in today’s experiment. It is important to know which terminal is giving off electrons and which terminal is taking in electrons. The anode, or (-) terminal, gives off electrons. If the terminals are color-coded, it is black, and the cathode, or (+) terminal, will be red. If you are not sure which terminal is which, as is the case with unmarked AC Power Adapters, you can use a voltmeter to check. When the common (-) lead from the voltmeter is attached to the (-) output from the power source, a positive voltage is shown. If the common (-) lead from the voltmeter is attached to the (+) output of t he power source, a negative voltage is shown. Connect the voltmeter’s leads to the power source in a manner that gives a positive voltage. The terminal that is connected to the common (-) lead of the voltmeter is the source of electrons, the anode, of the power source. It is often important to know the number of electrons that is coming out of the power source. Chemists count with moles. There are a few relationships which permit simple electron counting. The counting unit for electrons is the coulomb (C). One coulomb is 6.24 x 10 18 electrons. The ampere (A) is defined as a current of one coulomb per second. C A= or A⋅ s = C s (1)
tion that takes place is: 2e- + 2 H2O(l)
When there are alternate possibilities for reactions, the ones with the lowest energy requirements will usually take place. So when you run the electrolytic cell, iodine is produced at the anode, and hydrogen gas is produced at the cathode. You You will note the volume of hydrogen produced, and use the gas laws to calculate the moles produced. You will titrate the iodine with a standardized solution of sodium thiosulfate, and calculate the number of moles of iodine produced. From this data, you can check the number of moles of electrons run through the circuit against the number of moles of products produced in the cell. In the second electrolytic cell, you will anodize aluminum. Aluminum is a very reactive metal, but it forms a self-pro tecting layer of aluminum oxide in contact with the air. This layer is about 2 nanometers thick. If a strip of aluminum is oxidized in an electrolytic cell, the following takes place: 2 Al(s) + 3 H2O(l)
6 e- + 6 H+(aq)
3 H2(g)
You will measure the area of the aluminum anodized, and by monitoring the current and the time, you will calculate the quantity of aluminum oxide produced. Given the density of aluminum oxide, you will calculate the volume produced. Knowing the area of t he aluminum metal, you will calculate the thickness of the aluminum oxide coat on the surface of the aluminum metal. One note of advice to help you keep the connections straight. Oxidation takes place at the anode; electrons are coming out. When you connect the cell to the power source, those electrons will be going into the power source. You You connect the anode of the cell to the cathode of the power source. Likewise, reduction takes place at the cathode; electrons are coming in. The cathode of the cell must be connected to the source of electrons in the power supply, the power supply’s anode.
(3)
In today’s experiment, you will set up two electrolytic cells. The first one will be a solution of NaCl and KI in water. You You will measure the current flowing and t he time, and calculate the moles of electrons produced. At the anode of the cell, the iodide ions will be oxidized, even though chloride ions are also present. 2 I-(aq)
Al2O3(s) + 6 H+(aq) + 6 e -
If the solution is acidified, the aluminum oxide forms a porous layer on the surface of the aluminum. This layer gives further protection against corrosion, and also allows the surface to absorb dyes. At the cathode of the cell, the following takes place:
If we monitor the amperage in a circuit with an ammeter, and record the time the circuit is on, we know the number of electrons flowing. The Faraday (F) is the number of coulombs per mole of electrons. C F = 96500 mol e − (2) Combining equations one and two, we have A⋅s − mol e = F
H2(g) + 2 OH-(aq)
The general rule is that in an electrolytic cell, the power supply is cross-connected to the cell. The anode of the power supply is connected to the cathode of the cell, and the cathode of the power supply is connected to the anode of the cell.
I2(aq) + 2 e -
At the cathode of the cell, even though both sodium and potassium ions are present, and could be reduced, the reac1
ExpErImEnt Work in pairs Slies: • 1 power supply, 1 ammeter • 1 buret, 1 magnetic stirrer • 1 square of parafilm • 1 aluminum can, 1 aluminum square • 1 piece of cheese cloth • 2 ml 1 M HCl solution • 50 ml standardized Na2S2O3 solution, 1 ml starch • 250 ml 3 M H2SO4 solution • dye, acid, as described in Part 2 From your drawer • 1-150 ml, 1-400 ml beaker • 1-10 ml, 1-25 ml graduated cylinder • 1 thermometer From the storeroom, a kit containing: • 1 glass tube/nichrome wire, 1 platinum metal electrode • 2 alligator clip leads, 1-12 inch connector wire • 1 stir bar, 1 stopwatch
25 ml graduated cylinder, plastic parts removed
To anode of power source (through the ammeter)
ape To cathode of power source
Adjacent black&red terminals black
ape electrons in
electrons out
red
Ammeter VDC
150 ml beaker
Pt wire
black red Power Supply Ammeter and Power Supply are not to scale. Only the connecting terminals are shown. Magnetic Stirrer
pa 1: Elelysis f KI Set up an apparatus as shown in the right column. Place 120 ml of deionized water in the beaker. Add 5 grams of NaCl, and 0.5 grams of KI to the beaker. Stir with the magnetic stirrer until all the solid dissolves. Pour the solution from the beaker into the 25 ml graduated cylinder until it is brim full. Pull a square of parafilm over the top of the graduated cylinder, invert it, and place it into the beaker. Remove the parafilm from the graduated cylinder with a forceps and discard it in the trash. Place the bent glass t ube with the nichrome wire sealed in it into the beaker, and lift the graduated cylinder so that the wire is inserted into the cylinder. Be careful that you do not lift the cylinder above the level of water in the beaker, but place it closer to the top of the liquid than to t he bottom of the beaker. If there is air (no more than 5 ml) in the graduated cylinder, note how much, so that you can subtract the amount from the final gas volume of hydrogen. Clamp the graduated cylinder to a ringstand. Pull the bent glass tube up and tape it to the graduated cylinder.
expeditiously as possible. Keep the power on for 10 minutes. Adjust the knob as required to keep the current at 0.20 amps. You will see bubbles collecting in the graduated cylinder and a color swirling off of the platinum as the electrolysis proceeds. If bubbles are swept out of the graduated cylinder, turn the stirrer speed down. If you do not see this, or if the processes are on the wrong wires, then there is a problem with the connections. Call the instructor over if you don’t know what the problem is. While the electrolysis is going, rinse and half-fill a buret with the sodium thiosulfate solution. After 10 minutes, turn off the power and the stopwatch. Record the time the power supply was on. Read and record the volume in the graduated cylinder. Remember, it is upside down, and you want the volume of the gas in the cylinder. Pull the cylinder out of the beaker, allowing the liquid to drain into the beaker. Record the temperature of the liquid in the beaker. Leave the wires in the beaker. Add the HCl to the beaker to neutralize the OH- ions formed, increase the stirrer speed a bit, and then titrate it with the sodium thiosulfate. The reaction in the titration is: I2(aq) + 2S2O32-(aq) → 2I-(aq) + S4O62-(aq)
Make sure that the power source is off. You will use the DC terminals on the right side of the power supply. Connect the wires as shown in the diagram of the apparatus. Note that the diagram does not show the details of the appearance of the ammeter or the power supply, only the connections.
After the solution has gone from a reddish color to a yellow, add 1 ml starch. Continue to carefully add sodium thiosulfat e until 1 drop turns the solution from the iodine/starch blue to colorless. Record the final volume in the buret. Discard solutions into the sink. Rinse the buret and return it to the shelf. Put the platinum wire back into the baggie, and place both the glass tube/nichrome wire and the platinum wire into the kit. Do not return them yet. Keep the power supply connected as is for Part 2. pa 2: Aizig Ali Cut the aluminum square into three strips. Degrease the three strips of aluminum by immersing them for about 30 seconds in a 50:50 (v:v) mix of 1 M NaOH and methanol located
Dip the wire into the solution in the beaker, making sure that the alligator clamp does not dip into the solution. Turn the stirrer on to a low speed, and keep it on for the rest of Part 1. Make sure that the stir bar does not bump into anything while the stirrer is on. Turn on the power supply so that the ammeter reads 0.20 amp. Each small line is 0.05 amp. Turn on the stopwatch. Do the current adjustment and the stopwatch turn-on as 2
in the fume hood. Use a forceps to remove the strips, and then rinse them with deionized water. Lay them on a paper towel. Arrange them to give maximum exposed surface area by fanning them and bending slightly so they touch only where they are clamped. Place the aluminum can in a 400 ml beaker, and add 250 ml of 3 M H 2SO4 to the can. Connect the clamps as indicated in the diagram below. Use a clamp on a ring stand to hold the lead with the aluminum strips so that they are no more than half-immersed in the acid in the can.
Rit dye into the beaker. Heat the dye solution so that it is at a gentle simmer, ready for the aluminum when it is done. Don’t let the dye solution boil away. Also, fill one large test tube half full of 3 M HCl solution. You will use this to check the ability of the protective Al2O3 coating on the aluminum strips to slow down corrosion. After 10 minutes, turn off the power supply and the stop watch. Record the time. Remove the strips from the cell, rinse them, and lay them down on a piece of paper towel. Pour the sulfuric acid back into the reagent bottle for reuse (not the usual procedure!), and rinse out the aluminum can.
Leave the connections to the ammeter and power supply intact. The clamp that was on the platinum wire goes onto the strips of aluminum. The clamp that was on the glass tube device goes onto the can. electrons out
Bend two of the strips into a u-shape, and place them in the hot dye bath. Allow them to simmer for 10 minutes.
To cathode of power supply
During the simmering, place the other strip into the test tube with the HCl, anodized end down. HCl both dissolves Al2O3 and reacts with aluminum:
To anode of power supply (through the ammeter)
Al2O3(s) + 6 HCl(aq)
electrons in Clamp directly on can
→ 2 AlCl3(aq) + 3 H2O(l)
and
Cheesecloth, slit halfway
2 Al(s) + 6 HCl(aq) Do not allow
→ 2 AlCl3(aq) + 3 H 2(g)
The second reaction cannot take place until the first reaction strips the oxide coat off of the aluminum. Notice how the hydrogen bubbles begin to form on one half of the metal, and then after a while, on the other half of the metal. Record your observations.
strips to touch the can
Remove the strips from the dye with a forceps. Rinse and dry the aluminum. Describe how the anodized part of the strip differs from the untreated part of the strip. Measure the dimensions of the dyed portions of the strips, and calculate the total area that was dyed. Assume that the area of the strip that you reacted with HCl was anodized to the same extent as the two dyed strips, and calculate the total anodized area of the aluminum. You will use this data along with the quantity of Al2O3 formed during the anodization to calculate the thickness of the coating.
Do not let the strips touch the can. Notice that the aluminum strips are attached to the cathode of the power supply. That makes them the anode in the electrolytic cell. Turn on the power supply and quickly adjust the current to 0.40 amps. Turn on the stopwatch. Keep the power on for 10 minutes. Adjust the dial on t he power supply occasionally to keep the current at 0.40 amps. Bubbles will appear on the can, and not on the strips. If you see bubbles on t he strips, the connections are reversed. Once you are sure the setup is correct, carefully lay a square of cheesecloth over the top of the can. Cut the square halfway through so that you can slip it around the wire and clamp. This keeps the bubbles of hydrogen gas from carrying minute bits of spray from the sulfuric acid into the room. Make sure that you don’t move the aluminum strips and cause contact with the can.
Take the strips home as a trophy. Dispose of the solutions from this last part of the experi ment into the sink. Return the kit with the stopwatch to the storeroom. Clean up your workplace, and then work up your data.
While anodization proceeds, set up a ring stand with 2 rings and a wire gauze. You will heat up a dye bath in a 100 ml beaker to dye two of the pieces of anodized aluminum. Place 30 ml of deionized water and 10 ml of concentrated
3
Name_________________________________________ Grade___________ Date ___________
Data T able: E lectrolysis of K I E lectri cal Measur ements Ammeter reading
A Time on stopwatch
min Time in seconds
Number of coulombs (equation 1, page 85)
Moles of electrons C (equation 3, page 85)
Moles of I2 and H2 mol (half-reactions, page 85)
s mol
T itr ation Data Molarity of Na2S 2O3
M Initial buret volume
Volume of Na2S 2O3 used
Moles of Na2S 2O3 ml
( V in liters ⋅ M )
ml Final buret volume
ml
Moles of I2 mol (1/2 mol of Na2S 2O3)
mol
°C Atmospheric pressure
Torr
G as Volume Data Volume of H2 in cylinder Vapor pressure of water (see table)
ml Temperature of solution Pressure of H2 Torr ( PH 2 = Patm − PH 2 O )
Moles of H2 Torr
n = PV RT
mol
R = 62.4 LTorr/Kmol (V in liters, T in K)
% C orrespondence molI 2 (titration) − molI 2 (electron mol I
counting)
x100
2 (electron counting)
molH 2 (gas %
law)
mol H
− mol H
2
(electron counting)
x100
2 (electron counting)
%
Data T able: A nodization of Aluminum Area of dyed aluminum cm 2
(both sides of two strips)
Total area of Al anodized (3 strips, so 3/2 times area dyed)
cm2
Ammeter reading
A Time on stopwatch
min Time in seconds
Number of coulombs (equation 1, page 85)
Moles of electrons C (equation 3, page 85)
Moles of Al2O3 mol (half-reaction, page 85)
Grams of Al2O3 (use molecular mass)
Volume of Al2O3 g
(Density of Al2 O 3
Thickness of coating Volume of Al 2 O 3
Area anodized
= 4.0
g cm
3
)
Thickness in micrometers
cm
-6
(µm = 10 m)
Describe the Al strip dissolving in HCl:
Describe the appearance of the dyed Al strip:
4
sec mol
cm3
µm
Va pesse table Temp/°C VP/Torr 15 12.8 16 13.6 17 14.5 18 15.5 19 16.5 20 17.5 21 18.6 22 19.8 23 21.1 24 22.4 25 23.8 26 25.2 27 26.7 28 28.3 29 30.0 30 31.8
