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 the 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 1018 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)
H2(g) + 2 OH-(aq)
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
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-protecting 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)
Al2O3(s) + 6 H+(aq) + 6 eIf 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:
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 the aluminum metal, you will calculate
the thickness of the aluminum oxide coat on the surface of
the aluminum metal.
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 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 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
will measure the current flowing and the 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)
I2(aq) + 2 eAt the cathode of the cell, even though both sodium and
potassium ions are present, and could be reduced, the reac-
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.
1
Experiment
Work in pairs
Supplies:
• 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)
To cathode of
power source
Adjacent
black&red
terminals
black
Ammeter
Tape
red
VDC
black red
electrons
in
Tape
150 ml
beaker
electrons
out
Pt
wire
Power Supply
Ammeter and
Power Supply
are not to scale.
Only the connecting
terminals are shown.
Magnetic Stirrer
Part 1: Electrolysis of 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 tube
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 the 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)
After the solution has gone from a reddish color to a yellow,
add 1 ml starch. Continue to carefully add sodium thiosulfate
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.
Part 2: Anodizing Aluminum
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
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.
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 H2SO4 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 stopwatch. 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:
Al2O3(s) + 6 HCl(aq) → 2 AlCl3(aq) + 3 H2O(l)
To anode of power supply
(through the ammeter)
electrons in
Clamp directly on can
and
Cheesecloth, slit halfway
2 Al(s) + 6 HCl(aq) → 2 AlCl3(aq) + 3 H2(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.
Do not allow
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 the 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 the 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 experiment 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 lectrical Measurements
Ammeter reading
Number of coulombs
(equation 1, page 85)
A Time on stopwatch
min Time in seconds
Moles of electrons
C (equation 3, page 85)
s
Moles of I2 and H2
mol (half-reactions, page 85)
mol
T itration Data
Molarity of Na2S 2O3
Volume of Na2S 2O3 used
M Initial buret volume
ml Final buret volume
Moles of Na2S 2O3
ml ( V in liters ⋅ M )
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 
PV 
n =


RT 
mol
R = 62.4 LTorr/Kmol
(V in liters, T in K)
% C orrespondence
molI 2 (titration ) − molI 2 (electron counting)
x100
mol I2 (electron counting)
molH 2 (gas law ) − mol H2 (electron
mol H 2 (electron counting)
%
counting)
x100
%
Data T able: Anodization of Aluminum
Area of dyed aluminum
(both sides of two strips)
Ammeter reading
Number of coulombs
(equation 1, page 85)
Grams of Al2O3
(use molecular mass)
cm2
Total area of Al anodized
(3 strips, so 3/2 times area dyed)
A Time on stopwatch
min Time in seconds
Moles of electrons
C (equation 3, page 85)
g
Volume of Al2O3
(Density of Al2 O 3 = 4.0
Thickness of coating
 Volume of Al 2 O 3 


 Area anodized 
Moles of Al2O3
mol (half-reaction, page 85)
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
cm2
sec
mol
cm3
µm
Vapor Pressure 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