S K I L L
F O C U S
Predicting
Performing and recording
Analyzing and interpreting
Electrolysis of Aqueous
Potassium Iodide
When an aqueous solution is electrolyzed, the
electrolyte or water can undergo electrolysis. In
this investigation, you will build an electrolytic
cell, carry out the electrolysis of an aqueous
solution, and identify the products.
Safety Precautions
Questions
Procedure
What are the products from the electrolysis of
a 1 mol/L aqueous solution of potassium iodide?
Are the observed products the ones predicted
using reduction potentials?
Predictions
Use the relevant standard reduction
potentials from the table in Appendix E, and
the non-standard reduction potentials you used
previously for water, to predict the electrolysis
products. Predict which product(s) are formed
at the anode and which product(s) are formed
at the cathode.
Materials
25 cm clear aquarium rubber tubing (Tygon® ),
internal diameter 4 – 6 mm
1 graphite pencil lead, 2 cm long
2 wire leads (black and red) with alligator clips
600 mL or 400 mL beaker
sheet of white paper
1 elastic band
3 toothpicks
3 disposable pipettes
2 cm piece of copper wire (20 gauge)
1 drop 1% starch solution
10 mL 1 mol/L KI
1 drop 1% phenolphthalein
9-V battery or variable power source set to 9 V
532 MHR • Unit 5 Electrochemistry
Make sure your lab bench is dry before carrying
out this investigation.
1. Fold a sheet of paper lengthwise. Curl the
folded paper so that it fits inside the 600 mL
beaker. Invert the beaker on your lab bench.
2. Use the elastic to strap the aquarium tubing to
the side of the beaker in a U shape, as shown
in the diagram.
graphite
electrode
wire
electrode
battery
9V
aquarium tubing
elastic band
KI solution (1 mol/L)
600 mL beaker
rolled-up paper
inside beaker
3. Fill a pipette as completely as possible with
1 mol/L KI solution. Insert the tip of the
pipette firmly into one end of the aquarium
tubing. Slowly inject the solution into the
U-tube until the level of the solution is within
1 cm to 2 cm from the top of both ends. If
air bubbles are present, try to remove them
by poking them with a toothpick. You may
need to repeat this step from the beginning.
4. Attach the black lead to the 2 cm piece of
wire. Insert the wire into one end of the
U-tube. Attach the red electrical lead to the
graphite. Insert the graphite into the other end
of the U-tube.
5. Attach the leads to the 9-V battery or to a
variable power source set to 9 V. Attach the
black lead to the negative terminal, and the
red lead to the positive terminal.
6. Let the reaction proceed for three minutes,
while you examine the U-tube. Record your
observations. Shut off the power source and
remove the electrodes. Determine the product
formed around the anode by adding a drop of
starch solution to the end of the U-tube that
contains the anode. Push the starch solution
down with a toothpick if there is an air lock.
Determine one of the products around the
cathode by adding a drop of phenolphthalein
to the appropriate end of the U-tube.
7. Dispose of your reactants and products
as instructed by your teacher. Take your
apparatus apart, rinse out the tubing,
and rinse off the electrodes. Return your
equipment to its appropriate location.
(b) the anode and the cathode
(c) the positive electrode and the negative
electrode
(d) the movement of ions in the cell
2. Use your observations to identify the
product(s) formed at the anode and the
product(s) formed at the cathode.
3. Write a balanced equation for the half-reaction
that occurs at the anode.
4. Write a balanced equation for the half-reaction
that occurs at the cathode.
5. Write a balanced equation for the overall cell
reaction.
6. Calculate the external voltage required to carry
out the electrolysis. Why was the external
voltage used in the investigation significantly
higher than the calculated value?
Conclusion
7. What are the products from the electrolysis
of a 1 mol/L aqueous solution of potassium
iodide? Are the observed products the same
as the products predicted using reduction
potentials?
Applications
8. If you repeated the electrolysis using aqueous
sodium iodide instead of aqueous potassium
iodide, would your observations change?
Explain your answer.
9. To make potassium by electrolyzing
potassium iodide, would you need to
modify the procedure? Explain your answer.
Analysis
1. Sketch the cell you made in this investigation.
On your sketch, show
(a) the direction of the electron flow in the
external circuit
Chapter 11 Cells and Batteries • MHR
533
Spontaneity of Reactions
You know that galvanic cells have positive standard cell potentials, and
that these cells use spontaneous chemical reactions to produce electricity.
You also know that electrolytic cells have negative standard cell potentials,
and that these cells use electricity to perform non-spontaneous chemical
reactions. Thus, you can use the sign of the standard cell potential to predict whether a reaction is spontaneous or not under standard conditions.
Sample Problem
Predicting Spontaneity
Problem
Predict whether each reaction is spontaneous or non-spontaneous under
standard conditions.
(a) Cd(s) + Cu2+(aq) → Cd2+(aq) + Cu(s)
(b) I2(s) + 2Cl−(aq) → 2I−(aq) + Cl2(g)
Solution
(a) The two half-reactions are as follows.
Oxidation (occurs at the anode): Cd(s) → Cd2+(aq) + 2e−
Reduction (occurs at the cathode): Cu2+(aq) + 2e− → Cu(s)
The relevant standard reduction potentials are:
Cu(s) E˚ = 0.342 V
Cu2+(aq) + 2e− Cd2+(aq) + 2e− Cd(s)
E˚ = −0.403 V
E˚cell = E˚cathode − E˚anode
= 0.342 V − (−0.403 V)
= 0.745 V
The standard cell potential is positive, so the reaction is spontaneous
under standard conditions.
(b) The two half-reactions are as follows.
Oxidation (occurs at the anode): 2Cl−(aq) → Cl2(g) + 2e−
Reduction (occurs at the cathode): I2(s) + 2e− → 2I−(aq)
The relevant standard reduction potentials are:
2Cl−(aq) E˚ = 1.358 V
Cl2(g) + 2e− I2(s) + 2e− 2I−(aq)
E˚ = 0.536 V
E˚cell = E˚cathode − E˚anode
= 0.536 V − 1.358 V
= −0.822 V
The standard cell potential is negative, so the reaction is
non-spontaneous under standard conditions.
Practice Problems
17. Look up the standard reduction potentials of the following
half-reactions. Predict whether acidified nitrate ions will oxidize
manganese(II) ions to manganese(IV) oxide under standard conditions.
MnO2(s) + 4H+(aq) + 2e− → Mn2+(aq) + 2H2O()
NO3−(aq) + 4H+(aq) + 3e− → NO(g) + 2H2O()
534 MHR • Unit 5 Electrochemistry
18. Predict whether each reaction is spontaneous or non-spontaneous
under standard conditions.
(a) 2Cr(s) + 3Cl2(g) → 2Cr3+(aq) + 6Cl−(aq)
(b) Zn2+(aq) + Fe(s) → Zn(s) + Fe2+(aq)
(c) 5Ag(s) + MnO4−(aq) + 8H+(aq) → 5Ag+(aq) + Mn2+(aq) + 4H2O()
19. Explain why an aqueous copper(I) compound disproportionates
to form copper metal and an aqueous copper(II) compound under
standard conditions. (You learned about disproportionation in
Chapter 10.)
20. Predict whether each reaction is spontaneous or non-spontaneous
under standard conditions in an acidic solution.
(a) H2O2(aq) → H2(g) + O2(g)
(b) 3H2(g) + Cr2O72−(aq) + 8H+(aq) → 2Cr3+(aq) + 7H2O()
Rechargeable Batteries
In section 11.1, you learned about several primary (disposable)
batteries that contain galvanic cells. One of the most common secondary
(rechargeable) batteries is found in car engines. Most cars contain a
lead-acid battery, shown in Figure 11.18. When you turn the ignition,
a surge of electricity from the battery starts the motor.
When in use, a lead-acid battery partially discharges. In other words,
the cells in the battery operate as galvanic cells, and produce electricity.
The reaction in each cell proceeds spontaneously in one direction. To
recharge the battery, a generator driven by the car engine supplies
electricity to the battery. The external voltage of the generator reverses
the reaction in the cells. The reaction in each cell now proceeds nonspontaneously, and the cells operate as electrolytic cells. All secondary
batteries, including the lead-acid battery, operate some of the time as
galvanic cells, and some of the time as electrolytic cells.
As the name suggests, the materials used in a lead-acid battery
include lead and an acid. Figure 11.19 shows that the electrodes in each
cell are constructed using lead grids. One electrode consists of powdered
lead packed into one grid. The other electrode consists of powdered
lead(IV) oxide packed into the other grid. The electrolyte solution is
fairly concentrated sulfuric acid, at about 4.5 mol/L.
Figure 11.18 A typical car
battery consists of six 2-V cells.
The cells are connected in series
to give a total potential of 12 V.
(+)
(−)
cell connector
cell spacer
Pb
PbO2
cell with electrolyte,
H2SO4(aq)
Figure 11.19 Each cell of a
lead-acid battery is a single
compartment, with no porous
barrier or salt bridge. Fibreglass or
wooden sheets are placed between
the electrodes to prevent them
from touching.
Chapter 11 Cells and Batteries • MHR
535
When the battery supplies electricity, the half-reactions and overall cell
reaction are as follows.
Oxidation (at the Pb anode): Pb(s) + SO42−(aq) → PbSO4(s) + 2e−
Reduction (at the PbO2 cathode):
PbO2(s) + 4H+(aq) + SO42−(aq) + 2e− → PbSO4(s) + 2H2O()
Overall cell reaction:
Pb(s) + PbO2(s) + 4H+(aq) + 2SO42−(aq) → 2PbSO4(s) + 2H2O()
You can see that the reaction consumes some of the lead in the anode,
some of the lead(IV) oxide in the cathode, and some of the sulfuric acid.
A precipitate of lead(II) sulfate forms.
When the battery is recharged, the half-reactions and the overall cell
reaction are reversed. In this reverse reaction, lead and lead(IV) oxide are
redeposited in their original locations, and sulfuric acid is re-formed.
Reduction (at the Pb cathode): PbSO4(s) + 2e− → Pb(s) + SO42−(aq)
Billions of
rechargeable nicad batteries
are produced every year. They
are used in portable devices such
as cordless razors and cordless
power tools.
Figure 11.20
cap
vent ball
cover
seal
Oxidation (at the PbO2 anode):
PbSO4(s) + 2H2O() → PbO2(s) + 4H+(aq) + SO42−(aq) + 2e−
Overall cell reaction:
2PbSO4(s) + 2H2O() → Pb(s) + PbO2(s) + 4H+(aq) + 2SO42−(aq)
In practice, this reversibility is not perfect. However, the battery can go
through many charge/discharge cycles before it eventually wears out.
Many types of rechargeable batteries are much more portable than a car
battery. For example, there is now a rechargeable version of the alkaline
battery. Another example, shown in Figure 11.20, is the rechargeable
nickel-cadmium (nicad) battery. Figure 11.21 shows a nickel-cadmium cell,
which has a potential of about 1.4 V. A typical nicad battery contains three
cells in series to produce a suitable voltage for electronic devices. When
the cells in a nicad battery operate as galvanic cells, the half-reactions and
the overall cell reaction are as follows.
Oxidation (at the Cd anode): Cd(s) + 2OH−(aq) → Cd(OH)2(s) + 2e−
Reduction (at the NiO(OH) cathode):
NiO(OH)(s) + H2O() + e− → Ni(OH)2(s) + OH−(aq)
core
positive tab
KOH or
NaOH
electrolyte
NiO(OH)
electrode
pressed
powdered
cadmium
electrode
Overall cell reaction:
Cd(s) + 2NiO(OH)(s) + 2H2O() → Cd(OH)2(s) + 2Ni(OH)2(s)
Like many technological innovations, nickel-cadmium batteries carry risks
as well as benefits. After being discharged repeatedly, they eventually
wear out. In theory, worn-out nicad batteries should be recycled. In
practice, however, many end up in garbage dumps. Over time, discarded
nicad batteries release toxic cadmium. The toxicity of this substance
makes it hazardous to the environment, as cadmium can enter the food
chain. Long-term exposure to low levels of cadmium can have serious
medical effects on humans, such as high blood pressure and heart disease.
separators
insulating
washer
can
A nicad cell has a cadmium electrode and
another electrode that contains nickel(III) oxyhydroxide, NiO(OH).
When the cell is discharging, cadmium is the anode. When the
cell is recharging, cadmium is the cathode. The electrolyte is a
base, sodium hydroxide or potassium hydroxide.
536 MHR • Unit 5 Electrochemistry
Figure 11.21
Section Summary
In this section, you learned about electrolytic cells, which convert
electrical energy into chemical energy. You compared the spontaneous
reactions in galvanic cells, which have positive cell potentials, with the
non-spontaneous reactions in electrolytic cells, which have negative cell
potentials. You then considered cells that act as both galvanic cells and
electrolytic cells in some common rechargeable batteries. These batteries
are an important application of electrochemistry. In the next two sections,
you will learn about many more electrochemical applications.
Section Review
1
Predict the products of the electrolysis of a 1 mol/L aqueous
solution of copper(I) bromide.
2
In this section, you learned that an external electrical supply
reverses the cell reaction in a Daniell cell so that the products are
zinc atoms and copper(II) ions.
I
I
(a) What are the predicted products of this electrolysis reaction?
(b) Explain the observed products.
3
Predict whether each reaction is spontaneous or non-spontaneous
under standard conditions.
I
(a) 2FeI3(aq) → 2Fe(s) + 3I2(s)
(b) 2Ag+(aq) + H2SO3(aq) + H2O() → 2Ag(s) + SO42−(aq) + 4H+(aq)
4
Write the two half-reactions and the overall cell reaction for
the process that occurs when a nicad battery is being recharged.
5
What external voltage is required to recharge a lead-acid
car battery?
6
The equation for the overall reaction in an electrolytic cell
does not include any electrons. Why is an external source of electrons
needed for the reaction to proceed?
K/U
K/U
K/U
7 (a)
(b)
I
Predict whether aluminum will displace hydrogen from water.
I Water boiling in an aluminum saucepan does not react with the
aluminum. Give possible reasons why.
8
MC Research the impact of lead pollution on the environment. Do
lead-acid batteries contribute significantly to lead pollution?
9
C Lithium batteries are increasingly common. The lithium anode
undergoes oxidation when the battery discharges. Various cathodes
and electrolytes are used to make lithium batteries with different
characteristics. Research lithium batteries. Prepare a report describing
the designs, cell reactions, and uses of lithium batteries. Include a
description of the advantages and disadvantages of these batteries.
Chapter 11 Cells and Batteries • MHR
537