C2 7.1b – Student practical sheet
Electrolysis of copper chloride solution
Copper chloride melts at 620 °C, making it difficult to investigate the electrolysis of molten copper
chloride. Instead, you are going to investigate the electrolysis of copper chloride solution.
Aim
To investigate the electrolysis of copper chloride solution.
Equipment
●
eye protection
●
d.c. power supply and low voltage
●
electrolysis cell with carbon electrodes
●
leads with crocodile clips
●
teat pipette
●
copper chloride solution
●
blue litmus paper
Safety
●
Wear your eye protection.
●
Do not run the electrolysis cell for more than 5 minutes.
●
If you have asthma, tell your teacher and be very careful not to inhale any gas given off.
What you need to do
1
Set up the electrolysis cell.
2
Add enough copper chloride solution to cover the carbon electrodes but not to overfill it.
3
Connect the power supply.
4
Identify the positive electrode and the negative electrode.
5
Turn on the d.c. supply for no more than 5 minutes. Observe what happens at each electrode.
6
Test the gas collected at the positive electrode with blue litmus paper, dampened with water.
Record your observations.
Using the evidence
1
Identify the substances produced at each electrode. (2 marks)
2
Write a word equation and a balanced symbol equation for the electrolysis of copper chloride
solution, CuCl2 (aq). (3 marks)
Evaluation
3
Describe any problems you may have encountered when identifying the electrolysis products.
(1 mark)
Extension
4
Describe the properties observed for the substance produced at the positive electrode.
(2 marks)
5
Write the half-equation for the reaction that happens at:
a
the negative electrode (2 marks)
b
the positive electrode. (2 marks)
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257
C2 7.1c – Student worksheet
Electrolysis
Electrolysis of water
The diagram shows the Hofmann voltameter.
1
Identify the gases A and B. (2 marks)
2
Copy and complete these two half-equations:
a
H + e → H2
b
4 OH → 2 H2O + O2 + e
+
–
(1 mark)
–
–
(1 mark)
Oxidation and reduction
3
In terms of electrons, what are oxidation and reduction?
(2 marks)
4
At which electrode does reduction happen, the positive
electrode or the negative electrode? (1 mark)
Ions in molten ionic substances
5
6
Copy and complete the table to show the name of the product formed at each electrode.
(6 marks)
Molten
compound
Formula
Ions
Zinc chloride
ZnCl2
Zn
Lead bromide
PbI2
Pb
Aluminium
oxide
Al2O3
Al
Product at positive
electrode
2+
and Cl
2+
and I
3+
Product at negative
electrode
–
–
and O
2–
Copy these equations and correctly balance them. (6 marks)
+ e → Al
3+
–
a
Al
b
Cl → Cl2 + e
c
I → I2 + e
d
O
e
Pb
f
Zn
–
–
–
–
2–
→ O2 + e
–
2+
+ e → Pb
2+
+ e → Zn
–
–
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C2 7.1d – Student worksheet
Humphry Davy
Use the information about Humphry Davy’s experiments to answer the questions.
Humphry Davy
Humphry Davy used electrolysis to separate reactive metals from their compounds. He carried out
his experiments at the start of the 19th century. There were no power packs in those days, so
instead he used a type of battery called a voltaic pile. This had only just been invented 7 years
before.
Before Davy’s work, potassium and sodium compounds were known, but the free metals were not.
In 1807, Davy used his voltaic pile to pass an electric current through molten potassium hydroxide,
KOH. He was rewarded by seeing little bits of silver coloured potassium form, which burst into
flames when they reacted with oxygen in the air. Apparently he was so excited he danced around
his laboratory. Later that year, he discovered sodium by electrolysing molten sodium hydroxide,
NaOH.
1
Why did Davy use a voltaic pile in his experiments? (1 mark)
2
Explain why Davy had to use molten compounds, rather than solid ones. (2 marks)
3
Explain why hydroxide ions moved to the positive electrode in Davy’s experiments. (2 marks)
4
Correctly balance this half-equation for the reaction at the negative electrode:
K + … → K (1 mark)
+
5
Potassium ions were reduced to potassium atoms at the negative electrode. In terms of
electrons, what is reduction? (1 mark)
6
Hydroxide ions were oxidised at the positive electrode. In terms of electrons, what is oxidation?
(1 mark)
Extension
7
Write the half-equation for the reaction that occurs at the negative electrode during the
electrolysis of molten sodium hydroxide. (2 marks)
8
Correctly balance this half-equation for the reaction at the positive electrode:
OH → H2O + O2 + e (2 marks)
–
9
–
Most of the potassium formed during Davy’s electrolysis experiments burnt in air, forming a
lilac flame. Correctly balance this equation:
K + O2 → K2O (1 mark)
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259
C2 7.2b – Student practical sheet
Electrolysis of solutions
The electrolysis of solutions produces different substances, depending on the reactivity of the metal
and the nature of the non-metal ions involved.
Aim
To investigate the electrolysis of various ionic compounds in solution.
Equipment
●
eye protection
●
copper sulfate solution (0.5 mol/dm )
●
low voltage d.c. power supply
●
magnesium nitrate solution (0.5 mol/dm )
●
leads with crocodile clips
●
potassium iodide solution (0.5 mol/dm )
●
electrolysis cell with graphite
electrodes
●
sodium chloride solution (0.5 mol/dm )
●
sodium sulfate solution (0.5 mol/dm )
3
3
3
3
3
Safety
●
Wear your eye protection.
●
Toxic and harmful substances are involved. Do not run the electrolysis cell for more than
5 minutes.
What you need to do
1
Make a suitable table to record your observations at each
electrode.
2
Set up the electrolysis cell.
3
Add enough of one of the test solutions to cover the graphite
electrodes but do not overfill it.
4
Connect the power supply.
5
Identify the positive electrode and the negative electrode.
6
Turn on the d.c. supply for no more than 5 minutes. Observe
what happens at each electrode.
7
Clean the beaker and electrodes.
8
Repeat steps 3 to 7 for each of the remaining test solutions.
Using the evidence
1
In which experiments was a metal deposited at an electrode? Which electrode was this?
(2 marks)
2
In which experiments were gases released only? (2 marks)
3
In which experiments was a colour change observed? (1 mark)
Evaluation
4
Describe any patterns you may notice between the solution used and the electrolysis products.
(2 marks)
Extension
5
Explain why an a.c. power supply was not used. (2 marks)
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C2 7.2c – Student worksheet
Electrolysis rules
At the negative electrode
Metal ions and hydrogen ions are positively charged. They
are attracted to the negative electrode where they may be
discharged. During the electrolysis of solutions of ions,
hydrogen ions will be discharged unless the metal is less
reactive than hydrogen.
1
Use the reactivity series shown here to decide which
substance is produced at the negative electrode during
the electrolysis of the following solutions:
a
copper chloride. (1 mark)
b
magnesium nitrate. (1 mark)
c
sodium sulfate. (1 mark)
d
gold chloride. (1 mark)
e
zinc nitrate. (1 mark)
At the positive electrode
Non-metal ions are usually negatively charged. They are attracted to the positive electrode where
they may be discharged. If the solution contains halide ions (chloride, bromide or iodide ions),
these will be discharged as the corresponding halogen (chlorine, bromine or iodine). Otherwise,
oxygen will be released.
2
Name the substance produced at the positive electrode during the electrolysis of the following
solutions:
a
copper chloride. (1 mark)
b
potassium iodide. (1 mark)
c
sodium sulfate. (1 mark)
d
lead nitrate. (1 mark)
e
sodium bromide. (1 mark)
Predicting products
3
Predict the product formed at each electrode during the electrolysis of the following solutions:
a
copper nitrate. (2 marks)
b
sodium iodide. (2 marks)
c
potassium sulfate. (2 marks)
d
magnesium chloride. (2 marks)
e
silver nitrate. (2 marks)
f
sulfuric acid. (2 marks)
g
hydrochloric acid. (2 marks)
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261
C2 7.2d – Student worksheet
Electroplating and electrolysis
Electroplating
1
2
Describe how you would electroplate a steel object with silver.
a
What would you use for the electrolyte and why? (2 marks)
b
What would you use for the positive electrode and why? (2 marks)
c
What would you use for the negative electrode and why? (2 marks)
Explain why jewellery might be nickel electroplated with gold, rather than solid gold. (1 mark)
Electrolysis
3
4
Name the ions present in the following:
a
aqueous copper chloride. (2 marks)
b
aqueous potassium nitrate. (2 marks)
c
aqueous sodium sulfate. (2 marks)
d
dilute hydrochloric acid. (2 marks)
Use the reactivity series shown here, and your knowledge
of electrolysis, to predict the products formed at each
electrode during the electrolysis of the following solutions:
a
copper sulfate. (2 marks)
b
potassium iodide. (2 marks)
c
calcium nitrate. (2 marks)
d
sodium bromide. (2 marks)
e
silver nitrate. (2 marks)
f
dilute sulfuric acid. (2 marks)
Extension
5
Balance the following half equations:
a
Cl → Cl2 + e
(1 mark)
b
H + e → H2
(1 mark)
–
+
–
–
2+
+ e → Cu
–
c
Cu
d
Br → Br2 + e
–
–
(1 mark)
(1 mark)
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C2 7.3a – Student worksheet
Electrolysis of lead bromide – demonstration
Your teacher will demonstrate the electrolysis of molten lead bromide.
Answer these questions following the demonstration.
1
Why were the electrodes briefly touched together
at the start? (1 mark)
2
The lead bromide was solid at the start.
3
4
a
What happened to the lamp at this part of the
experiment? (1 mark)
b
Explain why this happened. (2 marks)
The lead bromide was melted during the
experiment.
a
What happened to the lamp when the lead
bromide was molten? (1 mark)
b
Explain why this happened. (1 mark)
2+
Lead bromide contains lead ions, Pb , and
–
bromide ions, Br .
a
Which ions were attracted to the negative
electrode? (1 mark)
b
Did they gain, or lose, electrons at this
electrode? (1 mark)
c
Were these ions reduced or oxidised at this
electrode? (1 mark)
d
Write a half-equation for the reaction that
happens at this electrode. (2 marks)
e
Write a half-equation for the reaction that
happens at the positive electrode. (2 marks)
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263
C2 7.3b – Student worksheet
Aluminium manufacture
3+
2–
Aluminium oxide is insoluble in water. It contains aluminium ions, Al , and oxide ions, O .
Aluminium is manufactured by electrolysing a molten mixture of aluminium oxide and cryolite. The
diagram shows a cross-section of a typical electrolysis cell.
1
Why is aluminium not produced by the electrolysis of an aqueous solution of aluminium oxide?
(1 mark)
2
Why must the mixture of aluminium oxide and cryolite be molten for electrolysis to happen?
(1 mark)
3
Why is cryolite used? (1 mark)
4
During electrolysis, aluminium ions are attracted to the negative electrode, where they gain
electrons.
a
Why are the aluminium ions attracted to the negative electrode? (1 mark)
b
Copy and correctly balance this half-equation:
3+
Al
c
5
+ e → Al
–
(1 mark)
Are the aluminium ions oxidised, or reduced, when they reach the negative electrode?
(1 mark)
The oxide ions are attracted to the positive electrode.
a
Copy and correctly balance this half-equation: O → O2 + e (1 mark)
b
Explain, with the help of a word equation and balanced symbol equation, why the positive
electrodes need replacing from time to time. (2 marks)
2–
–
6
Aluminium ores are abundant in the Earth’s crust, yet aluminium is an expensive metal.
Suggest a reason why aluminium is an expensive metal. (1 mark)
7
Write a balanced symbol equation to show the overall process that happens during the
electrolysis of molten aluminium oxide. (1 mark)
Sheet 1 of 1
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C2 7.3c – Student worksheet
Making aluminium
Read the information about making aluminium, then answer the questions.
Before electrolysis
Sir Humphry Davy discovered potassium and sodium in 1807, when he passed an electric current
through molten potassium hydroxide and molten sodium hydroxide. He later tried to extract
aluminium by passing an electric current through a hot mixture of aluminium oxide and potassium
hydroxide, but could not get this to work. Instead, aluminium was first extracted without electrolysis
by Hans Christian Oersted in 1825.
Oersted passed chlorine over a hot mixture of carbon and aluminium oxide, producing aluminium
chloride and carbon monoxide. He then heated the aluminium chloride with an alloy of potassium
and mercury. Since potassium is more reactive than aluminium, a reaction happened, producing
aluminium and potassium chloride. Unfortunately, the aluminium formed an alloy with the mercury.
The mixture was heated to boil away the mercury, leaving a tiny amount of aluminium. The process
was not very efficient.
Friedrich Wöhler improved upon Oersted’s method. He used potassium instead of a mixture of
potassium and mercury. This worked, but his aluminium was always contaminated with substances
from the crucible. It was not until 1845 that Wöhler managed to make reasonably pure aluminium.
Ten years later, Henri Deville was thinking about how to make commercial amounts of aluminium.
He could have heated aluminium chloride with potassium, or he could have electrolysed aluminium
chloride. Potassium was very expensive and dangerous to use, while at the time batteries were the
only way to make large amounts of electricity. In the end, Deville modified Wöhler’s method using
sodium instead of potassium. The amount of aluminium produced rapidly increased as a result. By
the time the Hall–Héroult process began in the 1890s, the cost of production had fallen by around
99%.
1
The first stage in Oersted’s process needed aluminium chloride, made with the help of chlorine
gas, Cl2.
a
Write a word equation for Oersted’s method of producing aluminium chloride, AlCl3, from
aluminium oxide, Al2O3. (1 mark)
b
Write a balanced symbol equation for the process. (2 marks)
2
Give one reason why Oersted’s process for making aluminium was not very efficient. (1 mark)
3
Wöhler improved upon Oersted’s process.
4
5
a
How did Wöhler improve the process? (1 mark)
b
What problem did he encounter? (1 mark)
c
Write a balanced symbol equation for the reaction between potassium and aluminium
chloride. (2 marks)
Deville improved upon Wöhler’s process.
a
What improvement did Deville make? (1 mark)
b
Explain why he decided upon this improvement. (2 marks)
c
Suggest two reasons why sodium was a better choice than potassium. (2 marks)
Explain how Davy’s research contributed to the eventual isolation of aluminium, even though
he did not manage this himself. (2 marks)
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265
C2 7.4a – Student practical sheet
Electrolysis of sodium chloride solution
The electrolysis of sodium chloride solution produces three substances – hydrogen, chlorine and
sodium hydroxide:
sodium chloride + water → hydrogen + chlorine + sodium hydroxide
2 NaCl (aq) + 2 H2O (l) → H2 (g) + Cl2 (g) + 2 NaOH (aq)
These products have important industrial uses.
Aim
To investigate the electrolysis of sodium chloride solution.
Equipment
●
eye protection
●
Bunsen burner
●
d.c. power supply and low voltage
●
electrolysis cell with carbon electrodes
●
2 × ignition tubes
●
leads with crocodile clips
●
splints
●
teat pipette
●
sodium chloride solution
●
blue litmus paper
●
universal indicator solution
●
pH colour chart
Safety
●
Wear your eye protection.
●
Toxic, harmful and flammable substances will be produced, both in the ignition tubes and in the
final solution.
●
Stop the electrolysis after about 5 minutes.
What you need to do
1
Set up the electrolysis cell.
2
Add enough sodium chloride solution to cover the carbon electrodes but not to overfill it.
3
Fill an ignition tube with sodium chloride solution. Put a finger over the tube and turn the tube
upside down over one of the electrodes. Remove your finger so that the tube stays filled with
sodium chloride solution.
4
Repeat step 3 for the other electrode.
5
Connect the power supply.
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C2 7.4a – Student practical sheet
6
Use universal indicator to measure the pH of the sodium chloride solution. Record your
observations.
7
Identify the positive electrode and the negative electrode. Turn on the d.c. supply for no more
than 5 minutes. Observe what happens at each electrode.
8
Test the gas collected at the positive electrode with blue litmus paper, dampened with water.
Record your observations.
9
Test the gas collected at the negative electrode using a lighted splint. Record your
observations.
10 Measure and record the pH of the solution in the electrolysis cell.
Using the evidence
1
What was the pH of the solution at the start and end of the experiment? What does this result
mean? (2 marks)
2
Identify each gas produced. (2 marks)
Evaluation
3
Describe any problems you may have encountered when identifying the electrolysis products.
(1 mark)
Extension
4
Write a balanced half-equation for the reaction at each electrode. (2 marks)
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267
C2 7.4b – Student worksheet
The chlor–alkali industry
Electrolysis of sodium chloride solution
1
Cut these cards out. Arrange them to show what happens during the electrolysis of sodium
chloride solution.
2
Make notes on the process, based on your arrangement of the cards.
Positive electrode
Negative electrode
Reduction happens
here
Oxidation happens here
Chlorine is produced
Hydrogen is produced
Sodium hydroxide
solution forms
Current passed through
sodium chloride
solution
Cl– ions attracted
H+ ions attracted
Na+ and OH- ions
remain in the solution
Cl– ions discharged in
preference to OH– ions
H+ (aq) + 2 e– →
H2 (g)
2 Cl– (aq) →
Cl2 (g) + 2 e–
Hydrogen is less
reactive than sodium
Membrane prevents
chlorine and sodium
hydroxide mixing
Products from the chlor–alkali industry
3
Cut these cards out. Arrange them to show the uses of the products from the chlor–alkali
industry.
4
Make notes on the products, based on your arrangement of the cards.
Hydrogen
Chlorine
Sodium hydroxide
Used in the manufacture of
polyvinyl chloride, PVC
Used to kill bacteria
in drinking water
Used to make bleach
Used in the manufacture
of soap
Used in the manufacture
of ammonia
Used in the manufacture
of paper
Used in the manufacture
of margarine
Used as a fuel
Used to kill bacteria in
swimming pool water
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C2 7.4c – Student worksheet
Products from salt
The diagram shows a type of cell used in the industrial electrolysis of sodium chloride solution.
1
Name each of the products A, B and C. (3 marks)
2
Write a word equation and a balanced symbol equation for the overall process. (3 marks)
3
Give the formulae of the ions present in the sodium chloride solution. (2 marks)
4
Copy and complete these half-equations.
a
Cl → Cl2 + e (1 mark)
b
H + e → H2 (1 mark)
–
+
–
–
5
Explain why an alkaline solution, solution C, forms. (2 marks)
6
Give one use for each of the three products A, B and C. (3 marks)
7
Suggest a reason for obtaining the sodium chloride by solution mining, rather than by
conventional mining. (1 mark)
8
Describe two hazards presented by the electrolysis of sodium chloride solution, and the
precautions that may need to be taken to protect the workers who look after the electrolysis
cells. (2 marks)
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