Mandatory Experiment 1.2
Investigation of
(a) Redox reactions of the halogens
(b) Displacement reactions of metals
Student Material
(a) Redox reactions of the halogens
Theory
The halogens, fluorine, chlorine, bromine, iodine and astatine are very reactive elements
and are too unstable to exist in nature in an uncombined form. They often react by taking
an electron from another element. This means they react as oxidising agents. The smaller
the halogen atom, the stronger the oxidising agent it is. So in terms of oxidising power
F > Cl > Br > I > At
Fluorine is extremely poisonous, and astatine is unstable and radioactive, so the
investigation here is confined to the other three halogens.
(i) Reactions with halides
Chlorine, being the strongest oxidising agent of the three, is capable of releasing the other
two elements from solutions of their salts:
Cl2 + 2Br- (aq)  2Cl- (aq) + Br2
Cl2 + 2I- (aq)  2Cl- (aq) + I2
Bromine can release iodine from a solution of its salts
Br2 + 2I- (aq)  2Br- (aq) + I2
Your task is to find the evidence to support this theory.
Chemicals and Apparatus
Chlorine solution
i
Bromine solution
Iodine solution
Sodium chloride solution
Sodium bromide solution
Potassium iodide solution
1
Safety glasses
PVC gloves
Fume cupboard or well-ventilated room
Pasteur pipettes
Test tubes
Test tube rack
Test tube brush
Quantities needed per working group
Name of solution
Quantity
Aqueous solutions of chlorine,
bromine and iodine.
Aqueous solutions of chloride, bromide,
and iodide salts.
2 cm3 aliquots (portions) per test
2 cm3 aliquots per test
Procedure
NB: Wear your safety glasses.
1. Copy the following table into your practical report book and fill in your
observations. By reference to this table you will be able to draw conclusions
from your later observations.
Name of solution
Chlorine in water
Bromine in water
Iodine in water
Chloride ions in water
Bromide ions in water
Iodide ions in water
Colour of solution
Table 1
2. Draw a second table into your practical report book with the following
headings:
Solutions added to the
test-tube
(a) Chlorine and bromide
ions
(b) Chlorine and iodide ions
(c) Bromine and iodide ions
Observation
Conclusion
Table 2
2
3. As there are as many as 16 reagent bottles in use in these experiments it is
vital to the success of the experiment not to mix the reagents by putting the
wrong stopper on the wrong bottle. Your teacher will demonstrate to you the
correct way to hold the stopper while using a reagent bottle. Always replace
stoppers immediately after use.
4. For each of the cases (a), (b) and (c) described in Table 2, add 2 cm3 of the
solutions mentioned to separate test tubes and mix.
Record your observations and conclusions. Retain the contents of the test tubes
for comparison purposes, ensuring that the test tubes are correctly labelled.
(ii) Reactions with iron(II) salts and with sulfites
All three halogen solutions are able to oxidise iron(II) ions to iron(III) ions, and to
oxidise sulfite ions to sulfate ions in aqueous solution. For example, chlorine reacts as
follows:
Cl2 + 2Fe2+ (aq)  2Cl- (aq) + 2Fe3+ (aq)
Cl2 + SO32- (aq)  2Cl- (aq) + SO42- (aq)
Your task is to find the evidence to support this theory.
Chemicals and Apparatus
Chlorine solution
Iron(II) sulfate solution
i
n
3
Iron(III) chloride solution
Sodium sulfite solution
i
Sodium hydroxide solution
Silver nitrate solution
Barium chloride solution
n
Dilute hydrochloric acid
Dilute ammonia solution
i
Safety glasses
PVC gloves
Fume cupboard or well-ventilated room
Pasteur pipettes
Test tubes
Test tube rack
Test tube brush
Quantities needed per working group
Name of solution
Quantity
Aqueous solutions of chlorine,
bromine and iodine.
Aqueous solutions of iron(II) sulfate,
iron(III) chloride and sodium sulfite.
Sodium hydroxide solution.
Aqueous solutions of silver nitrate,
barium chloride, dilute hydrochloric acid
and dilute ammonia.
2 cm3 aliquots (portions) per test
2 cm3 aliquots per test
2 cm3
2 cm3 aliquots per test
Procedure
NB: Wear your safety glasses.
1. Copy the following table into your practical report book and fill in your
observations. By reference to this table you will be able to draw conclusions
from your later observations.
Name of solution
Chlorine in water
Iron(II) sulfate in water
Iron(III) chloride in water
Colour of solution
4
Iron(II) ions in sodium hydroxide solution
Iron(III) ions in sodium hydroxide solution
Table 3
2. Draw another table into your practical report book with the following
headings:
Solutions added to the
test-tube
(d) Chlorine and iron(II)
sulfate followed by 10
drops of sodium hydroxide
(e) Chlorine and sodium
sulfite followed by the test
for the presence of chloride
ions or sulfate ions
Observation
Conclusion
Table 4
3. Check that the sulfite solution does not contain any sulfate ions, as follows:
Add 2 cm3 of sodium sulfite solution to a clean test tube. Using a dropping
pipette add a few drops of barium chloride solution. A white precipitate forms.
Now add 2 cm3 of dilute hydrochloric acid. The white precipitate should
dissolve. If any of the white precipitate does not dissolve, then sulfate ions are
present, and so this solution cannot be used.
4. Check that the iron(II) sulfate solution does not contain any iron(III) ions by
adding 10 drops of sodium hydroxide solution and studying the colour of
the precipitate formed and comparing it with the table set up in part 1.
5. For each of the cases (d) and (e) described in Table 4, add 2 cm3 of the
solutions mentioned to separate test tubes and mix.
5
6. Record your observations and conclusions. Retain the contents of the test
tubes for comparison purposes, ensuring that the test tubes are correctly
labelled.
7. Iron(II) ions and iron(III) ions form different coloured floating (flocculent)
precipitates with sodium hydroxide. This allows you to determine whether a
reaction has or has not taken place in (d) above.
8. In (e) above, the test for sulfate ions is carried out as described in 3 above.
The test for chloride ions is carried out as follows: Add the solution to be
tested to a clean test tube. Using a dropping pipette add a few drops of silver
nitrate solution. If chloride is present, a white precipitate forms. Now add 2
cm3 of dilute ammonia solution. The white precipitate should dissolve.
6
(b) Displacement reactions of metals
Theory
In this experiment, magnesium and zinc respectively are reacted with a solution of
copper(II) sulfate.
Metals higher up on the electrochemical series displace metals lower down from aqueous
solutions of their salts. The metal higher up in the series is oxidised in the process and
forms a soluble positive ion. The metal lower down in the series is reduced, gains
electrons and becomes a solid metallic element.
Use of this principle will allow, for example, the use of scrap iron (higher up in the
series) to liberate copper (lower down in the series) from an aqueous solution of
copper(II) sulfate.
Fe(s) + CuSO4(aq)  FeSO4(aq) + Cu(s)
Equations
Zn(s) + CuSO4(aq)  ZnSO4(aq) + Cu(s)
Mg(s) + CuSO4(aq)  MgSO4(aq) + Cu(s)
This experiment works best under acidic conditions; under these conditions, the
following reactions take place simultaneously:
Zn(s) + H2SO4(aq)  ZnSO4(aq) + H2 (g)
Mg(s) + H2SO4(aq)  MgSO4(aq) + H2 (g)
Your task is to find the evidence to support this theory.
Chemicals and Apparatus
Acidified copper(II) sulfate solution
Zinc powder
Magnesium ribbon
Safety glasses
Boiling tubes
Boiling tube rack
Test tube brush
Pasteur pipettes
7
Quantities needed per working group
2 cm3 aliquots per test
1.0 g
0.5 g
Acidified copper(II) sulfate solution
Zinc powder
Magnesium ribbon (cleaned with sandpaper)
Procedure
NB: Wear your safety glasses.
1. Copy the following table into your practical report book
Mg
Zn
(a) Colour of
CuSO4(aq) at the
beginning
(b) Colour of the
solution at the end of
the reaction
(c) Colour of the
precipitate formed
(d) Any other
observations
Table 5
2. Draw another table into your practical report book with the following
headings:
Mg
Zn
Observation
(a)
(b)
(c)
(d)
(a)
(b)
(c)
(d)
Conclusion
Table 6
8
3. Half fill two boiling tubes with the acidified copper(II) sulfate solution.
4. Add the magnesium ribbon to the solution in one boiling tube and the zinc
powder to the solution in the other boiling tube. Record your observations
and conclusions.
Questions relating to the experiment
1. Write out a full list of equations for all the redox reactions that took
place during this experiment. Using oxidation numbers, label the
species that have been oxidised and reduced, e.g.
Cl2 + 2Br-(aq)  2Cl-(aq) + Br2
0
-1
-1
0
r
o
2. The iodine/thiosulfate titration to be dealt with in experiment 4.7 is
another redox reaction involving a halogen. The equation for the
reaction is
2Na2S2O3 + I2  2NaI + Na2S4O6
Show that the iodine is acting as an oxidising agent there as well.
3. Describe the tests you would use to distinguish between sulfite and
sulfate anions in aqueous solution.
4. Explain why it is difficult to make an aqueous solution of iodine. What
particular method is used to overcome this problem?
5. An aqueous solution of chlorine is often made by reacting
concentrated hydrochloric acid with a diluted commercial bleach
9
solution. The active ingredient in commercial bleach is sodium
hypochlorite (NaOCl). The equation for the reaction is
NaOCl + 2HCl  Cl2 + NaCl + H2O
Show what species are oxidised and reduced during this reaction.
6. The position of zinc in the Periodic Table would allow you to predict
the colour of its sulfate salt solution. Explain.
7. Explain why magnesium is more reactive than zinc.
8. What would you expect to see happen if a piece of copper wire was
suspended in a solution of silver nitrate? (Silver nitrate is very
expensive but your teacher may be able to demonstrate this
experiment.)
9. Carry out some research to find out why commercial photography
laboratories might have a special interest in these kinds of reactions.
10
Teacher Material

This experiment needs quite an amount of preparation time. Additional value can
be got from it by making up the stable solutions during previous single periods,
thereby familiarising your students with the appearance of the solutions and with
the precautions needed when working with them.

Students generally find this material difficult. It can be useful to repeat some of
the reactions as demonstrations at a later stage when reviewing the laboratory
notebooks.

Silver nitrate solution is light-sensitive and should be stored in a dark container,
away from sunlight.

The following solutions must be prepared fresh: aqueous solution of chlorine,
aqueous iron(II) sulfate solution and aqueous sodium sulfite solution.

In the preparation of chlorine solution from sodium hypochlorite solution and
hydrochloric acid, the reaction is as follows:
2HCl + OCl- → Cl2 + H2O + Cl-

The full balanced equation for the reaction of chlorine with sulfite ions is as
follows:
Cl2 + SO32- + H2O  2Cl- + SO42- + 2H+

Clean the magnesium ribbon using a small piece of sandpaper before weighing it.
Measure the length of the weighed piece and use that approximate length for
everyone.

If the displacement of metals experiments are not carried out using acidic
conditions, a black precipitate rather than a copper-coloured precipitate is likely to
be formed.
Preparation of reagents
Distilled or deionised water should be used in making up all these solutions.
Aqueous solution of chlorine: This solution must be freshly prepared. Take 100 cm3
of commercial bleach. Add it to 500 cm3 of water. In a fume-cupboard, add
concentrated hydrochloric acid drop by drop with constant stirring until a drop of the
solution is just acid to litmus.
Put 50 cm3 of the solution into each of 12 reagent bottles.
Aqueous solution of bromine: This is best purchased as a solution because liquid
bromine itself is a poisonous liquid and difficult to work with. The solution purchased
11
may be diluted before use so as to match the concentrations of the other reagents.
Bromine water does deteriorate with time but lasts much longer than chlorine water,
provided that it is stored in a brown bottle.
In the event that the solution has to be made up directly using bromine, in a fumecupboard, shake 0.5 cm3 bromine with 100 cm3 water. Store in a tightly stoppered bottle.
Aqueous solution of iodine: Iodine crystals dissolve very poorly in water but dissolve
readily in an aqueous solution of potassium iodide forming KI3(aq). The I3- ion releases I2
molecules in reactions and so is always treated as an aqueous solution of iodine. Dissolve
20 g of potassium iodide in 500cm3 of water. Add about 10 g of iodine crystals, dissolve
and make up to 1 litre. This solution is quite stable.
Aqueous solution of sodium chloride (approximately 0.2 mol l-1): Dissolve about 12 g
in water and make up to a litre.
Aqueous solution of sodium bromide (approximately 0.2 mol l-1): Dissolve about 20 g
in water and make up to a litre.
Aqueous solution of potassium iodide (approximately 0.2 mol l-1): Dissolve about 33 g
in water and make up to a litre.
Aqueous sodium hydroxide solution: Dissolve 10 g of pellets in 200cm3 of cold water
and make up to 250 cm3 to make an approximately 1 mol l-1 solution. Special care is
required as this is a caustic solution.
Aqueous iron(II) sulfate solution: This solution must be freshly prepared. Dissolve
11.2 g of the crystalline salt in 100 cm3 of water containing 2 cm3 of concentrated
sulfuric acid, and dilute to 200 cm3 to make an approximately 0.2 M solution.
Aqueous iron(III) chloride solution: Dissolve 11 g of the hydrated salt in 100 cm3 of
water containing 4 cm3 of concentrated hydrochloric acid, and dilute to 200 cm3 to make
an approximately 0.2 M solution.
.
Aqueous sodium sulfite solution: This solution must be freshly prepared. Dissolve
5.2 g of the salt in 100 cm3 of water, and dilute to 200 cm3 to make an approximately 0.2
M solution.
Aqueous ammonia solution: In a fume-cupboard, dilute 40 cm3 of concentrated
ammonia solution to 250 cm3 to make an approximately 3 M solution.
Silver nitrate solution: Dissolve 4 g of the crystals and make up to 250 cm3. This makes
an approximately 0.1 mol dm-3 solution. The solution must be made with deionised water
and stored in a brown bottle. Light tends to reduce silver ions, and halide ions in tap
water would form a precipitate with the silver ions. This solution is also best freshly
prepared and in very small quantities.
12
Aqueous solution of barium chloride (approximately 0.2 mol l-1):
Dissolve about 40 g in water and make up to 1 litre.
Dilute aqueous solution of hydrochloric acid (approximately 2 mol l-1):
In a fume cupboard, add about 170 cm3 of concentrated hydrochloric acid slowly with
stirring to about 500 cm3 of water and make up to 1 litre.
Acidified copper(II) sulfate solution: Dissolve 5 g of copper(II) sulfate pentahydrate
(CuSO4.5H2O) in about 100 cm3 water and make up to 200 cm3. This makes an
approximately 0.1 mol l-1 solution. Carefully add 20 cm3 of concentrated sulfuric acid.
Quantities needed per working group
100 cm3 of each of the solutions should be placed in 125 cm3 reagent bottles.
Name of solution
Quantity
Aqueous solutions of chlorine,
bromine and iodine.
Aqueous solutions of chloride, bromide,
and iodide salts.
Aqueous solutions of iron(II) sulfate,
iron(III) chloride and sodium sulfite.
Aqueous starch solution.
Sodium hydroxide solution.
Aqueous solutions of silver nitrate,
barium chloride, dilute hydrochloric acid
and dilute ammonia.
Aqueous solution of copper(II) sulfate.
Zinc powder
Magnesium ribbon (cleaned with sandpaper)
2 cm3 aliquots (portions) per test
2 cm3 aliquots per test
2 cm3 aliquots per test
A few drops
A few drops
2 cm3 aliquots per test
2 cm3 aliquots per test
1.0 g
0.5 g
Safety Considerations

The chlorine and bromine solutions and their vapours are poisonous.
Consequently, a fume cupboard with proper ventilation should ideally be used for
some parts of this experiment.
Chemical hazard notes
Aqueous solution of chlorine
i:
Vapour attacks lungs, eyes and nose.
13
Bromine
: The vapour is highly toxic by inhalation. The liquid causes severe
burns to eyes and skin. The aqueous solution attacks lungs, eyes and nose.
Iodine
n:
Harmful by skin contact and by inhalation. Avoid eye contact.
Silver nitrate
: Solutions are very dangerous to the eyes and blacken skin.
Sodium hydroxide
protection.
: Caustic, harmful to skin and especially to eyes. Always wear eye
Ammonia solution
: Pungent vapour toxic by inhalation; irritating to eyes and
respiratory system; in case of contact with eyes wash immediately with plenty of water
and seek medical advice.
Barium chloride
n:
Harmful by ingestion or inhalation of dust.
Concentrated hydrochloric acid
very irritating to lungs.
: Very corrosive to eyes and skin, and its vapour is
Concentrated sulfuric acid
: Very corrosive to eyes and skin. Due to its very
considerable heat of reaction with water, it is essential that the acid be added to water
when it is being diluted.
Iron(II) sulfate
n:
Harmful if swallowed. Irritating to eyes and skin.
Iron(III) chloride
unattended.
Copper(II) sulfate
i: Eye
n:
and skin irritant. Severe eye burns may result if left
Skin and eye irritant. Harmful if ingested.
Magnesium
: Flammable; burns with an intense flame. Poisonous by ingestion.
Inhalation of dust harmful.
Zinc
: Zinc dust at the bottom of the container could be flammable.
Disposal of wastes
Solid products of the reactions such as barium sulfate and copper should be filtered,
mixed with sand and placed in a refuse bin. Except in the cases that follow, residual
liquid waste should be diluted with excess water and flushed to the foul water drain.
Waste containing bromine, as a result of the reaction between chlorine and bromide ions,
should be treated with 10% sodium carbonate solution before diluting with excess water.
Waste containing iodine, as a result of the reaction between bromine and iodide ions,
should be treated with 25% sodium thiosulfate solution before diluting with excess water.
Residual liquid waste from the replacement reactions of metals experiment should be
neutralised with sodium carbonate before diluting with excess water.
14
Specimen results (a)
Name of solution
Chlorine in water
Bromine in water
Iodine in water
Chloride ions in water
Bromide ions in water
Iodide ions in water
Iron(II) sulfate in water
Iron(III) chloride in water
Iron(II) ions in sodium hydroxide solution
Iron(III) ions in sodium hydroxide solution
Colour of solution
Pale green
Yellow/orange
Brown/red
Colourless
Colourless
Colourless
Pale green
Yellow
Green muddy precipitate
Brown precipitate
Conclusion
Solutions added to the
test-tube
(a) Chlorine and bromide
ions
(b) Chlorine and iodide ions
Observation
(c) Bromine and iodide ions
Darkening of solution
(d) Chlorine and iron(II)
sulfate followed by a few
drops of sodium hydroxide
(e) Chlorine and sodium
sulfite followed by the test
for the presence of sulfate
ion
or the test for the presence
of chloride ions
Orange/brown muddy
precipitate
Iodide ions oxidised to
iodine
Iron(II) ions oxidised to
iron(III) ions
Sulfate test will show a
permanent white precipitate
Sulfite ions oxidised to
sulfate ions
Orange/yellow solution
Bromide ions oxidised to
formed
iodide ions oxidised to iodine
bromine
Brown/red solution formed Iodide ions oxidised to
iodine
iron(II) ions oxidised to iron(III) ions
a white precipitate which
Chlorine has been reduced
dissolves with the addition
to chloride ions
of dilute aqueous ammonia
15
Specimen results (b)
(a) Colour of
CuSO4(aq) at the
beginning
(b) Colour of the
solution at the end of
the reaction
(c) Colour of the
precipitate formed
(d) Any other
observations
Mg
Zn
Mg
Blue
Zn
Blue
Colourless
Colourless
Brown
Brown
Gas
evolved
Gas
evolved
Observation
Blue solution
Colourless solution
Brown precipitate
Gas evolved
Blue solution
Colourless solution
Brown precipitate
Gas evolved
Conclusion
Cu2+(aq) present
Mg2+(aq) present
Copper metal powder
Hydrogen produced
Cu2+(aq) present
Zn2+(aq) present
Copper metal powder
Hydrogen produced
Solutions to student questions
1. Write out a full list of equations for all the redox reactions that took place
during this experiment. Using oxidation numbers, label the species that
have been oxidised and reduced, e.g.
Cl2 + 2Br-(aq)  2Cl-(aq) + Br2
0
-1
-1
0
r
o
The following are the equations to be labelled by the student:
Cl2 + 2Br-(aq)  2Cl-(aq) + Br2
0
-1
-1
0
r
o
Cl2 + 2I-(aq)  2Cl-(aq) + I2
0
-1
-1
0
r
o
16
Br2 + 2I-(aq)  2Br-(aq) + I2
0
-1
-1
0
r
o
Cl2 + 2Fe2+(aq)  2Cl-(aq) + 2Fe3+(aq)
0
+2
-1
+3
r
o
Cl2 + SO32-(aq)  2Cl-(aq) + SO42-(aq)
0 +4-2
-1
+6-2
r
o
Zn(s) + CuSO4(aq)  ZnSO4(aq) + Cu(s)
0
+2
+2
0
o
r
Mg(s) + CuSO4(aq)  MgSO4(aq) + Cu(s)
0
+2
+2
0
o
r
Mg(s) + H2SO4(aq)  H2(g) + Mg SO4 (aq)
0
+1
0
+2
o
r
Zn(s) + H2SO4(aq)  H2(g) + ZnSO4(aq)
0
+1
0
+2
o
r
2. The iodine/thiosulfate titration to be dealt with in experiment 4.7 is
another redox reaction involving a halogen. The equation for the reaction
is
2Na2S2O3 + I2  2NaI + Na2S4O6
Show that the iodine is acting as an oxidising agent there as well.
2Na2S2O3 + I2  2NaI + Na2S4O6
+1 +2–2 0
+1-1 +1+2.5-2
o
r
The iodine is reduced, and is therefore an oxidising agent.
17
3. Describe the tests you would use to distinguish between sulfite and sulfate
anions in aqueous solution.
Add 2 cm3 of the solution to be tested to a clean test tube. Using a dropping
pipette add a few drops of barium chloride solution. A white precipitate
forms. Now add 2 cm3 of dilute hydrochloric acid. The white precipitate
will dissolve if a sulfite is present, and will not dissolve if a sulfate is
present.
4. Explain why it is difficult to make an aqueous solution of iodine. What
particular method is used to overcome this problem?
Iodine crystals dissolve very poorly in water, because iodine is non-polar.
Iodine dissolves readily in an aqueous solution of potassium iodide,
because it reacts to form I3- ions. A solution of I3- ions is always treated as
an aqueous solution of iodine, as I3- ions release I2 molecules in reactions.
5. An aqueous solution of chlorine is often made by reacting concentrated
hydrochloric acid with a diluted commercial bleach solution. The active
ingredient in commercial bleach is sodium hypochlorite (NaOCl). The
equation for the reaction is
NaOCl + 2HCl  Cl2 + NaCl + H2O
Show what species are oxidised and reduced during this reaction.
NaOCl + 2HCl  Cl2 + NaCl + H2O
+1-2+1 +1-1
0
+1 -1 +1-2
r
o
6. The position of zinc in the Periodic Table would allow you to predict the
colour of its sulfate salt solution. Explain.
Zinc is a d-block metal but it is not a transition metal. It therefore will not
be expected to form coloured compounds.
7. Explain why magnesium is more reactive than zinc.
Magnesium atom has a larger atomic radius and a smaller nuclear charge
than a zinc atom. Because of this, and despite the fact that there is extra
screening of outer electrons by electrons in inner energy levels in a zinc
atom compared to a magnesium atom, the outer electrons in a magnesium
atom are not as tightly bound as those in an atom of zinc. Magnesium is
therefore higher on the electrochemical series and more reactive than zinc.
8. What would you expect to see happen if a piece of copper wire was
suspended in a solution of silver nitrate? (Silver nitrate is very expensive
but your teacher may be able to demonstrate this experiment.)
18
Crystals of silver should appear on the surface of the copper wire. The
solution should gradually take on a blue colour (Cu+2(aq)).
9. Carry out some research to find out why commercial photography
laboratories might have a special interest in these kinds of reactions.
Silver halides are reduced to silver when photographic film or paper is
exposed to light. More silver halide is reduced when the film or paper is
being developed. Unreacted silver halide is dissolved away near the end of
the processing. The silver should be recovered as it is a valuable metal and
would damage the environment as a waste chemical. One possible way is
to reduce the metal halide by reaction with a metal higher up on the
electrochemical series.
19