Mandatory Experiment 8.1
Simple experiments to illustrate Le Chatelier's Principle
Student Material
(a) The equilibrium between CoCl42- and Co(H2O)62+
Theory
Le Chatelier’s Principle states that when a disturbance is imposed on a system at
equilibrium, the equilibrium shifts in such a way as to minimise the effect of the
disturbance. Some reactions involving cobalt compounds are suitable for illustrating Le
Chatelier’s Principle because they involve clear colour changes.
Co(H2O)62+ is pink in aqueous solution and CoCl42- is blue.
The equilibrium between the two species is
CoCl42- + 6H2O
Blue
Co(H2O)62+ + 4ClPink
The forward reaction is exothermic. The equilibrium between the two species can be
disturbed by (i) adding Cl- ions or water or (ii) changing the temperature. In both cases
the changes that occur are as predicted by Le Chatelier’s Principle.
Chemicals and Apparatus
Cobalt(II) chloride (or cobalt(II) nitrate)
Deionised water
Concentrated hydrochloric acid
Ethanol
Crushed ice
Sodium chloride
Boiling tubes and racks
Dropping pipettes
250 cm3 Pyrex beakers
100 cm3 measuring cylinders
Electronic balance
Safety glasses
1
.
Procedure
NB: Wear your safety glasses.
1. Dissolve 4 g of cobalt chloride-6-water in 40 cm3 of deionised water. The
following equilibrium is set up when the crystals are added to water:
CoCl42- + 6H2O
Blue
Co(H2O)62+ + 4ClPink
Since the pink colour is predominant, we may conclude that the equilibrium lies
on the right hand side.
2. Using a fume cupboard, add concentrated hydrochloric acid, with stirring, until a
violet solution is formed.
3. Adding more concentrated hydrochloric acid produces a blue colour, while adding
water will restore the pink colour. By trial and error produce an “in between”
violet (or lilac) colour which will contain the two cobalt ions. Place this solution
in each of six boiling tubes to a depth of about 2 cm (Fig. 1).
Fig. 1
4. To study the effects of concentration changes on equilibrium, keep one tube as
a control. Add water to a second tube using a dropping pipette. The colour of the
solution should change to pink. The equilibrium has now shifted to the right hand
side as the forward reaction absorbs the stress of the increased concentration of
water.
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5. Using a fume cupboard add concentrated hydrochloric acid to a third tube using a
dropping pipette. The colour of the solution should change to blue. The
equilibrium has now shifted to the left hand side since adding the concentrated
HCl increases the concentration of Cl- ions and, in keeping with Le Chatelier’s
Principle, the concentration of these ions is decreased by the backward reaction
taking place.
.
6. To study the effects of temperature changes on equilibrium, keep one tube as
a control. Place another tube in a beaker of hot water (over 90 0C). Note that the
colour changes to blue. This is in keeping with Le Chatelier's Principle, i.e. the
endothermic reaction (reverse reaction) predominates in order to absorb the added
heat.
7. Place another tube in a beaker of crushed ice and water. Note that the colour
changes to pink. This is in keeping with Le Chatelier's Principle, i.e. the
exothermic reaction (forward reaction) predominates in order to replace the lost
heat.
Questions relating to the experiment
1. Why is a control used in this experiment?
2. Explain why there is a colour change in the mixture when a boiling tube
containing it is placed in ice.
3. Explain why there is a colour change in the mixture when a boiling tube
containing it is placed in hot water.
4. Explain why there is a colour change in the mixture when water is added.
5. Explain why there is a colour change in the mixture when concentrated
hydrochloric acid is added.
6. How can it be shown that it is the chloride ions in the hydrochloric acid that
cause this colour change?
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Teacher Material

It is advisable to have supplies of boiling water and ice available in advance of
the experiment.

Sheets of white paper provide a useful background when the colours of the
solutions are being observed.

The “in between” violet solution described in step 3 of the procedure is crucial to
the success of the experiment. It may therefore be advisable for the teacher to
prepare this solution in sufficient quantity for the class in advance, to ensure that
it has the correct colour. This would also save time; students could be given the
“in between” solution, and asked to continue from there. This approach would
also avoid potential problems caused by each group of students using 60 cm3 of
concentrated hydrochloric acid.

To avoid using too much concentrated hydrochloric acid for this experiment, the
following variation may be found useful. Dissolve the cobalt salt in a small
quantity of water and dilute this solution with ethanol, without precipitating the
salt. Since the solution contains less water, it will be found that less
hydrochloric acid is required.
Extension work

To show that it is the chloride ions in the hydrochloric acid that shift the
equilibrium, add a spatula of sodium chloride to the pink solution. This produces
a bluer colour eventually, although because the salt is slow to dissolve, it will
take some time for the colour change to occur.
Preparation of Reagents
Dilute sodium hydroxide solution (1 M): Carefully add, in stages, 40 g of sodium
hydroxide with constant stirring to about 800 cm3 of water. Continue stirring until all of
the solid has dissolved. Transfer the solution to a volumetric flask. Rinse the beaker and
add rinsings to the volumetric flask. Make the solution up to 1 litre with water. Stopper,
and mix thoroughly.
Dilute ethanoic acid solution (1 M): In a fume cupboard, add 57 cm3 to about 600 cm3
of water in a beaker. Stir, and pour the solution into a 1000 cm3 volumetric flask. Add
rinsings to the volumetric flask. Make the solution up to 1 litre with water. Stopper, and
mix thoroughly.
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Quantities needed per working group
4 g cobalt chloride-6-water
100 cm3 concentrated hydrochloric acid
Safety considerations



Safety glasses must be worn.
The use of gloves is recommended.
Operations involving the use of concentrated hydrochloric acid should be carried
out in the fume cupboard.
Chemical hazard notes
Concentrated hydrochloric acid is very corrosive to eyes and skin, and its vapour is
very irritating to lungs.
Solid cobalt chloride is harmful, and should not be allowed to be ingested or to
come into contact with the skin.
Solid sodium hydroxide is corrosive, and can cause severe burns to eyes and skin.
Always wear eye protection.
Concentrated ethanoic acid can cause severe burns. The vapour is very irritating to
lungs.
Ethanol - highly flammable; keep away from sources of ignition.
Disposal of wastes
Add 1 M sodium hydroxide solution to the waste. Filter off slurry and bag for landfill
waste. Neutralise the filtrate with 1 M ethanoic acid solution and flush to foul water
drain.
Suggested Solutions to Student Questions
1. Why is a control used in this experiment?
In order to be able to compare colours formed with the original colour.
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2. Explain why there is a colour change in the mixture when a boiling tube
containing it is placed in ice.
In the equilibrium
CoCl42- + 6H2O
Co(H2O)62+ + 4ClBlue
Pink
the forward reaction is exothermic. Lowering the temperature favours the
exothermic reaction, according to Le Chatelier’s Principle, and so the
colour changes to pink.
3. Explain why there is a colour change in the mixture when a boiling tube
containing it is placed in hot water.
In the equilibrium
CoCl42- + 6H2O
Co(H2O)62+ + 4ClBlue
Pink
the reverse reaction is endothermic. Raising the temperature favours the
endothermic reaction, according to Le Chatelier’s Principle, and so the
colour changes to blue.
4. Explain why there is a colour change in the mixture when water is added.
Adding water shifts the equilibrium to the right, according to Le Chatelier’s
Principle, and so the colour changes to pink.
5. Explain why there is a colour change in the mixture when concentrated
hydrochloric acid is added.
Adding hydrochloric acid shifts the equilibrium to the left, according to Le
Chatelier’s Principle, and so the colour changes to blue.
6. How can it be shown that it is the chloride ions in the hydrochloric acid that
cause this colour change?
Add solid sodium chloride to the “in-between” solution – a colour change to
blue occurs. Since chloride ions are the only type of ion found in both
hydrochloric acid and sodium chloride, the effect must be due to the chloride
ions.
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Student Material
(b) The equilibrium between CrO42- and Cr2O72Theory
Some reactions involving chromium compounds are also suitable for illustrating Le
Chatelier’s Principle because they involve clear colour changes.
One such equilibrium is Cr2O72- + H2O
orange
 2CrO42- + 2H+
yellow
This experiment will be used to demonstrate the effects of concentration changes on an
equilibrium mixture. Adding an acid will increase the concentration of H+, and adding a
base will reduce it.
Chemicals and Apparatus
Sodium dichromate solution
Sodium hydroxide solution (2 M)
Dilute hydrochloric acid (2 M)
i
Boiling tubes and racks
Dropping pipettes
Safety glasses
Procedure
NB: Wear your safety glasses.
1. Quarter fill a test tube with the solution of sodium dichromate provided. This
should have an orange colour. The following equilibrium exists:
Cr2O72- + H2O
orange
2CrO42- + 2H+
yellow
Since the orange colour predominates, the equilibrium must lie on the left hand
side of the equation.
2. Keep a second sample of the sodium dichromate solution in a test tube as a
control.
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3. Carefully add some bench dilute sodium hydroxide solution until the orange
colour changes to yellow. The sodium hydroxide removes the H+ ions giving rise
to a stress and therefore, in keeping with Le Chatelier's Principle, the forward
reaction predominates to produce more H+ ions.
4. Carefully add dilute hydrochloric acid until the yellow colour changes back to
orange. The added hydrochloric acid creates an excess of H+ ions that causes the
equilibrium reaction to be shifted to the left in order to absorb this excess of H+
ions.
Questions relating to the experiment
1. When sodium hydroxide solution is added to a solution of potassium
dichromate, a colour change occurs. Describe the colour change, and explain why
it happens.
2. Why does adding hydrochloric acid reverse the colour change referred to in
question 1?
3. Why is a control used in this experiment?
4. Describe and explain what happens when hydrochloric acid is added to a solution
of potassium chromate.
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Teacher material

Sheets of white paper provide a useful background when the colours of the
solutions are being observed.
Extension work

This experiment can also be carried out starting with a solution of sodium
chromate, to which first hydrochloric acid solution and then sodium hydroxide
solution are added. In this case a sample of the original sodium chromate
solution should be kept as a control.
Preparation of Reagents
Sodium dichromate solution (0.17 M): Dissolve 20 g of sodium dichromate in
deionised water and make up to 400 cm3. Stopper, and mix thoroughly.
Dilute hydrochloric acid (2 M): Using a fume cupboard, dilute 170 cm3 of concentrated
hydrochloric acid to 1 litre with deionised water. Stopper, and mix thoroughly.
Dilute sodium hydroxide solution (2 M): Carefully add, in stages, 80 g of sodium
hydroxide with constant stirring to about 800 cm3 of water. Continue stirring until all of
the solid has dissolved. Make the solution up to 1 litre. Stopper, and mix thoroughly.
Sodium chromate solution (0.21 M): Dissolve 20 g of sodium chromate in deionised
water and make up to 400 cm3. Stopper, and mix thoroughly.
The 1 M sulfuric acid solution for disposal of the waste from the reaction is prepared as
follows:
(Always dilute sulfuric acid by adding the acid to water and not the other way round.) 56
cm3 of the concentrated acid is added slowly to about 700 cm3 of deionised water
containing about 20 ice cubes. The mixture is stirred and made up to 1 litre in a
volumetric flask with deionised water. The flask is stoppered and inverted a number of
times.
Quantities needed per working group
20 cm3 sodium dichromate solution
30 cm3 dilute sodium hydroxide solution
30 cm3 dilute hydrochloric acid
Safety considerations

Safety glasses must be worn.
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

The use of gloves is strongly recommended, and they should definitely be worn
when the solutions are being made up.
The acid solution, dichromate solution and chromate solution should each be
made up in a fume cupboard.
Chemical hazard notes
Concentrated hydrochloric acid is very corrosive to eyes and skin, and its vapour is
very irritating to lungs.
Solid sodium hydroxide is corrosive, and can cause severe burns to eyes and skin.
Always wear eye protection.
Solid sodium dichromate is toxic, and should not be allowed to be inhaled,
ingested or to come into contact with the skin. The solid salt may cause cancer by
inhalation of dust. Avoid raising dust during preparation of the solution – prepare the
solution in a fume cupboard.
Solid sodium chromate is toxic, and should not be allowed to be inhaled, ingested
or to come into contact with the skin or eyes. Ulceration may occur on damaged skin.
n Sodium metabisulfite is harmful if swallowed. It produces the toxic gas sulphur
dioxide on contact with acids. Using this material in the disposal of waste material from
the experiment should therefore be carried out in a fume cupboard.
Concentrated sulfuric acid is 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.
i Dilute sulfuric acid is harmful to eyes and an irritant to skin.
Disposal of wastes
Add some 1 M sulfuric acid to the waste. In a fume cupboard add sodium metabisulfite to
produce a green solution of chromium(III). Dilute with excess water and flush to foul
water drain.
Suggested Solutions to Student Questions
1. When sodium hydroxide solution is added to a solution of potassium
dichromate, a colour change occurs. Describe the colour change, and explain
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why it happens.
The colour changes from orange to yellow. In the equilibrium
Cr2O72- + H2O
2CrO42- + 2H+
orange
yellow
adding sodium hydroxide solution removes H+, and so shifts the equilibrium
to the right, according to Le Chatelier’s Principle. Therefore the colour
changes to yellow.
2. Why does adding hydrochloric acid reverse the colour change referred to in
question 1?
The colour changes from yellow to orange. In the equilibrium
Cr2O72- + H2O
2CrO42- + 2H+
orange
yellow
adding acid increases the concentration of H+ ions and therefore shifts the
equilibrium to the left, according to Le Chatelier’s Principle. Therefore the
colour changes to orange.
3. Why is a control used in this experiment?
In order to be able to compare colours formed with the original colour.
4. Describe and explain what happens when hydrochloric acid is added to a
solution of potassium chromate.
The colour changes from yellow to orange. In the equilibrium
Cr2O72- + H2O
2CrO42- + 2H+
orange
yellow
adding acid increases the concentration of H+ ions and therefore shifts the
equilibrium to the left, according to Le Chatelier’s Principle. Therefore the
colour changes to orange.
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Student Material
(c) The equilibrium between Fe3+ and Fe(CNS)2+
Theory
Some reactions involving iron compounds are suitable for illustrating Le Chatelier’s
Principle because they also involve clear colour changes.
One such equilibrium is Fe3+ + CNSyellow
Fe(CNS)2+
red
This experiment will be used to demonstrate the effects of concentration changes on an
equilibrium mixture. Adding hydrochloric acid reduces the concentration of Fe3+ by
forming a complex ion containing iron and chlorine. This causes a shift of equilibrium to
the left. The equilibrium can be shifted to the right hand side by adding some potassium
thiocyanate solution.
Chemicals and Apparatus
Concentrated hydrochloric acid
Iron(III) chloride solution (0.05 M)
Potassium thiocyanate solution (0.05 M)
Boiling tubes and racks
Dropping pipettes
Safety glasses
Procedure:
NB: Wear your safety glasses.
1. Mix together about 5 cm3 respectively of solutions of iron(III) chloride and
potassium thiocyanate in a beaker. Note the formation of the red complex.
Fe3+ + CNSyellow
Fe(CNS)2+
red
Since the red complex above is formed, the equilibrium must lie on the right hand
side of the equation.
2. Divide the mixture into three portions in separate boiling tubes (Fig. 2). Keep one
of these as a control.
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Fig. 2
3. Using a fume cupboard, add some concentrated hydrochloric acid to the second
tube until the red colour disappears. In keeping with Le Chatelier's Principle, the
red colour disappears as the equilibrium is shifted to the left hand side to replace
the Fe3+ ions removed.
4. Add an equivalent amount of water to the third tube, and compare. This
comparison should indicate that the extent of lightening of the colour is not due to
a diluting effect.
5. To the second tube, add some potassium thiocyanate solution. The red complex
reforms because the equilibrium is shifted to the right.
Questions relating to the experiment
1. Why is a control used in this experiment?
2. When potassium thiocyanate solution is added to a solution of iron(III)
chloride, a colour change occurs. Describe the colour change, and explain why
it happens.
3. Why does adding hydrochloric acid reverse the colour change referred to in
question 2?
4. How can it be shown that it is the chloride ions in the hydrochloric acid that
cause this reversal?
5. Why is water added in step 4 of the procedure?
6. Name a substance other than hydrochloric acid that can reverse the colour
change referred to in question 2.
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Teacher material

Saturated potassium chloride solution can be used instead of hydrochloric acid in
this experiment, but because it contains much less chloride per unit volume than
hydrochloric acid, it is not as effective. This solution should be shaken
immediately before use.

Sheets of white paper provide a useful background when the colours of the
solutions are being observed.
Preparation of Reagents
Iron(III) chloride solution (0.05 M): Dissolve 13.5 g of iron(III) chloride in about 500
cm3 of deionised water containing about 20 cm3 of concentrated hydrochloric acid.
Make the solution up to 1 litre. Stopper, and mix thoroughly.
Potassium thiocyanate solution (0.05 M): Dissolve 5 g of potassium thiocyanate in
deionised water. Make the solution up to 1 litre. Stopper, and mix thoroughly.
Saturated potassium chloride solution: Add, with stirring, potassium chloride to 200
cm3 of deionised water until some remains undissolved, even after vigorous stirring.
Stopper, and mix thoroughly.
Quantities needed per working group
30 cm3 iron(III) chloride solution
30 cm3 potassium thiocyanate solution
10 cm3 saturated potassium chloride solution
10 cm3 concentrated hydrochloric acid
Safety considerations



Safety glasses must be worn.
The use of gloves is recommended.
Operations involving the use of concentrated hydrochloric acid should be carried
out in the fume cupboard.
Chemical hazard notes
Concentrated hydrochloric acid is very corrosive to eyes and skin, and its vapour
is very irritating to lungs.
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i Iron(III) chloride is an eye and skin irritant, and severe eye burns may result if not
attended to.
n Potassium thiocyanate is harmful if swallowed.
i Sodium carbonate (anhydrous) is an irritant to eyes and skin, and its dust irritates
lungs.
Disposal of wastes
Neutralise wastes with anhydrous sodium carbonate. Dilute with excess water and flush
to foul water drain.
Suggested Solutions to Student Questions
1. Why is a control used in this experiment?
In order to be able to compare colours formed with the original colour.
2. When potassium thiocyanate solution is added to a solution of iron(III)
chloride, a colour change occurs. Describe the colour change, and explain why
it happens.
The colour changes from yellow to red. In the equilibrium
Fe3+ + CNSFe(CNS)2+
yellow
red
some of the red complex Fe(CNS)2+ is formed, giving rise to the red colour.
3. Why does adding hydrochloric acid reverse the colour change referred to in
question 2?
In the equilibrium
Fe3+ + CNSFe(CNS)2+
yellow
red
adding hydrochloric acid causes the removal of Fe3+, due to the formation of a
complex ion containing iron and chlorine. This results in a shift of the
equilibrium to the left, according to Le Chatelier’s Principle, and so the colour
changes to yellow.
4. How can it be shown that it is the chloride ions in the hydrochloric acid that
cause this reversal?
Add saturated potassium chloride solution (or saturated sodium chloride
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solution) to the red solution – a colour change to yellow occurs. Again, this is
due to the removal of Fe3+ because of the formation of a complex ion
containing iron and chlorine. Since chloride ions are the only type of ion
found in both hydrochloric acid and sodium chloride, the effect must be due to
the chloride ions.
5. Why is water added in step 4 of the procedure?
To show by comparison that the extent of lightening of the colour is not due to
a diluting effect.
6. Name a substance other than hydrochloric acid that can reverse the colour
change referred to in question 2.
Potassium chloride, or sodium chloride.
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