Mill Hill County High School
Chemistry Department
Year 12
2006 – 2007
Chemistry Practical Book II
Name:_____________________________
Form:___________
Teachers: __________________________
This booklet contains 4 practical sessions and 5 homework practical exercises.
The results and calculations of each practical must be completed at each session.
1
YEAR 12 PRACTICAL 4
Determining the percentage purity of a sample of calcium carbonate
In this experiment, you will dissolve a piece of impure calcium carbonate in excess
hydrochloric acid. The calcium carbonate reacts with the acid as follows:
CaCO3(s) + 2HCl(aq)  CaCl2(aq) + CO2 + H2O(l)
You will then titrate samples of 1.0 moldm-3 NaOH against the excess hydrochloric acid.
Phase 1:
1.
Dissolving the calcium carbonate in excess hydrochloric acid
Weigh out accurately 3 – 4 g of impure calcium carbonate, and transfer it
carefully into a 250 cm3 volumetric flask.
2.
Carefully add 1.0 moldm-3 HCl to the volumetric flask until the calcium carbonate
dissolves.
3.
Then make the solution up to the 250 cm3 mark with 1.0 moldm-3 HCl and shake
vigorously to ensure the contents are evenly mixed.
Phase 2:
Titrating the sodium hydroxide against the acid
1.
Pipette 25.0 cm3 of 0.5 moldm-3 sodium hydroxide into a conical flask.
2.
Add a few drops of phenolphthalein indicator.
3.
Rinse the burette in the acid solution from the volumetric flask, and then fill the
burette with the acid solution.
4.
Titrate the alkali sample against the acid solution.
5.
Repeat until 2 concordant values are obtained.
6.
Record your titration results in a table.
1
Phase 3:
Determining the percentage purity of the calcium carbonate
1.
Determine the number of moles of acid used in the titration.
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Hence determine the number of moles of acid present in the 250 cm3 volumetric
flask, after the reaction with calcium carbonate.
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2.
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3.
Calculate the moles of hydrochloric acid originally added to the volumetric flask,
and hence the moles of hydrochloric acid used up in the reaction with calcium
carbonate.
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4.
Hence determine the moles of calcium carbonate present, and the mass of
calcium carbonate present.
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5.
Compare the mass of calcium carbonate with the mass of impure calcium
carbonate weighed out at the start of the experiment, and hence calculate the
percentage purity of the calcium carbonate.
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2
YEAR 12 PRACTICAL
HOMEWORK EXERCISE 4
DETERMINING THE CONCENTRATION OF A SOLUTION OF LACTIC ACID
PLANNING
/8
3
Introduction
Lactic acid is a monoprotic acid which reacts with sodium hydroxide. Using HA to
represent lactic acid, the equation for neutralisation is
HA + NaOH

NaA + H2O.
Phenolphthalein is a suitable indicator for the reaction.
Question
You are provided with a solution of lactic acid, the concentration of which is
approximately 0.05 moldm-3. Using the information given above, describe how you
would determine the exact concentration of the acid. Assume that you have access to
appropriate standard solution of sodium hydroxide.
Your answer must include
1.
The choice of a suitable concentration of sodium hydroxide solution used and
your reasons for choosing this concentration.
2.
A detailed description of the experiments you would perform.
3.
A clear explanation of how you would use your results to calculate the
concentration of the lactic acid solution.
4.
Details of the potential hazards, and the relevant safety precautions you would
take.
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6
YEAR 12 PRACTICAL 5 - USING GAS SYRINGES
Aim: Showing that one mole of magnesium produces one mole of hydrogen
according to the equation:
Mg(s) + 2HCl(aq)  MgCl2(aq) + H2(g)
Planning
1.
What is the maximum volume of hydrogen that it is safe to produce during this
experiment?
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2.
What mass of magnesium is needed to produce this volume?
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What volume of 1.0 moldm-3 hydrochloric acid is needed to react with this
magnesium?
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3.
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Method
4.
Weigh out a quantity of magnesium that is likely to give a sensible amount of
hydrogen gas and place it into a conical flask.
5.
Using a measuring cylinder, measure out enough acid to ensure it is in excess.
6.
Set up the gas cylinder as follows (ensure it is well-clamped):
7.
Add the acid to the magnesium and replace the bung as quickly as possible.
8.
When all the magnesium has reacted, note the volume of hydrogen produced.
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9.
Measure the temperature of the solution. ________________________
Hence deduce the number of moles of acid hydrogen produced.
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7
Analysis
10.
Hence deduce the number of moles of hydrogen produced per mole of
magnesium.
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11.
Deduce the percentage error in your answer.
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12.
Deduce the apparatus error from the following data:
Mass balance: ± 0.01 g
Gas syringe: ±1 cm3
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13.
Comment on your answers to 11 and 12.
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14.
Why is the apparatus error in the measuring cylinder not included in your error
estimation?
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8
15.
Suggest a way in which the accuracy of this experiment could be improved.
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16.
Outline an experiment that could be used to show that one mole of calcium
carbonate gives one mole of carbon dioxide, according to the equation:
CaCO3(s) + 2HCl(aq)  CaCl2(aq) + CO2 + H2O(l)
Indicate in your answer the mass of calcium carbonate you would use, and the
volume and concentration of acid you would use.
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9
YEAR 12 PRACTICAL
HOMEWORK EXERCISE 5
DETERMINATION OF THE RELATIVE MOLECULAR MASS OF A GROUP I METAL
CARBONATE
PLANNING
/8
10
Introduction
A Group I metal carbonate, M2CO3, (relative formula mass M, between 100 and 150)
reacts with hydrochloric acid as shown in the equation below:
M2CO3(aq) + 2HCl(aq)  2MCl(aq) + CO2(g) + H2O(l)
The M can be calculated by using the results from a titration in which standard
hydrochloric acid is added from a burette to 25.0 cm3 portions of a solution of the
carbonate. Methyl orange is a suitable indicator for this titration; it is red in acid and
yellow in alkali.
You are provided with a solid sample of the Group I metal carbonate and a 0.100
moldm-3 solution of hydrochloric acid.
Questions
Use the information above to answer the following questions in the space provided.
1
Suggest a suitable concentration for the Group 1 metal carbonate solution to be
used. Explain your reasons for choosing this concentration.
2
State the volume of the standard solution of the metal carbonate you would
prepare. Calculate the mass of the Group 1 metal carbonate you would need to
weigh out to prepare this solution.
3
Describe in detail how you would prepare your standard solution. You do not
need to describe the titration itself.
4
State how much indicator should be used and describe the colour change at the
end-point.
5
Show how you would use the titration results to calculate the M of the Group 1
metal carbonate.
6
State one potential hazard and the relevant safety precautions you would take.
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YEAR 12 PRACTICAL 6 – GRAVIMETRIC ANALYSIS
Determining the water of crystallisation of copper (II) sulphate
Copper sulphate exists in hydrated form, as CuSO4.xH2O. This form of copper (II)
sulphate is blue. On heating, the hydrated copper sulphate loses its water to become
anhydrous copper (II) sulphate. This process can be represented by the equation:
CuSO4.xH2O(s)  CuSO4(s) + xH2O(g)
By weighing the copper sulphate sample before and after strong heating it is possible to
calculate the value of x.
Method
1.
Weigh out accurately between 5 and 6 grams of hydrated copper sulphate, using
a 2 decimal place balance. Weigh a crucible without any solid in it, add the
correct amount of solid, and weigh the crucible again.
2.
Place the crucible into a clay pipe triangle and heat with a roaring flame for five
minutes. Then allow it to cool. When it is cool enough to touch weigh it again.
3.
Place the crucible back in the gauze and heat strongly for a further three
minutes. Then allow it to cool. When it is cool enough to touch weigh it again.
4.
Repeat the process, heating for two minutes each time, until there is no further
change in mass.
NOTES
14
Analysis
First method:
1.
Calculate the number of moles of anhydrous copper sulphate produced during
the reaction. Hence deduce the number of moles of hydrated copper sulphate
that was initially present.
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2.
Calculate the relative formula mass of hydrated copper sulphate and hence the
value of x.
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Second method:
1.
Calculate the mass of water lost during the experiment.
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2.
Calculate the moles of water lost and the moles of anhydrous copper sulphate
produced.
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3.
Calculate the ratio of the moles of water to the moles of copper sulphate. This is
the value of x.
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15
Evaluation
1.
Deduce the overall percentage apparatus error, given that the percentage error
in the mass balance is ±0.01 g and two critical measurements were made.
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2.
Using the first analytical method, work out the difference between your relative
formula mass and the correct formula mass based on x = 5.
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3.
Hence calculate the percentage error in the experiment, and compare this to you
the apparatus error.
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4.
Suggest two possible improvements, and explain why they would help.
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16
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YEAR 12 PRACTICAL
HOMEWORK EXERCISE 6
DETERMINING THE MOLE RATIO OF THE REACTION BETWEEN STRONTIUM
AND HYDROCHLORIC ACID
PLANNING
/8
17
Introduction
It is thought that strontium metal and hydrochloric acid reacts as follows:
Sr (s) + 2HCl (aq)

SrCl2 (aq) + H2 (g)
It is very difficult to weigh accurately the hydrogen gas produced in this reaction.
However, its volume can be measured by collecting the gas over water or in a gas
syringe.
Both strontium and hydrogen and hydrogen are flammable.
Question
You are provided with strontium and dilute hydrochloric acid.
Describe how, by
experiment you would confirm that one mole of strontium produces one mole of
hydrogen as shown in the equation above.
Your answer must include
1.
A suitable volume of hydrogen to be collected.
2.
A suitable mass of strontium to be used. Assume that under these conditions of
this experiment the volume occupies one mole of hydrogen is 24 dm3
3.
Details of the apparatus you would use, detailed description of the
measurements you would make and a diagram were appropriate to illustrate your
description.
4.
A brief explanation of how you would use your results to confirm the number of
moles of hydrogen produced for each mole of strontium reacted.
5.
Details of the potential hazards, and relevant safety precautions.
18
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20
YEAR 12 PRACTICAL 7 – GRAVIMETRIC ANALYSIS
Determining the empirical formula of magnesium oxide
Magnesium reacts with oxygen to produce magnesium oxide. By reacting a known
mass of magnesium with oxygen and weighing the mass of magnesium oxide produced,
the empirical formula of the oxide can be deduced.
Method
1.
Weigh a crucible with its lid on a 2dp balance.
2.
Add a small piece of magnesium ribbon to the crucible and weigh it again.
2.
Place the crucible into a clay pipe triangle and heat with a roaring flame with the
lid off. As soon as the magnesium starts to react, place the lid on the crucible to
prevent smoke escaping. After the reaction stops, heat for one more minute to
ensure the reaction is complete.
3.
Allow the crucible to cool, and then weigh it again.
Note: you must cool the hot crucible (on a heat proof mat) before weighing on the cold
surface of the balance.
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21
Analysis
1.
Deduce the mass of magnesium used in the experiment.
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2.
Deduce the mass of magnesium oxide produced, and hence the mass of oxygen
added.
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3.
Deduce the moles of magnesium and oxygen present in the oxide.
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4.
Hence deduce the empirical formula of the oxide.
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22
Evaluation
1.
Deduce the overall percentage apparatus error, given that the percentage error
in the mass balance is ±0.01 g and two critical measurements were made.
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2.
Using the first analytical method, work out the difference between the moles of
magnesium and the moles of oxygen, and express as a percentage of the moles
of magnesium.
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3.
Compare this to you the apparatus error.
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23
YEAR 12 PRACTICAL
HOMEWORK EXERCISE 7
DETERMINATION OF THE NUMBER OF MOLES OF WATER OF
CRYSTALLISATION
ANALYSING AND EVALUATING
/14
24
DETERMINATION OF THE NUMBER OF MOLECULES OF WATER OF
CRYSTALLISATION IN HYDRATED CALCIUM SULPHATE CRYSTALS
Introduction
The water of crystallisation in calcium sulphate crystals can be removed as water
vapour by heating as shown in the following erquation
CaSO4.xH2O (s)
Hydrated

CaSO4 (s)
+ xH2O (g)
Anhydrous
A student weighed a clean dry crucible. The student weighed 0.250 g of hydrated
copper sulphate to the crucible. The crucible was then heated. When the crucible and
its content had reached a constant mass, the mass was recorded. The mass was
repeated using different masses of hydrated calcium sulphate.
For each experiment the student recorded the original mass of hydrated calcium
sulphate and the mass of anhydrous calcium sulphate left after heating. The student’s
results are shown on the graph below.
25
26
Analysis.
1.
Draw a best fit straight line on the graph.
2.
Use the graph to determine the mass of hydrated calcium sulphate, which would
have formed 1.000 g of anhydrous calcium sulphate
Mass of hydrated calcium sulphate………………………………………
3.
Calculate the number of moles of CaSO4 present in 1.000 g of anhydrous
calcium sulphate.
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4.
Use your answer to part 2 and part 3 to calculate the M r of hydrated calcium
sulphate.
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5.
Use your answer to part 4 to calculate the value of x in CaSO4.xH2O.
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6. Assume that the maximum error in weighing 0.250 g on a balance was 0.100 g.
Calculate the percentage error in using the balance.
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Evaluation
1.
Consider your graph and comment on the results obtained by the student. Is
your line of best fit good enough for you to use with confidence? Identify any
anomalous results.
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2. Explain why it is necessary for the student to heat the crucible to constant mass.
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3.
Pure hydrated calcium sulphate has the formula CaSO 4.xH2O. Calculate the
difference between the Mr determined in part 4 of the analysis section and the Mr
of the literature value. Express this as a percentage of the Mr of the literature
value. (If you could not complete part 4 of the analysis section, you should assume that
the mass is 165.2. This is not the correct answer)
Difference ……………………………………………………………………
Percentage……………………………………………………………………
4.
Suggest one reason in each case why
(a)
small amounts of hydrated calcium sulphate, such as 0.100g should not
be used in this experiment.
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(b)
Large amounts of hydrated calcium sulphate, such as 50g, should not be
used in this experiment.
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28
YEAR 12 PRACTICAL
HOMEWORK EXERCISE 8
DETERMINING THE NUMBER OF MOLES OF WATER OF CRYSTALLISATION OF
BARIUM BROMIDE
PLANNING
/8
29
Introduction:
Barium chloride crystals contain water of crystallisation. Heating the crystals drives off
the water
BaCl2.2H2O (s)

BaCl2 (s)
+ 2H2O (g)
Hydrated barium chloride losses approximately 15% of its mass when fully dehydrated.
Barium compounds are toxic.
Questions
Using the information above, describe how you would confirm that there are two moles
of water of crystallisation in one mole of hydrated barium chloride. Assume that you
have access to a two decimal place balance
Your answer must include:
1. The choice of a suitable mass of barium chloride to be used and your reasons for
choosing this mass.
2. A description of the experiments you would perform, including details of the
apparatus you would use.
3. A clear explanation of how you would use the results to confirm that the formula
is BaCl2.2H2O.
4. Details of the potential hazards and the relevant safety precautions.
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