CHM1313practical Ans

CHM114/114A LABORATORY MANUAL Page 1 of 83

A. SAFETY 1. LABORATORY SAFETY RULES The science laboratory is one of the safest places to work if one follows the general safety rules prescribed below. 1. Do not eat or drink in the laboratory.

2. Do not taste any chemical.

3. Tie long hair back. This precaution will keep your hair out of burner flames and harmful liquids.

4. Wear shoes that cover all of your feet. Since broken glass on the laboratory floor is all too common, your feet will need more protection than that afforded by open-toed shoes or sandals.

5. Wear clothes that cover your torso and your legs to the knees. Clothing will give your body needed protection.

6. Do not smell chemicals directly. Use your hand to waft the odor to your nose.

7. Do not put flammable chemicals near open flame.

8. When heating substances in test tubes, always use wooden tongs to hold the tube, and point the tube away from yourself and your neighbors.

9. When heating liquids in test tubes, ensure that the tube is periodically removed from the flame. This will prevent the liquid from boiling and spurting out of the tube thus causing thermal and even chemical burns if the liquid is corrosive.

10. Do not engage in games or horseplay in the laboratory.

11. Be aware of your neighbors. Are they obeying the safety rules? A neighbor's accident may not hurt him or her, but it may injure you.

12. If the Bunsen burner has been used, check that the gas valve is closed once the experiment is completed. There is a tendency for the flame to go off and therefore do not assume that the gas valve is closed just because you cannot see the flame!

13. Treat all chemicals as potentially dangerous and therefore exercise extreme caution. Chemicals spilled onto the body should be treated immediately in the following ways, and report it to your laboratory instructor.

(a) If any chemical gets into the eyes, wash immediately with a gentle stream of water, holding the eyelids open. Continue washing for at least 15 minutes and then call for a doctor. If you wear contact lenses it is very important that you remove them before washing the eyes. Chemicals can become trapped behind them.

(b) If chemical is spilled on your skin, whether they are harmful or not, wash them off with water immediately and thoroughly. Further treatment will depend on the type of chemical and whether it causes burns. Even concentrated acids and bases may not seriously burn you if they can be washed off quickly and thoroughly.

(c) If chemical is spilled on your clothing, it may be necessary to remove the article of clothing, especially if the chemical is an acid or base or if a large quantity of any chemical is spilled. Use the safety shower in the toilets next to the laboratory if necessary.

14. Wash your hands before leaving the laboratory.

15. Adhere to any other instructions and precautions stated in each experiment.


2. LABORATORY HOUSEKEEPING RULES Good housekeeping will lead to a safe and pleasant workplace. Please observe the following safety rules. 1. Each laboratory bench should only accommodate four students, two on each side. Please

adhere to such an arrangement unless instructed otherwise by your laboratory instructor.

2. Clean up broken glass immediately. Inform your laboratory instructor who will direct you to place it in special containers.

3. Take beakers, test tubes, etc., to the stock containers. Do not bring stock chemicals to your laboratory bench.

4. Clean up solid and liquid chemical spills on the bench top, electronic balances, etc. immediately. Every bench is provided with a towel; wet the towel before using it to clean up the spill and remember to rinse the towel afterwards so that the next person using it will not come into contact with the chemical. If necessary, wipe the surface again so as to ensure that the spill has been completely removed. If a large quantity of chemical is spilled, consult your instructor concerning the best method to clean it up. Large quantities of concentrated acids and bases need to be neutralized first.

5. Read the label on a bottle carefully. Is it the correct chemical?

6. Do not take excessive chemicals from the stock container. Just take slightly more than what is required for your experiment.

7. Do not put hot containers on the bench top. Let the container cool down for a while to prevent scorching of the bench top.

8. Do not throw any solids into the sink but throw it into the rubbish bin. Any solids thrown into the sink will eventually clog the sink.

9. Rinse well with water all the glassware that you have used.

10. For any chemicals (including acids) that can be poured into the sink, pour them directly into the sink hole and flush the sink with a large amount of water. This will ensure that any vapor from the chemical will not be present in the laboratory and the flushing will dilute out the chemical, thereby preventing the corrosion of the pipes.

11. Wipe the bench top with a wet towel at the end of the experiment. Remember to rinse the towel well with water after you have wiped the bench top.

B. PRESENTATION OF LABORATORY REPORT

A laboratory report should include the following:

1. Title


Give the title of the experiment.

2. Aim

The aim(s) of the experiment must be stated clearly. Use your own words to formulate the aim(s); DO NOT plagiarize from the laboratory manual.

3. Procedure

The details of the procedure should be given but generally, students would just plagiarize from the laboratory manual. Therefore, a simple statement like “Please refer to the laboratory manual.” would be sufficient, unless your laboratory instructor requires you need to write the procedure. However, any precautions that you take in order to obtain accurate measurements must be given.

4. Result

Present your observations and data in a manner which is easy to understand, e.g., in a table form.

5. Calculations

All manipulations of data are to be given in this section. If any graphs have to be drawn, arrange them in a logical sequence, and assign a number to each one of them. For more details, refer to section D on “Handling of Data”.

6. Conclusion

A clear statement for the result is as essential as an accurate aim of an experiment.

7. Discussion

The rationale of the precautions that you have taken during the course of the experiment is stated here. You can also discuss any limitations inherent with the particular experimental procedure used that would lead in inaccurate result, and if possible, suggest ways of improving it. In addition, with some experiments where the experimental values obtained can be compared with theoretical values calculated, comment on the comparison. If there is a poor correlation between the theoretical and experimental values, provide scientifically sound explanations that can account for the discrepancy.


EXPERIMENT 1

Planning Skills I Do Not Carry Out Your Plan

1 Caesium nitrate, CsNO3, decomposes on heating.

The decomposition is represented by one of the following equations.

4CsNO3(s) → 2Cs2O(s) + 4NO2(g) + O2(g)

2CsNO3(s) → 2CsNO2(s) + O2(g)

You are to devise a method of heating the solid nitrate, collecting the gas given off and measuring its volume.

From the experimental results you are to determine which is the correct equation for the decomposition.

(a) Draw and label the apparatus you would use to heat the caesium nitrate, to collect the gas and to measure its volume.

When labelling your diagram include the volume of apparatus used (e.g. 250 cm3 beaker) where appropriate.

50

40

30

20

10

0

Information that may be used in the question.

The molar volume of gas, Vm, is 24.0 dm3 mol–1 under room conditions.

Nitrogen dioxide, NO2, a toxic gas, is soluble in water.

Oxygen, O2, is not soluble in water.

[Ar: Cs, 133.0; N, 14.0; O, 16.0.] (b) Which gas/gases are you collecting in your apparatus?

50 cm3 burette

500 cm3 beaker 250 cm

3

conical flask

CsNO3

Borosilicate

test-tube


………O2……………………………………………………………………………...

(c) Taking into consideration the gas/gases you will collect and the capacity of the collecting apparatus, use the equations to calculate an appropriate mass of caesium nitrate to be heated. [Show your working]

......50 cm3 of gas is to be collected. Moles of O2 = 50x10-3/ 24 dm3 mol-1 =0.0021 moles

…mass of CsNO3 required:…………………………………………………………………...

…If equation 1 is correct: moles of CsNO3 required = 4(0.0021) = 0.0084

mass of CsNO3 required = 0.0084 x 195 = 1.625 g

If equation 2 is correct: moles of CsNO3 required = 2(0.0021) = 0.0042

mass of CsNO3 required = 0.0042 x 195 = 0.813 g

…if 1.625 g of CsNO3 was used and equation 2 is correct, then 100 cm3 of O2 woul be produced

0.813 g of CsNO3 should be used in this eperiment ………………………………...

…………………………………………….……………………………………………………..

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(d) Indicate how you would use your results to find the correct equation for the thermal decomposition of caesium nitrate.

If heating 0.813 g of CsNO3 produces 50 cm3 of O2, then equation 2 is correct

But if only 25cm3 of O2 is produced then equation 1 is correct……………

…...………………………………………………………………………………………………

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(e) Suggest one safety precaution that should be undertaken during this experiment and the reason for it.

The gases must be passed through water to absorb the toxic NO2 gas. The experiment should be conducted in a fume chamber in case any toxic NO2 absorbs.

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………………...………………………………………………………………………………… 2 The aim of this experiment is to determine the rate of reaction between iron(III) ions and iodide

ions by carrying out four experiments. In each experiment, the iron (III) ions will react with iodide ions to produce iodine. At first, the iodine produced reacts with a fixed amount of sodium thiosulphate, Na2S2O3. As soon as all the sodium thiosulphate has been used up, the iodine will react with the starch that had been added, producing a blue colour.


The rate of the reaction can be calculated from the time taken for the blue colour to appear in your reaction mixture.

During the experiments, all of your solutions and the distilled water should be kept out of direct sunlight. The following solutions were provided:

K is an aqueous solution containing iron(III) sulphate

L is aqueous potassium iodide

M is aqueous sodium thiosulphate

Starch solution

Solutions L and M were poured into two separate burettes. A measuring cylinder was used to measure out the distilled water. The reactant mixtures used in the 4 experiments were prepared as follows:

Experiment 1 2 3 4

Solution L/ cm3 8.0 5.0 6.0 10.0

Solution M/cm3 5.0 5.0 5.0 5.0

H2O/cm3 21.0 24.0 23.0 19.0

Using a small measuring cylinder, 6.0 cm3 of K was added. As K is being added, the timer or stop clock was started. The mixture was swirled and the flask placed on a white tile. The time taken for the colour to change from yellow to blue was noted and recorded in the table below

Experiment 1 2 3 4

Volume of L/ cm3 8.0 5.0 6.0 10.0

Time( t ) / s 22 56 37 15

Initial rate of reaction/ s-1 45 18 27 67

(volume of L)2 /cm6 64 25 36 100

For each experiment, assume that the initial rate of reaction can be represented by Work out the initial rate of reaction for each experiment. The concentration of iodide ions in each experiment is proportional to the volume of aqueous potassium iodide, L, used. The (volume of L)2 is therefore proportional to [I-]2. Work out (volume of L)2 for each experiment. Use the grids to plot two best-fit graphs of: 1. initial rate of reaction against volume of L. 2. initial rate of reaction against ( volume of L )2

1000 t


0-

10-

20-

30-

40-

50-

60-

70-

| | | | | | | | | | |

0 1 2 3 4 5 6 7 8 9 10

X

X

volume of

L/cm3

X

X

volume of L/cm3

Init

ial

rate

of

rea

cion

/s-1

Initial rate vs volume of L


0-

10-

20-

30-

40-

50-

60-

70-

| | | | | | | | | | |

0 10 20 30 40 50 60 70 80 90 100

X

X

X

X

(volume of L)2/cm

6

Init

ial

rate

of

rea

cion

/s-1

Initial rate vs volume of L2


(a) From your graph, suggest the order of the reaction with respect to iodide ions. Justify your answer

……2nd order because the initial rate of reaction is proportional to L2. This is obtained from the graph

of initial rate vs (volume of L)2 . The graph gives a straight line passing through the origin.

(b) State and explain one safety precaution you took while doing your experiments.

Both KI and sodium thiosulphate are corrosive/oxidants. Therefore use gloves, safety goggles and lab

coat.…………………………………………………………………………………………………………………

……………………………………………………………………………………………………………………….

(c) Explain why the same total volume of solution was used in the reaction mixture in all four experiments of the rate investigation.

If a constant volume is used than the volume of L used in the experiment is proportional to its

concentration. There is no need to recalculate the [L] for each experiment and the voulme can be used

to calculate the rate order of the reactants.

………………………………………………………………………………………………………………………

(d) Identify three sources of inaccuracy in your experimental procedure for the rate investigation and suggest ways of reducing them.

1. measuring cylinders are not accurate and therefore use burettes or graduated pipettes.

2. volumes of solutions used are small which can result in large errors. Therefore use larger volumes

3. Temperature changes can affect rates of reaction. Therefore conduct the experiments in water

baths maintained at constant temperature

4. Colour changes at end points are not clear for reactions that are slow( or solutions that are dilute).

Therefore use higher concentrations to increase rate of reaction( or make the colour more darker)

5. Inaccuracy/difficulty in timing the exact start of reaction. Not all the reactants were added at exact

time the timer was started. Therefore et someone to start timing.

6. the volume of starch added in each experiment should be the same so that the total volume of

reacting mixture is constant.

7 Each experiment was conducted only once. Therefore repeat the experiments to get accurate

values

(e) Assess both the accuracy and the reliability of your results from the rate investigations, giving your reasons.

Ans can be ‘yes’ or ‘no’

The results are both accurate and reliable. All the points lie close to the line of best fit

The results are not accurate and reliable. There is insufficient evidence to draw a conclusion. A wider range of [L] should have been used/more experiments should have been used


(f) Outline how you could adapt the experimental procedure to find out whether the order of the reaction with respect to iron (III) ions is first order.

Conduct a series of experiment in which the volume of L and M are kept constant but volumes of K and water are varied.

Ensure that the total volume of reaction mixture is kept the same for all experiments by adding different volumes of water.

Record the time taken for the appearance of blue colour.

Plot a graph of initial rate (1000/t) against volume of K

The rate w.r.t to Fe3+ is 1st order if the graph is a straight line passing through the origin.

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EXPERIMENT 2

PLANNING SKILL II

Do Not Carry Out Your Plan

1 A solid known as „washing soda‟ consists of crystals of sodium carbonate-10-water, Na2CO3.10H2O.

On exposure to air, the crystals slowly lose their water of crystallisation to become Na2CO3.xH2O.

You are provided with a sample of „washing soda‟ that has lost some of its water of crystallisation.

You are to devise a method involving weighing and heating of the sample that will enable you to find the percentage of water that has been lost from the crystals on exposure to air.

You may assume that anhydrous sodium carbonate is not decomposed by heat.

The normal apparatus to be found in a school Chemistry laboratory is available.

Outline, by a series of numbered steps, the apparatus to be used and the experimental work you would carry out, including the measurements to be taken, to find the mass of water remaining in the crystals.

1 an empty, dry crucible with lid was weighed and its accurate mass was recorded.....

2 approximately 5 g of washing soda was placed in the crucible……………………….....

3 The crucible with its lid and contents was weighed and the accurate weight was recorded….........

4 the sample in the crucible was initially heated gently for three minutes and then heated strongly for

a further 5 minutes, without the lid........

5 the crucible and its content was then cooled with the lid on …

6 The cooled crucible with lid and the content was then weighed and the weight was recorded.

7 steps 4 to 6 was repeated until a constant weight was obtained

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Draw the table of results you would prepare to record the results of your experiment.

Before heating

After heating

1 2 3 4

Mass of crucible, lid and residue /g X Y Z z

Show how you would use the results of your experiment to find the percentage of water that has been lost by the crystals on standing in air.

Make use of the relative atomic masses in this section.

[Ar; Na, 23.0; C, 12.0; O, 16.0; H, 1.]

Mass of crucible, lid /g a

Mass of crucible, lid and sample /g b

Mass of sample /g c


……mass of anhydrous Na2CO3 in sample= z – a = f………………………………….........

……mass of H2O in Na2CO3.xH2O sample= c – d = h……………………………………………….........

…… moles of anhydrous Na2CO3 = f = g …………………………………………………

.…………………………………………106………………………………………………………

…… mole of H2O in Na2CO3.xH2O = h = j…………………………………………......... …………………………………………18

……x = j …which is small whole number ………… …… g

.…% water lost = (10 – x) × 100% …………………… …

x

………………………………………………………………………………………………….........

2 Calcium hydroxide is an ionic compound that has a low solubility in water.

The solubility of such a compound can be represented using a type of equilibrium constant called the solubility product, Ksp.

In this practical the solubility product of calcium hydroxide was determined by a titration method using sulphamic acid and an indicator that may be not familiar to you.

For the titration three materials were provided.

Solution Y is a saturated aqueous solution of calcium hydroxide, Ca(OH)2.

Solid Z is sulphamic acid, NH2SO3H, which is a strong monobasic acid.

Bromothymol blue is an indicator that changes colour at about pH = 7.

Approximately 1.0 g of Z was placed into a preweighed glass bottle. The accurate mass of glass bottle + sulphamic acid was recorded. The solid Z was tipped into 100 cm3 of deionised water and dissolved. The glass bottle + residue were reweighed and recorded. The solution was transferred to a 250 cm3 volumetric flask. Sufficient water was added until the 250 cm3 mark. The solution Z was then poured into a burette. 25 cm3 of solution Y was transferred into a conical flask using a pipette. 10 drops of bromothymol blue was added. Titration was then carried out. The titration was repeated until two consistent titres were obtained.

Results Weighings


Mass of empty bottle/ g 35.63

Mass of bottle + sulphamic acid /g 36.75

Mass of bottle + residue /g 35.71

Mass of sulphamic acid /g 1.04

Titrations

Titration 1 2 3 4

Final burette reading/ cm3 28.20 29.10 28.15 28.25

Intial burette reading / cm3 0.05 1.10 0.05 0.20

Volume of Z used / cm3 28.15 28.00 28.10 28.05

Use the additional columns of the table to record calculated values that will enable you to determine Ksp of calcium hydroxide. Label the columns you use, including the units. Analysis (a) Calculate the concentration, in mol dm-3, of NH2SO3H in your solution of Z. Moles of Z used = 1.04 g = 1.0107 mol 97 g mol-1

[NH2SO3H] = 0.0107 moles = 0.0428 mol dm-3

0.25 dm3

(b) Calculate the amount, in moles, of NH2SO3H in your mean titre. Mean titre = 28.15 + 28.00 + 28.10 + 28.05 = 28.075 cm3

4 Moles = (28.075 × 10-3)dm3 × 0.0428 mol dm-3 = 1.202 × 10-3 mols (c) Complete the balanced equation for the neutralization reaction between calcium hydroxide and

sulphamic acid in the titration. Give state symbols.

…Ca(OH)2..( s ) + 2 NH2SO3H( aq ) (NH2SO3)2Ca ( aq ) + 2H2O…..( l ) (d) Calculate the concentration, in mol dm-3, of Ca(OH)2 in solution Y. 1 mole Ca(OH)2 reacts with 2 moles NH2SO3H


moles of Ca(OH)2 = 1.202 × 10-3 = 0.601 × 10-3 2 [Ca(OH)2] = 0.601 × 10-3 mols = 0.024 mol dm-3

(25 × 10-3) dm3

(e) aqueous calcium hydroxide is fully ionized in solution Y.

Ca(OH)2(aq) Ca2+(aq) + 2OH-(aq)

(i) Use your answer to (d) to deduce the calcium ion concentration, [Ca2+(aq)], in mol dm-3 in solution Y.

[Ca2+(aq)] = 0.024 mol dm-3 (ii) Deduce the hydroxide ion concentration, [OH-(aq)], in mol dm3, in solution Y. [OH-(aq)] = 0.024 × 2 = 0.048 mol dm-3 (f) (i) Write the expression for the solubility product of calcium hydroxide. Ksp = [Ca2+(aq)] [OH-(aq)]2

(ii) Use this expression to calculate the solubility product, Ksp, of Ca(OH)2. Deduce the units. (0.024)(0.048)2 = 5.55 × 10-5 mol3 dm-9

Questions 5 cm3 of saturated aqueous calcium hydroxide, solution Y, was transferred into a test tube. Approximately an equal volume of aqueous sodium hydroxide, of concentration 2.0 mol dm-3, was added. The tube was allowed to stand. (a) What observation would you make? ……white precipitate will be formed

…………………………………………………………………………….............................


(b) Use Le Chaterlier‟s principle to explain why the observation in (a) was made.

When NaOH is added it increases [OH-].

This causes the equilibrium to shift to the left/reverse in order to reduce [OH-]

Ca(OH)2 solid is formed

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(c) Why is special care needed during the test you have just carried out?

1 NaOH is corrosive and should not be spilled on hands.

2 Ksp is temperature dependent. a constant temperature should be maintained.

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(d) Bromothymol blue indicator was used in this experiment. It changes colour at about pH = 7. A student suggested that phenolphthalein, which changes colour at about pH = 9, is a better indicator to use.

State and explain whether or not you agree with this suggestion. (Note: calcium hydroxide is a strong alkali and sulphamic acid is a strong acid) Both may be used.

Titration of strong acid with strong base results in pH change form 3 to 10 at end point.

A suitable indicator changes colour within about two pH units

The colour change of bromothymol is hard to discern but colour change of phenolphthalein is

very sharp.

A student attempted to determine the solubility of Ca(OH)2 by a different method. He added an excess of aqueous copper(II) nitrate to 25 cm3 of a saturated solution of calcium hydroxide. The reaction below took place.


Cu(NO3)2(aq) + Ca(OH)2(aq) Cu(OH)2(s) + Ca(NO3)2(aq)

The student filtered and then weighed the precipitated copper(II) hydroxide. He obtained 0.05 g. (a) Write the ionic equation for the reaction. Cu2+(aq) + 2OH-(aq) → Cu(OH)2(s) (b) Calculate the concentration of Ca(OH)2, in mol dm-3, in the saturated solution, using the

student‟s data. Moles of Cu(OH)2 = 0.05 g = 5.13 × 10-4 moles 94.5 g mol-1

moles of Ca(OH)2 = 5.13 × 10-4 moles [Ca(OH)2] = 5.13 × 10-4 moles = 0.025 mol dm-5

(25 × 10-3) dm3

(c) Compare the results obtained in both methods. Explain which of the two methods gives a more reliable determination of the solubility of

Ca(OH)2. Titration method is more reliable: 1 burettes and pipettes are more accurate/ less error

2 easy to determine when reaction is over, using indicators

3 Titration readings are repeated – consistent readings are evidence of reliability

Gravimetric method: 1 masses are too small so high error

2 filtration and drying can lead to loss in mass

3 if not dried completely, ppt will have water of hydration and mass will be overestimated


EXPERIMENT 3

PLANNING SKILL III 1 Do Not Carry Out Your Plan

Copper carbonate

decomposes on strong heating,

CuCO3(s) → CuO(s) + CO2(g)

reacts with hydrochloric acid.

CuCO3(s) + 2HCl(aq) → CuCl2(aq) + CO2(g) + H2O(l) Carbon dioxide, CO2, is an acidic oxide. Copper oxide, CuO, is a basic oxide. Malachite is a rock containing a high percentage of copper carbonate. You are to plan an experiment to determine the % of copper carbonate in a specimen of malachite.

In your plan you may use any equipment normally found in a school laboratory.

You may assume that any other material present in the malachite is unaffected by heating and is neither acidic or basic.

In your plan you must include masses, volumes, concentrations of materials used as appropriate.

You will need to use some or all of the following information. [Cu, 63.5; C, 12.0; O, 16.0.] 1 mole of any gas occupies a volume of approximately 24.0 dm3 at room temperature and pressure.

Plan Method 1

1 an empty, dry crucible with lid was weighed and its accurate mass was recorded.....

2 approximately 5 g of malachite was placed in the crucible……………………….....

3 The crucible , its lid and contents was weighed and the accurate weight was recorded….........

4 the sample in the crucible was heated strongly for a r 5 minutes, without the lid........

5 the crucible and its content was then cooled with the lid on …

6 The cooled crucible with lid and the residue was then weighed and the weight was recorded.

7 steps 4 to 6 was repeated until a constant weight was obtained

method 2

1 an empty, dry weighing bottle was weighed and its accurate mass was recorded.....


2 approximately 5 g of malachite was placed in the weighing bottle and the weighing bottle with it contents was weighed and the accurate weight was recorded….........

3 A conical flask containing 100 cm3 of 2.0 mol dm-3 HCl solution was placed on electronic balance and the reading was zeroed. ........

4 the malachite sample was tipped into the conical flask and reaction allowed to proceed.

5 the final mass of conical flask and malachite was recorded …

6 the mass of the weighing bottle + malachite residue was then recorded .

Show how you would tabulate your results. All necessary measurements should be shown in the table(s). Method 1 Before heating

mass of empty crucible+ lid /g B

Mass of malachite + crucible + lid/g A

After heating

1 2 3 4 5

Mass of residue+ crucible + lid/g x y z z

Method 2

mass of malachite sample + weighing bottle /g e

Mass of residual malachite + weighing bottle /g F

Mass of conical flask + acid/g 0(zero)

Mass of conical flask + acid + malachite residue/g H

Outline how you would process your results to find the % of copper carbonate in the sample of malachite.

Method 1

Mass of malachite used = A – B = K

Mass of CO2 lost from malachite sample= A – z = L

Moles of CO2 lost = L/44 = R

moles of CuCO3 in malachite sample = R

mass of CuCO3 in malachite sample = R × 123.5 = S

mass % of CuCO3 in sample = S × 100% K

Method 2


Mass of malachite used = e – F = M

Mass of CO2 lost from malachite sample= M – H = N

Moles of CO2 lost = N/44 = P

moles of CuCO3 in malachite sample = P

mass of CuCO3 in malachite sample = P × 123.5 = Q

mass % of CuCO3 in sample = Q × 100% M

2 Part 1 We can find out how much nickel is present in a nickel(II) compound by a procedure called back-titration. This titration uses edta. Edta is a ligand which forms complex ions.

Excess edta is added to a solution of a nickel(II) compound.

One mole of nickel(II) ions reacts with one mole of edta.

The amount of edta left over is estimated by titration with magnesium sulphate.

One mole of magnesium ions reacts with one mole of edta. In this practical you will determine the amount of nickel(II) ion in a sample of hydrated nickel(II) sulphate, NiSO4.xH2O. You will then determine the number, x, in this formula Five labelled materials are provided.

E is hydrated nickel (II) sulphate. (harmful)

F is aqueous edta of concentration 0.110 mol dm-3.

G is aqueous magnesium sulphate of concentration 0.0800 mol dm-3.

H is a buffer solution. (irritant)

A solid indicator called solochrome black. Approximately 2.90 g of E was placed into a preweighed glass bottle. The accurate mass of glass bottle + sulphamic acid was recorded. The solid E was tipped into 150 cm3 of deionised water and dissolved. The glass bottle + residue were reweighed and recorded. The solution was transferred to a 250 cm3 volumetric flask. Sufficient water was added until the 250 cm3 mark. 25 cm3 solution E was then transferred into a conical flask using a pipette Then the following were added to E:

25 cm3 of solution F (edta), using another pipette ;

About 25 cm3 of H, the buffer solution, using a measuring cylinder;

A few specks of solochrome black indicator. The burette was filled with solution G, aqueous magnesium sulphate. The titration was then carried out. The procedure was repeated to obtain two accurate titres

Results


Weighings

Mass of empty bottle/ g 35.63

Mass of bottle + E /g 38.58

Mass of bottle + residue /g 35.67

Mass of E used /g 38.58 -35.67 = 2.91

Titrations

Titration 1 2 3 4

Final burette reading/ cm3 28.20 21.80 20.65 20.85

Intial burette reading / cm3 0.05 1.10 0.05 0.20

Volume of G used / cm3 28.15 20.70 20.60 20.65

Use the additional columns of the table to record calculated values that will enable you to determine the molar mass of the nickel (II) compound. Label the columns you use, including the units. Analysis: In this section all your working must be shown clearly (a) Calculate the amount, in moles, of MgSO4 used in the mean titre. Mean titre = 20.70 + 20.60 + 20.65 = 20.65 cm3

3 Moles of G = 0.08 × (20.65 × 10-3) = 0.00615 moles (b) What amount of edta, in moles, reacted with this amount of MgSO4? 1 mole of MgSO4 reacts with 1 mole edta

mole of edta reacted = 0.00615 moles (c) Calculate the total amount, in moles, of edta used in each titration.


0.11mole dm-3 × (25 × 10-3) dm3 = 0.00275 moles (d) Calculate the amount, in moles, of edta which reacted with 25.0 cm3 of solution E. Moles of edta reacted with E = total moles used(c) - moles reacted with MgSO4 (b) = 0.00275 – 0.00615 = 0.0011 (e) Calculate the amount, in moles, of Ni2+ in 250 cm3 of your solution. 0.0011 × 250 = 0.011 moled 25 (f) Calculate the mass of one mole of NiSO4.xH2O. 0.011 moles NiSO4.xH2O weighs 2.91 g 1.0 mole NiSO4.xH2O weighs 2.91 g = 264.55 g mol-1

0.11 mols xH2O weighs (264.55 –[58.7 + 32.1 + 64]) = 109.75 g x = 109.75 = 6.10 18 Part 2 A student determined the value of x in NiSO4.xH2O by another method. This method involved heating the hydrated compound in order to drive off the water of crystallization and obtain anhydrous nickel(II) sulphate.

The following readings were obtained Mass of empty crucible = 20.00 g Mass of crucible + hydrated nickel(II) sulphate = 21.31 g Mass of crucible and residue, after heating = 20.77 g

(a) Assuming that the residue is anhydrous nickel(II) sulphate, calculate the moles of anhydrous nickel(II) sulphate obtained.

Mass of anhydrous NiSO4 = 20.77 – 20.00 = 0.77


Moles of NiSO4 = 0.77 = 4.97 × 10-3

154.8 (b) Calculate x in the formula NiSO4.xH2O. Mass of H2O in sample = 21.31 – 20.77 = 0.54 g Moles of H2O in sample = 0.54 = 0.03 18 x = 0.03 = 6.03 4.97 × 10-3

(c) State and explain three precautions that the student should take to ensure that the experiment

was as accurate as possible.

1 Initially cover the crucible with a lid during heating to prevent spitting/spluttering of solid. Later remove the lid to allow H2O to escape. During cooling cover the lid to prevent condensation /absorption of water.

2 dehydration maybe incomplete. Heat strongly for a longer period . reheat and reweigh until a constant mass is obtained.

3 Mass of salt used is small resulting in large percentage error. Use a larger mass of NiSO4.xH2O

4 Heat used may be too strong and cause NiSO4 to decompose, forming oxides of sulphur

5 Expt was done only once. Repeat the experiment until consistent values are obtained.

6 ensure that there is no carbon (from bunsen) accumulating on the crucible, resulting in higher mass of crucible and residue

………………………………………………………………………………………………………………………

………………………………………………………………………………………………………………………

………………………………………………………………………………………………………………………

………………………………………………………………………………………………………………………

………………………………………………………………………………………………………………………

………………………………………………………………………………………………………………………

……………………………………………………………………………….

(d) Which result for the value of x is probably more reliable, your own titration or the result from the

student‟s experiment? Justify your answer.

Titration method is more reliable


Burettes and pipettes are more accurate.

Titration is repeated until consistent titres are obtained

End point of reaction is visible through use of indicator

Student‟s method

1. no evidence the sample is completely dry

2 involves too many weighing and can result in cumulative error


EXPERIMENT 4

PLANNING SKILL IV 1

You are provided with solid sodium hydrogen carbonate, NaHCO3; sodium carbonate, Na2CO3; HCl and distilled water.

Using these materials alone, you are to plan and carry out one experiment to determine a

further enthalpy change, H4, which can be put together with the other enthalpy changes to

find the enthalpy change for the following reaction.

Na2CO3(s) + CO2(g) + H2O(l) → 2NaHCO3(aq) H4

using

Na2CO3(s) + 2HCl(aq) → 2NaCl(aq) + CO2(g) + H2O(l) H1

NaHCO3(s) + HCl(aq) → NaCl(aq) + CO2(g) + H2O(l) H2

Na2CO3(s) + CO2(g) + H2O(l) → 2NaHCO3(s) H3

Give your plan to determine H1 and H2 as a series of numbered steps

…About 5.0g of Na2CO3 (0.047 moles) was weighed into a preweighed weighing bottle and the accurate mass recorded

…50 cm3 of 2.0 mol dm-3 HCl solution was placed into a polystyrene cup, using a burette.

…The intial temperature of the HCl solution was measured and recorded

…The NaCO3 was tipped into the HCl solution and covered with a lid that has a hole.

….The solution was gently strirred with a thermometer and the final temperature of the solution was recorded

…The weighing bottle with residue was reweighed

…The mass of Na2CO3 tipped into the HCl solution was determined

…The heat change for rection was calculated : q = mc

…….…………………………………where q = heat of reaction

…….…………………………………………m = mass of solution (50 g)

…….…………………………………………c = 4.18…………………………………

…….………………………………………… = Change in temperature

…….……………ΔH = q …

…….…………………moles of Na2CO3………………………………………

…The experiment was repeated using NaHCO3 and H2 was estimated


……………………………………………………………………………………………………….

…………………………………………………………………………………………………….…

Use the equations

Na2CO3(s) + 2HCl(aq) → 2NaCl(aq) + CO2(g) + H2O(l)

NaHCO3(s) + HCl(aq) → NaCl(aq) + CO2(g) + H2O(l)

and the calculated values of H1 and H2 to calculate the enthalpy change, H3, for the

following reaction, where the enthalpy change cannot be measured directly by experiment.

Na2CO3(s) + CO2(g) + H2O(l) → 2NaHCO3(s)

Using Hess Law

……..…… Na2CO3(s) + 2HCl(aq) → 2NaCl(aq) + CO2(g) + H2O(l) = H1

……………2NaCl(aq) + 2CO2(g) + 2H2O(l) → 2NaHCO3(s) + 2HCl(aq) = -2( H2)

………………… H3 = H1 - 2( H2)

………………………………………….……………………………………………………………………………

………………………….……………………………………………………………………………………………

Show how you can use your results to determine the enthalpy change for the reaction

Na2CO3(s) + CO2(g) + H2O(l) → 2NaHCO3(aq)

…A known mass of NaHCO3 solid was completely dissolved in excess water.

The temperature change was determined and enthalpy change for the reaction

NaHCO3(s) → NaHCO3(aq) was estimated …= H5…

using Hess Law : Na2CO3(s) + CO2(g) + H2O(l) → 2NaHCO3(s)

2NaHCO3(s) → 2NaHCO3(aq)

… Na2CO3(s) + CO2(g) + H2O(l) → 2NaHCO3(s)

H = H3 + 2( H5) …………….……………………………………………………………………………………………………….

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2 Part 1 Introduction Ethanoic acid and ethanol react reversibly to form ethyl ethanoate and water according to the following equation.

CH3COOH + CH3CH2OH → CH3COOCH2CH3 + H2O

A mixture of 0.420 g of ethanoic acid and 0.0100 mol of ethanol was allowed to reach equilibrium at 20 °C. The ethanoic acid remaining in the equilibrium mixture reacted exactly with 3.00 cm3 of 0.500 mol dm–3 sodium hydroxide solution added from a burette. Analysis Full marks can only be scored in calculations if you show all of your working. a Write an expression for the equilibrium constant, Kc, for the reaction between ethanoic acid and

ethanol. [CH3COOH] . [CH3CH2OH] [CH3COOCH2CH3] . [H2O] b Calculate the number of moles present in 0.420 g of ethanoic acid (Mr = 60.0). 0.420 = 0.007 60 c Calculate the number of moles of sodium hydroxide in 3.00 cm3 of 0.500 mol dm–3 sodium

hydroxide solution. Hence calculate the number of moles of ethanoic acid present in the equilibrium mixture.

Moles of sodium hydroxide (3.0 × 10-3)dm3 × 0.500 mol dm-3 = 0.0015 Moles of ethanoic acid 1mol NaOH reacts with 1 mole CH3COOH

moles of ethanoic acid = 0.0015 moles

Kc =


d Use your answers from parts b and c to calculate the number of moles of acid that have reacted with the ethanol. State the number of moles of ester and of water present in the equilibrium mixture.

Moles of acid reacted 0.007 – 0.0015 = 0.0055 Moles of ester 1 mole acid reacts to form 1 mole ester 0.0055 moles acid will produce 0.0055 mole ester 0.0055 mols Moles of water 0.0055 mols e Calculate the number of moles of ethanol present in the equilibrium mixture and hence

calculate the value of Kc. 0.0055 moles acid reacted

0.0055 moles of ethanol will have reacted Moles of ethanol left in equilibrium mixture = 0.0100 – 0.0055 = 0.0045 f Assume that the maximum errors for the apparatus used in this experiment were

balance total error ± 0.001 g burette total error ± 0.15cm3 (from two readings and an end-point error)

Calculate the maximum percentage error in using each piece of apparatus, and hence the maximum overall percentage error.

Balance : 0.001 × 100% = 0.24% 0.42 Burette : 0.15 × 100% = 5% 3 Overall error = (0.24)2 + (5)2 = 5.006% Part 2


Full marks can only be scored in calculations if you show all of your working. a A data book value for Kc for this reaction is 3.92 at 20 °C. Calculate the difference between the

value calculated in part e of the Part 1 section and the data book value. Express this difference as a percentage of the data book value. (If you could not complete the calculation in part e of the Part 1 section, you should assume a value of 3.40 for Kc. This is not the correct value.)

Difference 4.48 – 3.9 = 0.56 Percentage 0.56 × 100% = 14.4% 3.92 b Comment on the quality of the experimental result for Kc.

.....The percentage error is greater than apparatus error.

This may be due to

(1) a disturbance in equilibrium due to temperature changes,

(2) some procedural error

(3) operator error ...................................................................................................................................................................

...................................................................................................................................................................

................................................................................................................ c The volume of the sodium hydroxide solution used in this titration was very small. Suggest an

alternative concentration for the sodium hydroxide solution, and explain how this would improve the accuracy of the experiment.

...use 0.1 mol dm-3 solution of HCl

Then larger volumes of solution can be used to reduce burette error. It also gives a more accurate end-point. ...................................................................................................................................................................

d State why it is necessary to maintain a constant temperature in this experiment. Suggest one method for maintaining a constant temperature in this experiment.

............Kc is temperature dependent

use a thermostated water bath

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EXPERIMENT 5

PLANNING SKILL V 1 Hydrochloric acid reacts with solid calcium carbonate in the form of marble chips. The products of

the reaction are aqueous calcium chloride, carbon dioxide gas and water.

2HCl(aq) + CaCO3(s) CaCl2(aq) + CO2(g) + H2O(l)

You are to plan an investigation to find how changes in one variable affect the rate of reaction.

(a) (i) The concentration of the hydrochloric acid is one independent variable that could be

investigated. Identify two other independent variables that could be investigated. temperature and size of marble chips

(ii) Identify a dependent variable that you could measure to follow the rate of reaction. volume or mass of CO2

(b) Propose a hypothesis that can be tested experimentally and that links the rate of reaction

and the concentration of the hydrochloric acid.

(c) Design a laboratory experiment that you would use to investigate your hypothesis. You should draw a labelled diagram showing the arrangement of your apparatus. Your plan should include the following.

(i) a step-by-step description of the procedure you would follow

(ii) how you would measure the independent variable

(iii) how you would vary the independent variable and how you will ensure accuracy in your results

(iv) how you would control the other two variables that you identified in part (a)

(v) any safety precautions that you would take

(vi) appropriate quantities (masses, volumes, concentrations) to use in the experiment [Ar: Ca, 40.1; C, 12.0; O, 16.0]

Concentrated hydrochloric acid, a corrosive liquid, has an approximate concentration of 11 mol dm

-3.

Labelled diagram

50

40

30

20

10

0


Plan …

1. The apparatus was set up as shown in the diagram and placed in a thermostated water bath

2. 0.5 g of sodium carbonate chips of similar size was weighed and placed in the conical flask…

3. A thistle flask containing 100 cm3 of 0.20 mol dm-3 HCl solution was firmly attached to mouth of

conical flask……

4. The HCl was then released into the conical flask containing the sodium carbonate and the delivery

tube placed below the burette……

5. The stopwatch was immediately started and the volume of gas collected was noted every 15

seconds for 3 minutes……

6. the experiment was repeated using 100 cm3 of 0.4, 0.6, 0.8 and 1.0 mol dm-3 HCl solutions but

keeping the mass of sodium carbonate the same

…or …………………………………………………………………………………………………

1. The apparatus was set up as shown in the diagram and placed in a thermostated water bath

2. 0.5 g of sodium carbonate chips of similar size was weighed and placed in the conical flask…


conical flask……

4. The HCl was then released into the conical flask containing the sodium carbonate and the delivery

tube placed below the burette……

5. The stopwatch was immediately started and the volume of gas collected in 15 seconds was noted


keeping

…….…………………………………………………………………………………………………

……..………………………………………………………………………………………………...

……………………………………………………………………………………………………….

…………………………………………………………………………………………………….…


(d) Prepare a table showing all of the data you wish to record in the experiment and any columns for processing the data recorded. Explain how you will use the data to confirm or reject the hypothesis proposed in (b).

For method 1

Time/s [HCl]/ mol dm-3

Volume of CO2 collected/cm3

15 30

45

60

75

90

105

120

135

150

165

180

0.2

0.4

0.6

0.8

1.0

For method 2

[HCl]/ mol dm-3

Volume of CO2 collected/cm3

Rate / cm3 s-1

(vol./15 s)

0.2

0.4

0.6

0.8

1.0

Method 1:

A graph of time(s) vs volume of CO2 collected was drawn for each concentration of HCl used.

A tangent was drawn at the initial part of the curve and the rate of reaction was calculated.

The rates of reaction were then compared to see how [HCl] affects rate of reaction

Method 2……………………………………………………………………………………

The rate of reaction was determined by using, rate = vol. of gas collected(cm3) time(s)

The rates for different concentrations of HCl were compared to see how [HCl] affects rate of reaction


2 A student makes the following prediction.

“The amount of carbon dioxide evolved when a metal carbonate, X2CO3, reacts with an excess of an acid is directly proportional to the moles/mass of the carbonate used.”

If this statement is true the relative molecular mass of X2CO3 can be determined by the following simple experiment.

• The mass of a weighing bottle + X2CO3 is measured.

• 50 cm3 of 2.0 mol dm-3 hydrochloric acid (an excess) is transferred from a measuring cylinder to a 250 cm3 conical flask.

• Sodium carbonate is added to the acid in the flask to saturate the acid with carbon dioxide so that none of the gas given off when X2CO3 reacts with the acid will dissolve in the acid.

• A cotton wool plug is placed in the neck of the flask and the flask + acid is weighed.

• The weighed X2CO3 is added to the acid in the flask and the cotton wool plug quickly replaced in the neck of the flask to prevent any loss of acid spray.

• The empty weighing bottle is weighed.

• When reaction in the flask has stopped, the flask is left for 10 minutes to allow carbon dioxide to diffuse from the flask.

• The flask and its contents after the reaction are weighed.

The results from a group of students are given in the table overleaf. A B C D E F G H

mass of weighing bottle + X2CO3 /g

mass of empty weighing bottle /g

mass of flask + acid before the reaction /g

mass of flask + solution after the reaction /g

mass of X2CO3 used /g

(A–B)

mass of CO2 evolved /g

(E + C– D)

14.29 11.48 221.35 223.21

2.81 0.95

16.41 11.76 209.71 212.91

4.65 1.45

12.24 11.34 210.45 210.97

0.90 0.38

16.77 11.27 214.38 217.80

5.50 2.08

16.48 10.68 211.63 215.59 5.80

1.84

14.85 11.15 217.18 219.68

3.7 1.2

13.81 11.61 212.12 213.76 2.2

0.56

18.81 11.06 218.65 224.25

7.75 2.15

16.18 10.94 206.77 210.31

5.24 1.7

17.93 11.53 221.49 225.84

6.40 2.05

14.49 11.09 217.18 219.68

3.4 0.9

18.19 10.87 215.33 220.31

7.32 2.34


(a) Use the additional columns of the table to record calculated values that will enable you to determine Mr for X2CO3 graphically. You may use some or all of the columns. Label the columns you use, including the units and an equation to show how the value is calculated. You may use the column headings A to H in these equations e.g. = C - B. (Remember that the gas is lost from the flask after the weighed carbonate is added to the weighed acid.)

(b) Present the data calculated in (a) in graphical form. Draw the line of best-fit.


0 1 2 3 4 5 6 7 8 9 Mass of XCO3/g

2.8

2.6

2.4

2.2

2.0

1.8

1.6

1.4

1.2

1.0

0.8

0.6

0.4

0.2

0

Mass o

f CO

2 form

ed/g

×

×

×

×

×

×

×

×

×

×

×

×


(c) Circle any anomalous points on the graph that you did not use when drawing the best-fit line. By reference to the description of the experiment, suggest an explanation for any anomaly.

...........The volume of CO2 evolved is too high. The cotton plug may not have been weighed

...................................................................................................................................................................

...................................................................................................................................................................

...................................................................................................................................................................

(d) Identify any part of the range of readings where the data appears to be less reliable and suggest why this should be the case.

.......When low masses are used the data appears less reliable. Low mass gives higher percentage

error and are generally unreliable

...................................................................................................................................................................

................................................................................................................................................................

(e) Use data obtained from the best-fit line you have drawn to calculate Mr for X2CO3. X2CO3(s) + 2H+(aq) → 2X+(aq) + CO2(g) + H2O(l)

...........1.4 g CO2 formed from 4.5 g X2CO3

..........nCO2 = 1.4/44 = 0.0318

Therefore 4.5 g X2CO3 is = 0.0318 moles

1 mole X2CO3 weighs 5.5/0.0318 = 141.4 g .

(f) By considering the data you have processed and the graph you have drawn, explain whether the experimental data supports the initial prediction and whether the experimental procedure described is suitable for the determination of Mr.

........It is suitable as most of the points plotted fall on or near the line of best fit

...................................................................................................................................................................

...................................................................................................................................................................

(g) Suggest an improvement to the experimental procedure that has been described.

.....more experiments should be done using lower masses of X2CO3

...................................................................................................................................................................

.............................................................................................................


(h) Use your knowledge of acid-base chemistry to describe a more appropriate way by which the Mr of a metal carbonate might be determined.

......use titration method

Titrate the X2CO3 with an acid using a suitable indicator

...................................................................................................................................................................

..........................................................................................................................................................


EXPERIMENT 6

Planning Skill VI 1 Hydrochloric acid and magnesium ribbon react to produce hydrogen gas.

2HCl(aq) + Mg(s) →MgCl2(aq) + H2(g)

You are to plan the details of an experiment, based on the volume of gas produced in the reaction, to investigate how the rate of reaction depends on the concentration of the hydrochloric acid.

(a) Using scientific knowledge linking molecular or ionic collisions to rates of reaction,

predict the relationship between

(i) the rate of formation of hydrogen gas and the concentration of the hydrochloric acid,

..... the rate of formation of H2(g) increases as the concentration of HCl increases

..........................................................................................................................................

...........................................................................................................................................

(ii) the rate of formation of hydrogen gas and the temperature of the reaction. ............ rate of reaction increaseswith increasing temperature

............................................................................................................................................

...........................................................................................................................................

(b) In an experiment to determine the rate of reaction with respect to HCl identify the

independent variable. ............. Concentration of HCl............................................................................

(c) Identify one variable, other than temperature, that must be controlled in the experiment. length/mass/surface area of the magnesium ribbon must be kept constant


(d) Design a laboratory experiment that you would use to investigate your prediction in (a)(i).

You are provided with the following.

magnesium ribbon 2 mol dm−3 hydrochloric acid

Your plan should give a step-by-step description of the method, including (i) a list of the apparatus you would use,

(ii) how you would measure the independent variable and how you would ensure that its value has been measured accurately,

(iii) how you would mix the reactants without loss of gas from the apparatus,

(iv) how you would measure the dependent variable,

(v) how you would control the other variables,

(vi) appropriate quantities (volumes, relative concentrations) to use in the experiment.

..Apparatus: thistle funnel, conical flask burette/gas syringe, beaker, delivery tube, stopwatch.

1. A piece of magnesium ribbon weighing 0.5 g was weighed and placed in a 250 cm3 conical

flask…

2. A thistle flask containing 100 cm3 of 0.20 mol dm-3 HCl solution was firmly attached to the mouth

of conical flask……

3. The HCl was then released into the conical flask containing the magnesium and the delivery tube

placed below the burette……

4. The stopwatch was immediately started and the volume of gas collected was noted every 15

seconds for 3 minutes……


keeping the mass of magnesiume the same

…or …………………………………………………………………………………………………

1. 0.5 g of magnesium ribbon was weighed and placed in the conical flask…


conical flask……


3. The HCl was then released into the conical flask containing the magnesium and the delivery tube

placed below the burette……

4. The stopwatch was immediately started and the volume of gas collected in 15 seconds was noted


keeping

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(e) Prepare a table that you could use to record all of the data from the experiment. Include columns for processing the data recorded. Explain how you would use the data to confirm or reject your prediction in (a)(i).

For method 1

Time/s [HCl]/ mol dm-3

Volume of gas collected/cm3

15 30

45

60

75

90

105

120

135

150

165

180

0.2

0.4

0.6

0.8

1.0

For method 2

[HCl]/ mol dm-3


0.2

0.4

0.6

0.8

1.0

Method 1:

A graph of time(s) vs volume of gas collected was drawn for each concentration of HCl used.

A tangent was drawn at the initial part of the curve and the rate of reaction was calculated.

The rates of reaction were then compared to see how [HCl] affects rate of reaction

Method 2……………………………………………………………………………………

The rate of reaction was determined by using, rate = vol. of gas collected(cm3) time(s)

The rates for different concentrations of HCl were compared to see how [HCl] affects rate of reaction


(f) Describe how you would alter the method to investigate how the rate of reaction varies with changing temperature.

..conduct a series of experiments in which the [HCl] and amount/length of magnesium are kept the

same but temperature is changed for each experiment

.............................................................................................................................................................

...............................................................................................................................................................

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. 2 A chloride of mercury has the formula HgClx.

The formula of the chloride can be determined experimentally. A solution of the chloride in water is reduced to mercury metal by the addition of hypophosphorous acid, H3PO2. Method • A 100 cm3 beaker is weighed empty and then with HgClx. • The solid is dissolved in distilled water. • The resulting solution is heated in a water bath. • 10 cm3 of hypophosphorous acid is added from a measuring cylinder. • The mixture is stirred and heated for a further 5 minutes. • The aqueous solution remaining after the reaction is poured off to leave droplets of

mercury in the bottom of the beaker. • The mercury is washed several times with distilled water, discarding the wash water each

time. • The beaker and mercury are dried by rinsing several times with propanone which

dissolves any remaining water drops. • The remaining propanone is evaporated by warming the beaker. • The beaker and mercury are weighed. • The experiment is repeated using different masses of HgClx.

Observations • HgClx dissolves in water to form a colourless solution. • On adding H3PO2 a white suspension forms which rapidly turns to a grey suspension f droplets of metallic mercury.

• On further heating, the suspended droplets merge to form large drops of mercury at the bottom of the beaker.


Results

A B C D E F

experiment m a s s o f beaker/g

m a s s o f beaker + mercury chloride /g

m a s s o f beaker + mercury /g

Mass of mercury chloride/g (B – A )

Mass of mercury/g (C – A )

Mass of chlorine/g (B – C )

1 54.87 55.52 55.30 0.65 0.43 0.22

2 54.64 55.88 55.59 1.24 0.95 0.29

3 56.70 58.38 57.94 1.68 1.24 0.44

4 51.03 53.34 52.53 2.31 1.50 0.81

5 55.33 58.74 57.84 3.41 2.51 0.90

6 53.05 57.20 56.10 4.15 3.05 1.10

7 53.92 58.57 57.17 4.65 3.25 1.40

8 55.26 61.09 59.57 5.83 4.31 1.52

(a) Process the results in the table to produce data that will enable you to plot a graph from which the formula of HgClx can be determined.

Record this data in the additional columns of the table. You may use some or all of the columns. Label the columns you use. For each column, include the units and an equation to show how your values are calculated. You may use the column headings A to F in the equations e.g. C – B.

(b) Present the data calculated in (a) in graphical form. Draw the line of best-fit.


(c) Indicate clearly any anomalous points on the graph that you did not use when drawing the line of best-fit. By reference to the instructions for the experiment suggest an explanation for these anomalies.

.......Mass of mercury is higher than expected. This may be due to insufficient drying/ incomplete

reaction with phosphoric acid.

...............................................................................................................................................................

Mass of chlorine/g

0 0.5 1.0 1.5 2.0 2.5 3.0 3.5 4.0 4.5

Mass of mercury/g

0

0.2

0.4

0.6

0.8

1.0

1.2

1.4

×

×

×

×

×

×

×

×


(d) Explain why the mass of HgClx used in experiment 8 is more appropriate than that

used in experiment 1.

.....low masses will result in high percentage errors when weighing...................................................

...............................................................................................................................................................

...............................................................................................................................................................

(e) Draw construction lines on the graph to derive values to enable you to calculate the value of x in HgClx.

[Ar: Hg, 201.0; Cl, 35.5] 3 g Hg is combined with 1.08 g Cl 3 mole Hg combines with 1.08 moles Cl 201 35.5 0.0149 mole Hg combines with 0.0304 moles Cl 1 mole Hg combines with 2 moles Cl

Experimental value of x is ……2…………… Formula of HgClx is … HgCl2………

(f) By considering the data you have processed and the graph you have drawn, explain

whether the experimental procedure described is suitable for the determination of the formula of HgClx.

Suitable experimental method. The graph is a straight line, (passing through the origin), with most of

the points on or near the line

...............................................................................................................................................................

(g) The mass of chlorine in HgClx can also be determined by precipitation of an insoluble

chloride. Use your knowledge of halogen chemistry to suggest a suitable reagent for this reaction.

AgNO3.......................................................................................................................

...............................................................................................................................................................


EXPERIMENT 7

PLANNING SKILL VII 1 The hydroxides of Group I metals (LiOH, NaOH, KOH, RbOH, CsOH) are highly corrosive

white solids which rapidly absorb water vapour on exposure to the atmosphere. All of these solids dissolve exothermically in water.

The enthalpy change of solution, ΔHsoln, is the energy change associated with the following reaction. M represents the Group I metal.

MOH(s) + (aq) → M+(aq) + OH–(aq)

The following diagram represents theoretical stages in the formation of aqueous MOH.

Lattice energy and hydration enthalpy are both more exothermic when ions carry a higher charge and/or ions have a smaller radius. When comparing Group I hydroxides, changes in ΔHhydration are more significant than changes in lattice energy.

(a) By considering trends in the size and charge of the ions, predict the likely trend in ΔHsoln

from LiOH to CsOH and sketch your prediction.

.. ionic size increases down the group but ionic charge remains the same. charge density decreases This results in decreasing ΔHsoln from LiOH to CsOH

[2]

lattice energy hydration enthalpy

of the ions ΔH hydration

ΔH soln

MOH(s) + (aq)

M+(g) + OH

-(g) + (aq)

M +(aq) + OH

-(aq)


2. The enthalpy change of solution, ΔHsoln, for any Group I hydroxide can be measured experimentally in the laboratory.

In experiments to compare ΔHsoln for LiOH, NaOH, KOH, RbOH, CsOH state

the independent variable .. mass/weight/moles of MOH (not metal), 2. Group I hydroxide.

The dependent variable .. 1. Temperature change, 2. Heat produced/energy change, 3. Δhsoln.

The other variable to be controlled …..1. volume/weight of water, (not amount) 2. mass of hydroxide – only allowed if “Group I hydroxide” is

given as the independent variable. .

[3]

(c) Draw a labelled diagram to show the apparatus you would use to obtain data from which ΔHsoln could be determined.

From the information given on page 2 and the apparatus you plan to use, identify two possible sources of error in the experiment and state how you would minimise the effect of each.

error 1 ...the solid hydroxide absorbs water……………………….....................................

............Store the solid in closed container………….....................................................

error 2 ....heat loss to the surrounding………………………...............................................

............cover the sides with insulation such as wool/cover the top....................................

[3] (d) Identify a health and safety risk in the experiment and explain how you would minimise

it when carrying out the experiment.

............Group I hydroxides are corrosive so use a pair of gloves/goggles......................................

...............................................................................................................................................................

........................................................................................................................................................[2]

thermometer

Polystyrene cup

water


(e) Describe the procedure you would carry out to find the enthalpy change of solution, ΔHsoln, for one of the Group I hydroxides.

Your plan should give a step-by-step description of the method, including

• how you would measure the independent variable, • how you would measure the dependent variable, • appropriate masses and volumes of reagents.

The following data may be of use in planning the detail of your experiment.

1. Measure out 100 cm3 of distilled water into a measuring cylinder and measure the initial

temperature of the water..................

2. Weigh the polystyrene cup and note its exact weight

3. Weigh out 1.0 g of LiOH into the polystyrene cup and note the exact weight of cup and LiOH

4. Add the water in the measuring cylinder into the cup with LiOH and start stirring with a

thermometer

5. Note the highest temperature reached

6. Repeat steps 1 to 5 with each of the other hydroxides using new dry polystyrene cups.

.....................................................................................................................................................[4]...

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Ar: Li, 6.9; Na, 23.0; K, 39.1; Rb, 85.5; Cs, 133.0, O, 16.0; H, 1.0 4.3 J are required to raise the temperature of 1.0 cm

3 of any solution by 1°C.


...............................................................................................................................................................

...............................................................................................................................................................

...............................................................................................................................................................

...............................................................................................................................................................

...............................................................................................................................................................

..............................................................................................................................................................

(f) The procedure is repeated for each of the Group I hydroxides. Show how you would tabulate the results for all the experiments.

Metal hydroxide

Initial Temperature/

oC

Final temperature/

oC

Mass of empty cup/g

Mass of cup + MOH/g

Change in temperature/

oC

Mass of hydroxide used/g

LiOH 25.0 39.0 13.45 14.45 14.0 1.0

NaOH

KOH

RbOH

CsOH

[1] (g) Show how you would use the results of one experiment to calculate the enthalpy

change of solution, ΔHsoln, for the reaction.

Q = m × c × T where q = energy change m = volume of water

T = change in temperature (final temp – initial temp) c = 4.3 J = 50 × 4.3 × 14 = 3010 J

Hsoln = q . Mole of hydroxide used Mole = (Mass of cup + MOH) – (mass of empty cup) Molar mass of MOH

[1]


2 Students were asked to investigate how the rate of reaction between magnesium ribbon and hydrochloric acid varied with change in concentration of the acid.

2HCl(aq) + Mg(s)→ MgCl2(aq) + H2(g)

Student 1, looking at the equation, suggested the following.

Rate of production of hydrogen gas = k[HCl]2

This student used the following apparatus to investigate the rate of production of hydrogen gas, H2.

The student used a 500 cm3 measuring cylinder to measure 100 cm3 of dilute acid into a conical flask. A 1 cm length (0.01 g) of magnesium ribbon was dropped into the acid in the flask and the stopper quickly replaced in the flask. The stop-clock was started and the volume of gas collected was measured at 0.5 minute intervals. The results of the experiment were recorded as shown in the table below.

Time / min volume of H2/ cm3 time/ min volume of H2/ cm3

0.5 15.5 5.5 80.0

1.0 25.0 6.0 82.5

1.5 34.0 6.5 85.0

2.0 43.0 7.0 87.0

2.5 51.0 7.5 87.5

3.0 59.0 8.0 91.0

3.5 65.0 8.5 92.5

4.0 69.5 9.0 93.5

4.5 74.0 9.5 94.5

5.0 75.0 10.0 95.0

(a) Plot a graph of volume of hydrogen produced against time.


[2]

(b) Identify clearly on your graph any anomalous readings and suggest a reason for these anomalous readings.

..may be due to sticking of syringe piston……………………………………………………………

..................................................................................................................................[1]…

(c) On the graph you have plotted, construct a line from which you can calculate the initial rate of reaction. Calculate the initial rate of reaction and show your working.

90 = 18.95 4.75

initial rate = .......18.95...... cm3 min–1 [2]

×

0 1 2 3 4 5 6 7 8 9 10

Time/mins

10

20

30

40

50

60

70

80


×

×

×

×

×

×

×

×

90

×

× ×

× ×

× ×

× ×

× × ×


(d) By considering the experimental method described, explain why the plotted line does not pass through 0,0.

1. stopwatch started after magnesium ribbon dropped into flask, (flask sealed and gas collected), 2. gas pushed into the syringe when inserting the bung. ……………………………………………..

[1]

(e) Identify a further source of error in the method described and suggest a change to the method to reduce this error.

1. use of 500 cm3

measuring cylinder to measure 100 cm3 of acid, use a 100 cm3 measuring cylinder

2. 0.01 g of magnesium weighed on a balance weighing to 2 decimal places. Use a larger mass 3. loss of gas before the bung is inserted

[1]Any one…..

(f) Students 2-8 carried out similar experiments with different concentrations of acid. The initial rate of reaction was calculated for each of their experiments and is shown in the table below. Enter in the table the initial rate you have calculated for Student 1.

Student

volume of acid used / cm3

Volume of water used / cm3

mass of magnesium used / g

initial rate / cm3 min–1

Relative concentration of acid

1 100 0 0.01 1.0

2 45 55 0.01 8.0 0.45

3 90 10 0.01 16.4 0.90

4 60 40 0.01 11.5 0.60

5 100 100 0.01 9.3 0.50

6 35 65 0.01 6.8 0.35

7 80 20 0.01 15.2 0.80

8 60 240 0.01 3.3 0.20

[1]

Use the additional column of the table to record calculated values for the relative concentration of the hydrochloric acid.

relative concentration of acid = volume of acid used

total volume


(g) Plot the initial rate of reaction against the relative concentration of the acid.

[3]

(h) Is the prediction made by Student 1 consistent with your graph? Explain your answer.

Prediction is wrong.

1. the rate of reaction = k[HCl], 2 reaction is directly proportional to [HCl].

3 The graph/line for a 2nd order reaction would be a curve/parabola. .........................................................................................................................................

.........................................................................................................................................

(i) Do any of the points you have plotted in (g) reduce your confidence in the conclusion that can be drawn? Justify your answer.

No. The points fall on or near the line of best fit

..............................................................................................................................[2].......

.........................................................................................................................................

.........................................................................................................................................

0 0.1 0.2. 0.3 0.4 0.5 0.6 0.7 0. 8 0. 9 1.0

Relative concentration

0

2.5

5

7.5

10

12.5

15

17.5

Initial rate /cm3min

-1

×

×

×

×

×

×

×

×


EXPERIMENT 8

PLANNING SKILL VIII 1 The equations for the complete combustion of the first five members of the alcohol homologous

series are shown below.

methanol CH3OH(l) + 1½O2(g) → CO2(g) + 2H2O(l)

ethanol H3CH2OH(l) + 3O2(g) →2CO2(g) + 3H2O(l)

propan-1-ol H3CH2CH2OH(l) + 4½O2(g) → 3CO2(g) + 4H2O(l)

butan-1-ol CH3CH2CH2CH2OH(l) + 6O2(g) → 4CO2(g) + 5H2O(l)

pentan-1-ol CH3CH2CH2CH2CH2OH(l) + 7½O2(g) → 5CO2(g) + 6H2O(l)

You are to plan an experiment to determine the enthalpy change of combustion under laboratory conditions, ΔHc, for each of the five alcohols using the apparatus and information below.

[4.2 J of heat energy raise the temperature of 1 cm3 (1 g) of water by 1 °C.]

(a) Inspection of the products of the five equations shows that from ethanol to pentan-1-ol each equation has one more mole of CO2 and one more mole of H2O produced per mole of alcohol than the equation above it.

What difference between the structures of adjacent alcohols is responsible for these additional combustion products?

..........additional –CH2- group............................................................................................[1]


(b) There is a trend in ΔHc for these five alcohols. Use your answer to (a) to sketch the trend indicated by the equations.

[1]

(c) Identify each of the following variables in the experiment to determine the enthalpy change of combustion of an alcohol using the apparatus in the diagram.

The independent variable is …mass/moles of alcohol or type of alcohol The dependent variable is …temperature rise or change Two variables to be controlled in this experiment are

1 ......mass of alcohol/water..................................................................

2 ....same metal can / position of metal can above the flame .................................[3]


(d) You are to plan a step-by-step method to determine the enthalpy change of combustion, ΔHc, for each of the alcohols in turn. You should use the apparatus shown in the diagram on page 2 and any other apparatus required that is normally found in a school or college laboratory.

You need only explain how the method will be used for ethanol. Your method should cover the following points. • the precision of the apparatus to be used • the way in which the independent variable is varied between experiments • measurement of the dependent variable • appropriate quantities to be used • the raw data to be collected • how this data is to be processed for ethanol

[4.2 J of heat energy raise the temperature of 1 cm3 (1 g) of water by 1 °C.] Method ..100 cm3 of water was placed in a metal can and the temperature of the water was measured using a thermometer capable of measuring 1oC......... A spirit lamp was filled with ethanol and the mass of spirit lamp + alcohol was measured using a balance capable of measuring 0.01 g....................... The spirit lamp was ignited and placed below the water

The water was heated for 10 minutes with constant stirring.............................

After 10 mins the spirit lamp was extinguished

The highest temperature reached by the water was recorded...........................

The spirit lamp with the remaining alcohol was weighed and the mass recorded

.............................................................................. [4]

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................................................................................................................................................................ Processing of raw data [Mr: CH3CH2OH, 46]

q = m × c × T where q = energy change m = volume of water

T = change in temperature (final temp – initial temp) c = 4.2 J

Hc = q . Mole of ethanol burned Mole = (initial mass of lamp + alcohol ) – (final mass of lamp + alcohol) 46 [2]

(e) Prepare a single table to show, for each alcohol, all of the measurements obtained directly from the experiment and the information processed from these data.

Your table should show the correct units for each of the headed columns.

Alcohol Initial Mass of spirit lamp + alcohol /g

Final Mass of spirit lamp + alcohol /g

Initial temperature /oC

Final temperature/ oC

Mass of alcohol burned /g

Change in temperature/

oC

Methanol

ethanol

Propan-1-ol

Butan-1-ol

Pentan-1-ol


(f) The apparatus shown in the diagram on page 57 and used in the experiment you have described is inefficient as much of the energy released from the burning alcohol is lost to the atmosphere. You are to plan two simple changes to the apparatus shown on page 57 that will reduce the loss of heat energy to the atmosphere. Draw a diagram of your improved apparatus, indicating and explaining how these changes will reduce the loss of heat energy.

improvement 1 .insulate the sides of the beaker with glass wool. This prevents heat loss to the

surrounding through conduction

...............................................................................................................................................................

...............................................................................................................................................................

...............................................................................................................................................................

...............................................................................................................................................................

improvement 2 .Use a wind(draught) shield or raise the spirit lamp or lower the height of the

stand ( or cover the top ).This prevents loss of heat to the surrounding through convection

...............................................................................................................................................................

...............................................................................................................................................................

...............................................................................................................................................................

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............................................................................................................................................................[2] [total : 15]

Glass wool

Wind shield


2 Basic copper(II) carbonate is a mixture of copper(II) carbonate and copper(II) hydroxide with varying amounts of water of crystallisation. The formula may be represented as CuCO3.Cu(OH)2.xH2O. A group of students conduct an experiment to determine the relative formula mass and the actual value of x in the formula of the basic copper(II) carbonate, CuCO3.Cu(OH)2.xH2O. This basic carbonate decomposes on heating. CuCO3.Cu(OH)2.xH2O (s) → 2CuO(s) + CO2(g) + (x+1)H2O(g) Each student is provided with a sample of the solid basic carbonate and the following apparatus. • a boiling-tube • a test-tube holder • a chemical balance, reading to 2 decimal places • a Bunsen burner

Method • The empty boiling-tube is weighed.

• The sample of basic copper(II) carbonate is tipped into the boiling-tube which is reweighed.

• The carbonate is heated, gently at first, as carbon dioxide and water vapour are given off. The carbonate is then heated strongly.

• Heating is continued until all of the green colour of the initial solid has turned to the black colour of copper(II) oxide.

• The apparatus is cooled and the mass of the tube and residue is measured.

The results of the experiment are recorded below.


A B C D E F

student mass of empty boiling-tube /g

mass of boiling-tube and basic carbonate before heating /g

mass of boiling-tube and residue after heating /g

Mass of basic carbonate used / g

Mass of residue (CuO) /g

1 10.32 11.19 10.92 0.87 0.60

2 10.05 11.56 11.09 1.51 1.04

3 10.11 12.33 11.24 2.22 1.13

4 9.94 12.39 11.62 2.45 1.68

5 9.99 13.73 12.56 3.74 2.57

6 10.23 14.68 13.28 4.45 3.05

7 10.01 15.19 13.56 5.18 3.55

8 9.87 15.80 13.97 5.93 4.10

9 9.96 16.62 14.66 6.66 4.70

10 9.84 17.83 15.46 7.99 5.62

[2]

(a) Use the additional columns of the table to record the calculated values that will enable you to determine the relative formula mass, Mr , of basic copper(II) carbonate and the value of x in CuCO3.Cu(OH)2.xH2O. You may use some or all of the columns. Label the columns you use, including the units and an equation to show how the value is calculated. You may use the column headings A to F in the equations e.g. C – B. Make certain that your processed values show the precision of the apparatus used.

(b) Present the values calculated in (a) in graphical form. Draw the line of best-fit.


[4]

Mass of CuO/g

0 1.0 2.0 3.0 4.0 5.0 6.0 7.0 8.0 9.0

Mass of basic carbonate / g

0

0.5

1.0

4

1.5

2.0

2.5

3.0

3.5

×

×

×

×

×

×

×

× 4.0

4.5

5.0

5.5

6.0

×

×


(c) Give the co-ordinates of the most significant anomalous point. By reference to the instructions for the experiment suggest an explanation for this anomaly. co-ordinates .......2.22, 1.13 ..................................................................................... explanation ...mass of CuO lower than expected. It could be due to solid being blown

out of tube due to strong heating ...........................................................................

.................................................................................................................................[2]

(d) The students note that points plotted when using higher masses of the basic carbonate lie to one side of the line of best-fit.

Suggest an explanation for this observation.

...Mass of CuO is higher than expected. This may be due to incomplete decomposition

of basic carbonate ................................................................................................

..................................................................................................................................[1]

(e) How would you modify the experiment to overcome the problem described in (d)?

...heat the basic carbonate longer until constant mass is obtained ..................................

.........................................................................................................................................

....................................................................................................................................[1]

0.� Draw construction lines on the graph to derive values to enable you to calculate the relative formula mass of the basic copper(II) carbonate and the value of x in CuCO3.Cu(OH)2.xH2O.

Values read from graph, including units

x-axis …...........4.0…..................... y-axis …...........2.75….................................

Calculation of the relative formula mass, Mr [Ar: H, 1.0; C, 12.0; O, 16.0; Cu, 63.5]

nCuO = 2.75 = 0.0346 79.5

n CuCO3.Cu(OH)2.xH2O = 0.0346 2 = 0.0173 Mr CuCO3.Cu(OH)2.xH2O = 4.0 g = 231 g mol-1

0.0173 mol


(f) continued Calculation of x in the formula

Mass of water = 231 – 2(63.5) – 12 – 5(16) – 2 = 10 x = 10/18 = 0.56

[2]

(g) By considering the data you have processed and the graph you have drawn, explain if the experimental procedure described is suitable for the determination of the relative formula mass of the basic carbonate.

..it is suitable. Most of the points fall on or near the line of best fit……… .....................

.........................................................................................................................................

.........................................................................................................................................

.........................................................................................................................................

...................................................................................................................................[1]

(h) What other measurements could be made during the course of the experiment to provide alternative data to determine the relative formula mass and the value of x ?

.....estimate the volume of CO2 produced, after condensing the water……..................

.........................................................................................................................................

.........................................................................................................................................

....................................................................................................................................[1]

(i) Explain why the precision of the experiment would improve if students used a chemical

balance weighing to 3 decimal places.

...Since the masses used are small , it will decrease the percentage error..........................

.........................................................................................................................................

.........................................................................................................................................

....................................................................................................................................[1]

[Total:15]

EXPERIMENT 9

PLANNING SKILL IX


1 The solubility of cerium(IV) sulphate, at a particular temperature, is defined as: the mass of cerium(IV) sulphate that will dissolve in and just saturate 100 g of solvent at that temperature.

Cerium(IV) sulphate crystals also dissolve in dilute sulphuric acid, H2SO4, a corrosive aqueous solution. You are to plan an experiment to investigate how the solubility of cerium(IV) sulphate crystals in dilute sulphuric acid depends on the concentration of the acid. (a) By considering the ions present in a solution of cerium(IV) sulphate in sulphuric acid

predict and explain how the solubility of the cerium(IV) sulphate will be affected by the concentration of the acid.

Prediction .....solubility decreases..........................................................

........................................................................................................................................

Explanation ....due to common ion effect of sulphate ions from sulphuric acid ...............

........................................................................................................................................

........................................................................................................................................

Display your prediction in the form of a graph.

solubility of cerium(IV) sulphate

concentration of sulphuric acid [2]

(b) In the experiment you are about to plan, identify the following.

(i) the independent variable ….....concentration of sulphuric acid….......................

(ii) the dependent variable …......solubility of cerium sulphate solid ….....................

Cerium(IV) sulphate will dissolve in water. A saturated solution is one in which no more solid can dissolve at a particular temperature. In a saturated solution in contact with undissolved solid, the following equilibrium is established.

Ce(SO4)2(s) + aq → Ce4+(aq) + 2SO42–(aq)

Cerium(IV) sulphate is unusual. Its solubility decreases as the temperature of the solution

increases


(ii) another variable to be controlled ….......temperature…..................

[3]

(c) Design a two-part laboratory experiment to investigate your prediction in (a).

Part 1 – solubility of cerium(IV) sulphate in water

In addition to the standard apparatus present in a laboratory you are provided with the following materials.

60 cm3 distilled water solid cerium(IV) sulphate, Ce(SO4)2

Give a step-by-step description of how you would

(i) prepare a saturated solution using all of the 60 cm3 of distilled water,

(ii) control the variable given in (b) (iii),

(iii) separate the saturated solution from undissolved solid,

(iv) obtain the mass of cerium(IV) sulphate and mass of water in the saturated solution,

(v) calculate the solubility of cerium(IV) sulphate from the experimental results.

....................................................................................................................................................

1 Using a measuring cylinder, 60 cm3 of distilled water was measured into a 250 cm3 beaker.

2 The beaker with water was placed in a thermostated water bath maintained at 25oC

3 CeSO4 solid was added to the water with continuous stirring until no more solid dissolved and there was some solid at the bottom of the beaker

4 10 cm3 of the saturated solution was then placed in a pre-weighed beaker and weighed.

5 The solution was heated until all the water evaporated 6 The beaker + residue was then weighed

7 The residue was then repeatedly heated and weighed until a constant mass was obtained

8 The mass of CeSO4 dissolved was obtained : (Mass of beaker + solid residue) – (mass of beaker)

9 The mass of water(solvent) = (mass of beaker + saturated solution) – (Mass of beaker + solid residue)

10 Solubility = Mass of CeSO4 dissolved × 100

Mass of water

[4]

or

.........................................................................................................................................................


1 Using a measuring cylinder, 60 cm3 of distilled water was measured into a 250 cm3 beaker.

2 Weigh the beaker + water and record the accurate weight.

3 The beaker with water was placed in a thermostated water bath maintained at 25oC

4 CeSO4 solid was added to the water with continuous stirring until no more solid dissolved and there was some solid at the bottom of the beaker

5 The beaker + content was weighed and the accurate mass recorded

6 The solid was then filtered off and the solid was weighed

7 The mass of CeSO4 dissolved was estimated : [( mass of beaker + content) – (mass of beaker + water) – (mass of solid)

8 Solubility = Mass of CeSO4 dissolved × 100

60

Part 2 – solubility of cerium(IV) sulphate in sulphuric acid In addition to the standard apparatus present in a laboratory you are provided with the following materials.

5 mol dm–3 sulphuric acid distilled water solid cerium(IV) sulphate

Give a step-by-step description of how you would prepare a range of solutions of sulphuric acid of different concentrations. These solutions could then be used to investigate the solubility of cerium(IV) sulphate in different concentrations of sulphuric acid. See section (e). Your plan should include details of the following.

(i) the number and concentrations of the solutions to be prepared

(ii) the volumes of acid and water used to prepare the solutions

(iii) the apparatus and method you would use in their preparation


The water and 5 M sulphuric acid solutions were placed in two separate burettes

100 cm3 each of 1.0, 2.0, 3.0 , 4.0 and 5.0 mol dm-3 H2SO4 solution were prepared by mixing 5M H2SO4 and water in a beaker as shown below

[H2SO4]/ mol dm-3 Volume of 5 M H2SO4 /cm3 Volume of water / cm3

1.0 20.0 80.0

2.0 40.0 60.0

3.0 60.0 40.0

4.0 80.0 20.0

5.0 100.0 0.0

..........................................................................................................................................................[4]

….............................................................................................................................................................

….............................................................................................................................................................

…............................................................................................................................................................

…............................................................................................................................................................

…............................................................................................................................................................

….............................................................................................................................................................

…............................................................................................................................................................

…............................................................................................................................................................

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…............................................................................................................................................................

…...........................................................................................................................................................

….............................................................................................................................................................

…............................................................................................................................................................

…...........................................................................................................................................................

….............................................................................................................................................................

…............................................................................................................................................................


…...........................................................................................................................................................

….............................................................................................................................................................

…............................................................................................................................................................

…...........................................................................................................................................................

….............................................................................................................................................................

(c) State a risk that must be taken into consideration when planning the experiment. …...sulphuric acid is corrosive…………………………..........................................................................

…...........................................................................................................................................................

…............................................................................................................................................................. [1]

(d) Draw up a table with appropriate headings to show the data you would record when investigating the solubility of cerium(IV) sulphate in the solutions of sulphuric acid planned in C, Part 2.

Method 1

[H2SO4] Mass of empty beaker /g (A)

Mass of beaker + saturated solution/g (B)

Mass of beaker + dried solid/g (C)

Mass of solid dissolved /g C – A = D

Mass of H2O/g B – C = E

Solubility /g per 100 cm3

D g x 100 E cm3

1.0

2.0

3.0

4.0

5.0

Method 2

[H2SO4] Mass of beaker + water/ g (A)

Mass of beaker + water + solid before filtering/g (B)

Mass of solid filtered/g (C)

Mass of solid dissolved /g B – A –C = D

Mass of H2O/g

Solubility D g 100 cm3

1.0 100

2.0 100

3.0 100

4.0 100

5.0 100

[1]


[total 15]

2 A student reads about the electrolysis of aqueous copper(II) sulphate and makes the following notes.

The student ealizes that a value for the Faraday constant can be determined experimentally using the following apparatus.

Method • The cathode is cleaned and weighed before being placed in the copper(II) sulphate

solution.

• The circuit is completed and the current set at 0.3 A by adjusting the variable resistor.

• The current is maintained at 0.3 A for exactly 40 minutes at which point the circuit is broken.

• The cathode is removed from the solution and carefully washed with distilled water to remove any copper(II) sulphate solution.

• Distilled water is removed from the cathode by rinsing it with propanone in which the water dissolves.

• The cathode is finally dried by allowing the propanone to evaporate from its surface.

• The cathode is reweighed and placed back in the solution.

• A constant current of 0.3 A is passed for a further 40 minutes when the rinsing, drying and weighing are repeated.

• This procedure is repeated a further 8 times.

• reaction at the cathode Cu2+

(aq) + 2e– → Cu(s)

• reaction at the anode Cu(s) →Cu2+

(aq) + 2e–

• 1 mol of Cu(s) is deposited on the cathode by 2 mol of electrons.

• The Faraday constant is the charge in coulombs, C, on 1 mol of electrons.

• The Faraday constant = 9.5 × 104 C mol

–1.



A B C D E F

Time /minutes

mass of cathode /g

Copper deposited/g

Charge/C (A×60×0.3)

0 115.74 0.00 0.00

40 115.97 0.23 720

80 116.22 0.48 1440

120 116.46 0.72 2160

160 116.70 0.96 2880

200 116.94 1.20 3600

240 117.19 1.45 4320

280 117.49 1.75 5040

320 117.67 1.93 5760

360 117.92 2.18 6480

400 118.14 2.40 7200

[2] (a) Use the additional columns of the table to record the charge passed and the mass of

copper deposited on the cathode. [ charge in coulombs = time in seconds × current in amps; 1 minute = 60 seconds ] You may use some or all of the columns. Label the columns you use, including the units

and an equation to show how the value is calculated. (b) Present the data calculated in (a) in graphical form. Draw the line of best-fit.


[4]

2.6

×

×

×

1.6

1.8

2.0

2.2

2.4

Mass of Cu /g

0 1000 2000 3000 4000 5000 6000 7000 8000

Charge/ C

0

0.2

0.4

0.6

0.8

1.0

1.2

1.4

×

×

×

×

×

×

×

×


(c) Indicate clearly one anomalous point on the graph that you did not use when drawing the line of best-fit. By reference to the instructions for the experiment suggest an explanation for the anomaly.

.......Mass of copper is higher than expected. The copper on cathode was not completely

dried ..

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(d) The balance used by the student weighed to 2 decimal places. By reference to the results of the experiment explain why it would have been more appropriate to use an analytical balance weighing to 4 decimal places.

.. % error is decreased since small masses are involved..................

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1 Draw construction lines on the graph to derive relevant values and use them to calculate a

numerical value for the Faraday constant. Reaction at the cathode Cu2+(aq) + 2e– → Cu(s) [Ar: Cu; 63.5]

3650 C deposits 1.2 g of Cu 3650 C deposits 1.2 = 0.0189 moles of Cu 63.5 1 mole Cu requires 3650 =193122 C = 2 Faradays 0.0189 1 faraday = 193122 = 96560 C 2

The value of the Faraday constant determined in this experiment is

....... 96560................. C mol–1. [2]

(f) By considering the data you have processed and the graph you have drawn, decide if the experimental procedure described is suitable for the determination of the Faraday constant. Explain your reasoning.


..It is suitable. The line of best fit is linear and passes through the origin. Most of the points fall on

or near the line of best fit...............................................

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(g) What other measurements could be made during the course of the experiment to

provide alternative data to confirm the determined value of the Faraday constant?

...Weigh the loss in mass of anode..................................................

.............................................................................................................................................[1]..

(h) Another student, performing the same experiment, plotted the mass of copper

deposited against time, and obtained the results below.

Making reference to the experimental method, suggest an explanation for the shape of this graph. ....there was a decrease in current after some time..................................

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EXPERIMENT 10

PLANNING SKILL X


1 You are to plan an investigation into the thermal decomposition of caesium nitrate, CsNO3.

You may make use of some or all of the following data when planning your investigation.

Group I element cation ionic radius/ nm

lithium Li+ 0.060

sodium Na+ 0.095

potassium K+ 0.133

rubidium Rb+ 0.148

caesium Cs+ 0.176

nitrogen dioxide gas

oxygen gas

NO2 brown in colour soluble in water poisonous

O2 colourless almost insoluble in water powerful oxidant

Equations for the thermal decomposition of lithium nitrate and sodium nitrate are given below. 4LiNO3(s) → 2Li2O(s) + 4NO2(g) + O2(g) 2NaNO3(s) → 2NaNO2(s) + O2(g)

1 mol of any gas occupies a volume of approximately 24 dm3 at room temperature

and atmospheric pressure. Ar: Cs, 133; N, 14.0; O, 16.0


(a) Predict which of the equations below will represent the thermal decomposition of caesium nitrate. Place a tick against the equation of your choice.

4CsNO3(s) → 2Cs2O(s) + 4NO2(g) + O2(g)

2CsNO3(s) → 2CsNO2(s) + O2(g)

Use data provided on page 76 to explain your prediction. Cs+

has a larger ionic radius than Li+ or Na+. Therefore it has less polarising power than Na+ or Li+. .........................................................................................................................................

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(iii) You are to plan an experiment in which • caesium nitrate is heated, • gas is collected, • the volume of gas collected is measured, • the experimental results are used in a calculation to confirm or reject your prediction. (iv) Identify the independent variable in the experiment.

Mass of caesium nitrate used…................................................................................

(v) Identify the dependent variable in the experiment. Volume of gas collected….........................................................................

[2] (c) Draw a diagram of the apparatus you would use in this experiment.

Your apparatus should use only standard items found in a school or college laboratory. Show clearly how the solid will be heated, the gas collected and its volume measured. Label each piece of apparatus used, indicating its size or capacity, e.g. 250 cm3 beaker.


50

40

30

20

10

0

Assuming that either equation in (a) might be correct, which gas or gases would you expect to collect in your apparatus. Explain your answer.

O2. NO2 is soluble in water .............................................................................

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(d) Calculate the volume of gas you would expect to collect in your apparatus if 1 mol of

caesium nitrate completely decomposed according to your predicted equation in (a). 1 mole CsNO3 produces ½ mole O2 At R.t.p. ½ mole of gas occupies 24 × ½ = 12 dm3

[1]

(e) Use your answer to (d) and the size of the apparatus selected in (c) to calculate the maximum mass of CsNO3 that can be used in your experiment.

50 × 10-3dm3 × 195 = 0.8125 g 12

[1]

(f) Outline, in a series of numbered steps, the method to be used in the experiment.

50 cm3 burette

500 cm3 beaker 250 cm

3

conical flask

CsNO3

Borosilicate

test-tube


Make certain that the steps you describe are in the correct order. You need not explain how the apparatus is assembled. Indicate clearly how you will know when decomposition is complete.

1 0.8 g of CsNO3 was weighed into a boiling tube and connected to the transfer tube ...............

2 The CsNO3 was then heated gently and later strongly and the gas was collected in the inverted burette

3 The heating was conducted until no more gas was collected

4 The delivery tube was removed from the burette and the bunsen flame was stopped

5 The volume of gas collected was then estimated by subtracting the final volume from initial volume of water in the burette (since burette is inverted)

6 If (0.8 /0.8125 × 50 = 49.2) cm3 of gas is collected then the predicted (2nd ) equation is correct.

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g) What should be done when decomposition is complete to ensure that the volume of gas

measured in the apparatus is the “correct” volume?


…Cool the gas to room temperature……………………………….........................................................

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(h) Identify a risk present in the method you have described.

...1. NO2.is poisonous O2 is easily combustible …………………………….........

.....2. Potential suck back if gas is collected over water ……………………... ...................................

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Describe how you would minimise this risk.

....1 conduct the experiment in a fume chamber

2 Remove the delivery tube from water before stop heating .................................................

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2 .

When a solution of sodium iodide, NaI, saturated at 100 °C is cooled, crystals of NaI form and sink to the bottom of the solution. As the solution continues to cool a temperature is reached below which crystals of NaI.2H2O are deposited. The temperature when this change takes place is called the transition temperature. Above and below this temperature the way in which the solubility changes with temperature is noticeably different. A group of students carried out an experiment to determine the transition temperature.

Solubility is defined as the amount of a substance that will dissolve in and just saturate 100 g of solvent at a particular temperature. solubility = mass of solid in a saturated solution × 100

mass of water in the saturated solution When a solution is saturated, dissolved solid is in equilibrium with undissolved solid. To make a saturated solution add the solid to water at a particular temperature until no more dissolves. Then add more solid and leave the mixture in a thermostatically controlled water bath to establish equilibrium This question uses sodium iodide as the solid to be dissolved and water as the solvent


The instructions given for the experiment were as follows. • Prepare a saturated solution at a temperature between 20 °C and 100 °C. • Record the temperature of the saturated solution. • Weigh an empty evaporating basin. • Transfer some of the saturated solution, but no solid, into the weighed evaporating

basin. • Weigh the evaporating basin and solution. • Evaporate the water from the solution by placing the evaporating basin on top of a

beaker of boiling water. • When all of the water in the solution has evaporated, cool and reweigh the evaporating

basin. • Repeat the heating, cooling and weighing until a constant mass is obtained.


A B C D E F G H

temperature at which the saturated solution as prepared / °C

mass of evaporating basin / g

mass of basin + solution / g

final constant mass of basin + solid / g

Mass of solid dissolved/g (D-B)

Volume of water used/ cm3 (C-D)

Solubility E/F x 100

20 23.7 145.1 101.6 77.9 43.5 179

30 31.8 182.0 130.2 98.4 51.8 190

40 33.4 172.5 126.9 93.4 45.6 205

45 25.9 214.3 154.3 128.4 60.0 214

50 31.9 229.1 166.7 134.8 62.4 216

55 27.6 217.0 160.8 133.2 56.2 237

60 33.3 242.9 184.2 150.9 58.7 257

65 31.6 298.7 228.6 197 70.1 281

70 28.5 225.7 175.4

146.9 50.3 292

75 29.1 203.6 159.2

130.1 44.4 293

80

30.0 220.4 172.2

142.2 48.2 295

85 27.8 242.4 188.2

160.4 54.2 296

90 36.6 226.0 178.4

141.8 47.6 298

95 31.4 247.0 193.1

161.7 53.9 300

100 32.9 225.9 177.9

145 48 302


(a) Process the results in the table to produce values that will enable you to plot a graph to show the solubility of sodium iodide in water at different temperatures. Record these values in the additional columns of the table. You may use some or all of the columns. Label the columns you use. For each column you use include the units and an expression to show how your values are calculated. You may use the column headings A to H in the expressions e.g. C – B. [3]

(b) Present the values calculated in (a) in graphical form. Draw two separate lines and extrapolate them to a point of intersection. The line at higher temperatures represents the solubility of NaI. The line at lower temperatures represents the solubility of NaI.2H2O. Do not start either scale at zero.


[4]

solubility

×

10 20 30 40 50 60 70 80 90 100

Temperature/ oC

170

180

190

200

210

220

240

260

×

×

×

× ×

×

×

270

280

290

300

310

×

× × × ×

× ×


(c) Read from the graph the temperature where the two lines intersect. This is the transition temperature.

The transition temperature is …66…… °C.

[1]

(d) Circle, on the graph, any point(s) you consider to be anomalous. For any point circled on the graph suggest an error in the conduct of the experiment that might have led to the anomalous result.

..at T= 45 the mass dissolved is larger than expected. Some of the water has not been evaporated ......

At T = 30, 50, 60 the mass dissolved is lower than expected. Loss of solid during heating. Loss of

water of hydration from NaI. H2 O

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(e) Suggest additional measurements that could be made to enable a more precise value

of the transition temperature to be determined.

........more measurements around the transition temperature……………….........................................

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(f) Describe the difference in the variation of solubility with temperature above the

transition temperature,

......solubility increases gradually …………………………………….....................................................

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below the transition temperature.

Solubility increases rapidly................................................................................................................

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The solubility curve represents equilibrium conditions between solid sodium iodide and dissolved sodium iodide. The position of equilibrium can be influenced by temperature change and whether a change is exothermic or endothermic. From the shape of your graph, comment on the likely enthalpy change for solid sodium iodide dissolving under equilibrium conditions. Explain your answer.

.Endothermic, because solubility increases with temperature.........................................................

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Revised: December, 2009

Documents

CHM1313practical Ans