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PART A

Practical work

Experiments 3Solution preparation quantities 4

Experiments for each chapter 6

Results, answers and activities 87Each chapter contains:

• expected results for each experiment

• suggested answers to experiment questions

• additional learning activities

• internet resources.

Experiments

Solution preparation quantities 4

EXPERIMENT 1.1 : Iron content of steel wool 6

EXPERIMENT 2.1 : Sulfate in lawn food 7

EXPERIMENT 3.1 : Ammonia content of cloudy ammonia — a problem-solving exercise 9

EXPERIMENT 3.2 : Nitrogen content of lawn fertiliser 10

EXPERIMENT 3.3 : Standardisation of hydrochloric acid 11

EXPERIMENT 3.4 : Indicators — a problem-solving exercise 13

EXPERIMENT 4.1 : Separation of food dyes using paper chromatography 16

EXPERIMENT 4.2 : Separating mixtures using column chromatography 18

EXPERIMENT 4.3 : A model GC detector 19

EXPERIMENT 5.1 : Flame colours 20

EXPERIMENT 5.2 : Zinc content of cornfl akes using atomic absorption spectroscopy 21

EXPERIMENT 5.3 : Iron content of wine using UV–visible spectroscopy 22

EXPERIMENT 5.4 : Phosphates in detergents 23

EXPERIMENT 6.1 : Constructing models of alkanes and alkenes 25

EXPERIMENT 6.2 : Constructing models of organic compounds 26

EXPERIMENT 6.3 : Identifying alkanes and alkenes 27

EXPERIMENT 7.1 : Carbohydrate models 28

EXPERIMENT 7.2 : Modelling the structure and denaturation of proteins 29

EXPERIMENT 7.3 : A stomach enzyme 30

EXPERIMENT 7.4 : The action of bile 31

EXPERIMENT 7.5 : Testing for nitrates in soils 32

EXPERIMENT 8.1 : Electrophoresis simulations 33

EXPERIMENT 8.2 : Preparation of aspirin 35

EXPERIMENT 9.1 : Reaction rates 37

EXPERIMENT 9.2 : Investigating changes to the position of an equilibrium 39

EXPERIMENT 9.3 : Rate of hydrogen production — a problem-solving exercise 41

EXPERIMENT 9.4 : Temperature and the equilibrium constant 42

EXPERIMENT 9.5 : Modelling an equilibrium 43

EXPERIMENT 10.1 : The acidity constant of acetic acid 45

EXPERIMENT 10.2 : Buffers 47

EXPERIMENT 11.1 : Catalytic oxidation of ammonia 48

EXPERIMENT 11.2 : Cracking of paraffi n oil 49

EXPERIMENT 11.3 : Fractional distillation of ethanol/water mixture 51

EXPERIMENT 11.4 : Catalytic dehydration (cracking) of ethanol 53

EXPERIMENT 11.5 : Constructing models of esters 54

EXPERIMENT 11.6 : An esterifi cation experiment 55

EXPERIMENT 13.1 : Investigating heat changes in reactions 57

EXPERIMENT 13.2 : Calorimetry: Determining the heat of reaction of a simple redox reaction 60

EXPERIMENT 13.3 : The energy content of a peanut 62

EXPERIMENT 14.1 : Investigating the Daniell cell 64

EXPERIMENT 14.2 : Galvanic cells and redox potentials 66

EXPERIMENT 15.1 : Looking at a dry cell 68

EXPERIMENT 15.2 : Lead storage batteries 70

EXPERIMENT 15.3 : Investigating a hydrogen–oxygen fuel cell 71

EXPERIMENT 16.1 : Electrolysis of aqueous solutions of electrolytes 73

EXPERIMENT 16.2 : Factors affecting electrolysis 75

EXPERIMENT 17.1 : Electroplating 76

EXPERIMENT 17.2 : Anodising aluminium 78

EXTENDED INVESTIGATIONS

Vitamins and minerals 80

An investigation of iron in a liquid iron supplement 81

Analysis of vitamin C 82

Gravimetric determination of salt in salt tablets 83

Investigating iron content in iron tablets using atomic absorption spectroscopy 85

Chemistry 2: Practical work

studyon Chemistry 2 TEACHER SUPPORT KIT © John Wiley & Sons Australia, Ltd 2008

This sheet may be photocopied for non-commercial classroom use.

4

Solution preparation quantities

Chemical Formula Make up to 1 L with distilled water

acetic acid — 0.1 mol L–1 CH3COOH 6 mL

ammonia (0.88) — 5.0 mol L–1 NH4OH 280 mL

ammonium chloride — saturated NH4Cl Add NH4Cl to water until no longer soluble.

ammonium peroxydisulfate

— 0.1 mol L–1

(NH4)2S2O8 22.8 g

barium chloride — 0.5 mol L–1 BaCl2.2H2O 122 g

bromine solution Br 1 mL diluted in 300 mL water

copper sulfate — 0.01 mol L–1

0.2 mol L–1

0.5 mol L–1

1.0 mol L–1

CuSO4.5H2O 2.5 g

50 g

125 g

250 g

disodium hydrogen orthophosphate

— 0.1 mol L–1

Na2HPO4 14.2 g

ferric nitrate — 0.1 mol L–1 Fe(NO3)3.9H2O 40.4 g + 50 mL 2 mol L–1 HNO3

ferrous sulfate — 1.0 mol L–1 FeSO4.7H2O Dissolve 278 g in water containing 20 mL

conc. H2SO4 and make up to 1 L.

hydrochloric acid (12 mol L–1 36%)

— 0.1 mol L–1

1.0 mol L–1

5.0 mol L–1

HCl

8.75 mL

87.5 mL

437.5 mL

lead nitrate — 1.0 mol L–1 Pb(NO3)2 331 g

nickel(II) sulfate — 0.2 mol L–1 NiSO4.6H2O 52.6 g

nitric acid (16 mol L–1 70%)

— 0.5 mol L–1

3.0 mol L–1

HNO3

32 mL

192 mL

potassium chloride — 1.0 mol L–1 KCl 74.6 g

potassium hydroxide — 0.5 mol L–1

1.0 mol L–1

KOH 28.0 g

56.1 g

potassium iodide — 0.1 mol L–1 KI 16.6 g

potassium permanganate

— 0.02 mol L–1

0.005 mol L–1

KMnO4 3.2 g

Dilute 250 mL 0.02 mol L–1 potassium

permanganate to 1 L

potassium thiocyanate — 0.1 mol L–1 KSCN 9.7 g

silver nitrate — 0.1 mol L–1 AgNO3 17.0 g

sodium bromide — 0.5 mol L–1 NaBr 51.5 g

sodium carbonate — 0.1 mol L–1 Na2CO3 11.0 g

sodium fluoride — 0.1 mol L–1 NaF 4.2 g

Chemistry 2: Practical work

© John Wiley & Sons Australia, Ltd 2008 studyon Chemistry 2 TEACHER SUPPORT KITThis sheet may be photocopied for non-commercial classroom use.

5

Chemical Formula Make up to 1 L with distilled water

sodium hydrogen carbonate

— 0.1 mol L–1

NaHCO3 8.4 g

disodium hydrogen orthophosphate

— 0.1 mol L–1

Na2HPO4 14.2 g

sodium hydroxide — 0.1 mol L–1

1.0 mol L–1

NaOH 4.0 g

40 g

sodium iodide — 0.5 mol L–1 NaI 74.9 g

sodium phosphate — 0.1 mol L–1 NaH2PO4.2H2O 15.6 g

sodium thiosulfate — 0.0025 mol L–1 Na2S2O3.5H2O 0.6 g

sulfuric acid (18 mol L–1 98%)

— 1.0 mol L–1

2.0 mol L–1

H2SO4

55 mL

110 mL

zinc chloride — 1 mol L–1 ZnCl2 136 g

zinc sulfate — 0.1 mol L–1

1.0 mol L–1

ZnSO4.7H2O 28.8 g

288 g



6

Chemistry 2: Experiments EXPERIMENT 1.1

Iron content of steel wool

In this redox titration, the iron in steel wool is fi rst reacted with sulfuric acid

to change it into iron(II) ions. These ions are then titrated with a standardised

solution of potassium permanganate.

The equation for the reaction is:

5Fe2+(aq) + MnO4–(aq) + 8H+(aq) 5Fe3+(aq) + Mn2+(aq) + 4H2O(l)

AIMTo determine the percentage (by mass) of iron in steel wool

APPARATUSburette with stand

measuring cylinder

electronic balance

standardised 0.02 mol L–1 potassium permanganate

1.0 mol L–1 sulfuric acid

3 × 100 mL conical fl asks

small watchglass

heating equipment

wash bottle

METHOD

1. Add the potassium permanganate solution to the burette.

2. Accurately weigh out about 0.1 g of steel wool and place into a 100 mL

fl ask.

3. Add about 20 mL of 1.0 mol L–1 sulfuric acid. After placing a watchglass

over the top of the fl ask, heat the contents until the steel wool has completely

reacted. Note: Some unreacted impurities may be evident, this is quite

normal.

4. Use a little water from a wash bottle to wash any splashed solution from

the sides of the fl ask back down into the bulk of the liquid.

5. Add the potassium permanganate solution from the burette until a faint pink

colour persists for at least 30 seconds. Note the volume required. Note: Due

to the intensity of colour in the potassium permanganate solution, you may

fi nd it easier to make all burette readings from the top of the meniscus.

6. Repeat for two more samples of steel wool.

7. For each sample, calculate the percentage of iron that was present in the

steel wool.

QUESTIONS

1. Why was it not necessary to add an indicator in this experiment?

2. Why was the volume of acid used in step 3 only measured

approximately?

3. Explain why reading the burette volumes to the top of the meniscus (as

in this experiment) would not cause an error, yet reading the volume of a

pipette to the top of the meniscus would cause an error.

EXTENSIONDesign an experiment to see if different brands of steel wool rust at different

rates. Rust contains Fe3+ ions, which should not interfere with the method of

titration described above.

Take care when handling the


solution. If spilt, it can stain skin

and clothes. Special care should

be exercised with the hot sulfuric

acid solution in step 3 as it is

corrosive.



solution. If spilt, it can stain skin

and clothes. Special care should

be exercised with the hot sulfuric

acid solution in step 3 as it is

corrosive.


7

Chemistry 2: Experiments

Sulfate in lawn food

Sulfur in the form of sulfate is an important plant nutrient and is therefore

found in many fertiliser preparations. As sulfates easily form precipitates with

suitable metal cations, a gravimetric method based on this feature is a common

way to analyse for sulfate. The metal ion commonly used to form such a

precipitate is barium, owing to the very low solubility of barium sulfate.

The equation for this reaction is:

Ba2+(aq) + SO42–(aq) BaSO4(s)

AIMTo determine the amount of sulfur (as sulfate) in a sample of lawn food

APPARATUSelectronic balance

beakers

stirring rod (with one end covered by rubber)

fi lter funnel and stand

fi lter paper

Buchner funnel and fl ask

mortar and pestle

heating equipment

samples of lawn food

1 mol L–1 hydrochloric acid

0.5 mol L–1 barium chloride solution

methylated spirits

METHOD

1. Grind up some lawn food in a mortar and pestle, then accurately weigh out

about 1 g.

2. Add this to about 100 mL of distilled water and dissolve as much as

possible.

3. Filter the solution into a 400 mL (or larger) beaker. Wash the residue in the

fi lter paper with several portions of distilled water.

4. Add about 5 mL of 1 mol L–1 hydrochloric acid to the fi ltrate, followed by

suffi cient distilled water to give a total volume of 200 mL.

5. Heat to near boiling and then add about 15 mL of the barium chloride

solution, stirring throughout.

6. Allow the precipitate to settle and then add a few drops of the barium

chloride solution. If a precipitate forms, add a further 2 mL of the barium

chloride solution. Repeat this step until no further precipitate forms.

7. Allow the mixture to stand for as long as possible before fi ltering.

8. Weigh a clean and dry suction funnel, together with a piece of fi lter

paper.

9. Carefully fi lter the supernatant liquid and then the precipitate. Scrape

diffi cult-to-remove traces of precipitate into the funnel. Rinse the beaker a

few times with water and add these washings to the fi lter paper.

10. Wash the precipitate with several portions of warm water followed by

methylated spirits.

11. Dry the crucible and contents in an oven at 110°C overnight.

12. Weigh when dry and calculate the mass of barium sulfate produced.

Exercise care with the

hydrochloric acid. Use

methylated spirits away from

any naked fl ames.

Exercise care with the

hydrochloric acid. Use

methylated spirits away from

any naked fl ames.

EXPERIMENT 2.1

Chemistry 2: ExperimentsCONTINUED



8

EXPERIMENT 2.1

13. Use your results to calculate the percentage of sulfur in the lawn food

tested.

QUESTIONS

1. Compare your result against the manufacturer’s specification (if this is

available).

2. In what parts of the method are errors likely to occur? Discuss the effect of

these on your final result.

3. Why were the contents of the filter paper washed with water in step 3?

4. Why was the precipitate washed with water in step 10?


9


Ammonia content of cloudy ammonia — a problem-solving exercise

Cloudy ammonia contains ammonia (a base), as its ‘active ingredient’. In use,

the ammonia helps by reacting with greasy surfaces to form a type of soap.

The presence of a few other additives makes this product suitable for use as a

cleaning agent.

You are to design a volumetric procedure to check the amount of ammonia

in some commercially available cloudy ammonia. This can then be compared

with the claims on the label.

CONSIDERATIONSAs ammonia is a base, a standardised solution of hydrochloric acid can be

used in this titration. The equation for the reaction that then occurs is:

HCl(aq) + NH3(aq) NH4+(aq) + Cl–(aq)

As ammonia is volatile, your procedure should involve handling techniques

that keep such losses to a minimum.

Methyl orange is a suitable indicator for this reaction.

•

•

•

Ammonia vapour may irritate

the skin, eyes and respiratory

system. The laboratory should be

well ventilated and safety glasses

should be worn.


hydrochloric acid solution.

Ammonia vapour may irritate

the skin, eyes and respiratory

system. The laboratory should be

well ventilated and safety glasses

should be worn.





10


Nitrogen content of lawn fertiliser

Nitrogen is an important nutrient for plant growth, and many nitrogenous

fertilisers are currently on the market. These usually contain the required nitrogen

as ammonium sulfate. Other nutrients may or may not also be present.

In this experiment, the manufacturer’s claims are checked by adding an

excess of sodium hydroxide to a sample of lawn food solution. The excess is

then determined by titration against standardised hydrochloric acid.

AIMTo determine the percentage of available nitrogen in a brand of lawn fertiliser

APPARATUSburette and stand

small funnel

white tile

wash bottle

2 pipettes and a pipette fi ller

volumetric fl ask

weighing bottle

balance


hydrochloric acid solution

standardised 0.1 mol L–1 sodium hydroxide solution

phenolphthalein indicator

METHOD

1. Place an accurately weighed mass of about 1.2 g of the fertiliser into a

250 mL volumetric fl ask.

2. Add some distilled water and dissolve as much of the fertiliser as possible.

Make up to the mark with more distilled water.

3. Use a pipette to deliver 20.00 mL of sodium hydroxide (of accurately

known concentration) into a conical fl ask. Note: If the concentration of the

sodium hydroxide is not known, it will need to be standardised.

4. Use another pipette to add 20.00 mL of the fertiliser solution to the same

conical fl ask.

5. Add about 50 mL of distilled water to this fl ask, and then boil the mixture

for about 10 minutes. A funnel placed on top of the fl ask will help prevent

the hot mixture from splashing out and will help when adding more water

from time to time, so that the volume in the fl ask stays fairly constant.

6. Repeat steps 1 to 5 with the other two conical fl asks.

7. After the fl ask has cooled, add 3 drops of phenolphthalein indicator and

titrate using standardised 0.1 mol L–1 hydrochloric acid.

8. Repeat step 7 with the other two fl asks (or more if necessary), until three

concordant titres are obtained.

9. Calculate the percentage of nitrogen in the fertiliser sample.

QUESTIONS

1. Why is the mixture boiled in step 5 above?

2. If the sodium hydroxide was incorrectly standardised, would this represent

a source of systematic or random error in this experiment?

Care should be taken with the

boiling process. Hot sodium

hydroxide solution will be

present.

Wear safety goggles and

immediately wash off any spills

with water.

Take care while handling the


Care should be taken with the

boiling process. Hot sodium

hydroxide solution will be

present.

Wear safety goggles and

immediately wash off any spills

with water.

Take care while handling the


EXPERIMENT 3.2


11


Standardisation of hydrochloric acid

Hydrochloric acid solution is used in many volumetric procedures. Prior to its

use, however, its concentration needs to be accurately determined.

A common method for doing this is to react it against a solution of sodium

carbonate — a substance that makes a good primary standard.

The equation for the reaction is:

2HCl(aq) + Na2CO3(aq) 2NaCl(aq) + H2O(aq) + CO2(g)

Once the accurate concentration is established it may be used in other

experiments.

AIMTo determine the accurate concentration of a solution of hydrochloric acid


20.00 mL pipette and pipette fi ller


white tile

weighing bottle

spatula

small funnel

wash bottle

250 mL standard fl ask

balance

anhydrous sodium carbonate

0.1 mol L–1 hydrochloric acid

methyl orange indicator

METHOD

1. Calculate the amount of anhydrous sodium carbonate that is required to

make up 250 mL of a 0.05 mol L–1 solution.

2. Accurately weigh out close to this amount of sodium carbonate in a

weighing bottle or other suitable container. Using a dry funnel, transfer

the solid into the volumetric fl ask. Any particles remaining in the weighing

bottle should be carefully washed into the fl ask using the wash bottle.

3. About half-fi ll the volumetric fl ask and swirl to dissolve the sodium

carbonate. Make up to the mark on the fl ask with distilled water and swirl

to ensure uniform concentration.

4. Rinse the burette with the hydrochloric acid solution, then fi ll the burette

with this acid.

5. Rinse the pipette with a small portion of the sodium carbonate solution.

6. Pipette 20.00 mL of sodium carbonate solution into a conical fl ask. Add

3 drops of methyl orange indicator and titrate with acid from the burette

until the end point occurs. A white tile placed under the titration fl ask will

make this easier to detect. Note the volume of acid used.

7. Repeat step 6 until three burette readings within 0.1 mL are obtained.

8. Use your results to calculate the accurate concentration of the hydrochloric

acid.

Take care with the hydrochloric

acid. Any spills should be

washed immediately with a large

volume of water.

Take care with the hydrochloric

acid. Any spills should be

washed immediately with a large

volume of water.

EXPERIMENT 3.3




12

QUESTIONS

1. The sodium carbonate would have been heated prior to the experiment.

Suggest a reason for this.

2. Why can’t sodium hydroxide be used as the primary standard in this

experiment?

3. If a conical flask needs to be reused for a second titration, it should be

rinsed out thoroughly with water. However, it is not necessary that it be

dried before re-use. Explain why.

EXPERIMENT 3.3


13


Indicators — a problem-solving exercise

The following problems concern indicators and pH.

Part AStudents were provided with 0.10 mol L–1 solutions of hydrochloric acid and

sodium hydroxide. Using pipettes, they prepared mixtures of these two solutions

and then measured their pHs using narrow range universal indicator paper. The

results are tabulated below.

Acid (mL) Base (mL) pH

0 100 13.0

10 90 12.9

20 80 12.8

30 70 12.6

40 60 12.4

45 55 10.6

50 50 7.0

55 45 3.0

60 40 1.7

70 30 1.4

80 20 1.2

90 10 1.1

100 0 1.0

1. Plot a line graph of pH versus the volume of acid added to the mixture.

2. Determine the pH of the mixture formed on mixing:

(a) 48 mL HCl and 52 mL NaOH

(b) 52 mL HCl and 48 mL NaOH.

3. The following table shows the colour ranges and pH of different

indicators.

Indicator pH range Colour range

methyl orange 3.1–4.4 red–yellow

bromocresol green 3.8–5.4 yellow–blue

chlorophenol red 5.2–6.8 yellow–red

bromothymol blue 6.0–7.6 yellow–blue

phenol red 6.8–8.4 yellow–red

Predict the colour of each of these indicators in the following mixtures:

(a) 60 mL HCl + 40 mL NaOH

(b) 55 mL HCl + 45 mL NaOH

(c) 50 mL HCl + 50 mL NaOH

(d) 48 mL HCl + 52 mL NaOH

EXPERIMENT 3.4




14

Part BClare collected some flower petals and extracted the coloured dye pigment into

water. The mixture was filtered and the yellow filtrate poured into three test

tubes (X, Y and Z). She performed the following tests on each tube.

Tube Test Observation

X drops of hydrochloric

acid added

yellow extract

turned orange

Y drops of sodium

hydroxide added

yellow extract

turned green

Z water added remained yellow

1. Identify the colour of the extract in basic solution.

2. Predict the colour of the extract in soda water.

3. Predict the colour of the extract in ammonia solution.

4. Predict the colour of the extract in salt water.

Part CTwo indicators have been codenamed DING and DONG. The following

information was collected about these two indicators.

DING turns green when substance P is present.

DONG turns orange when substance Q is present.

DING and DONG were added to samples of four colourless solutions (I, II, III

and IV) that have been made using P, Q, or both. The results of this experiment

are tabulated below.

Solution DING DONG

I blue orange

II green yellow

III green orange

IV blue yellow

Use this information to decide whether each of the following statements is supported or not supported by the data.

1. I contains Q but not P.

2. II contains both P and Q.

3. III contains P but not Q.

4. IV contains Q but not P.

•

•

EXPERIMENT 3.4



15

EXPERIMENT 3.4

Part DTwo indicators are codenamed PING and PONG. The colours of these indicators

are determined in the presence of two substances (D and E). The results are

tabulated below.

Indicator

Substance

added Colour

PING E red

PING D + E orange

PONG D yellow

PONG D + E yellow

PING and PONG were then added to four colourless solutions (M, N, P, R).

These solutions contained D; E; D and E; or neither. The results are tabulated

below.

Solution added PING PONG

M no change no change

N turns orange turns yellow

P turns red no change

R no change turns yellow

1 . Identify the solution that contains substance E but none of substance D.

2. Account for the change in PING’s colour when D is added to E.

3. There was no change to the indicator colours with solution M. Suggest a

reason for this observation.



16


Separation of food dyes using paper chromatography

Dyes are used in foods to make the food more attractive and to give it a

consistent colour. In the confectionery industry in particular, brightly coloured

sweets and lollies are purchased by the consumer. Hence, in this area of the

food industry, a large number of food dyes are used to achieve the many and

varied colours displayed by today’s sweets.

The separation and identification of these dyes is especially suited to the

technique of chromatography. This technique has the added advantage that it

usually requires only small amounts of material.

In this experiment, paper chromatography is used to examine some

purchased food dyes, and then a mixture made from some of these dyes is run

to demonstrate how separation occurs. As an extension, the method suggested

is easily adapted to identifying the dyes present in Smarties.

AIMTo examine the separation by paper chromatography of a number of commercially

available food dyes

APPARATUSlarge beaker

filter paper

Gladwrap (or similar)

pencil

ruler

capillary tubes

samples of commercial food dyes

at least one mixture made from the dyes that are being tested

METHOD

1. Obtain the largest beaker available, and prepare a sheet of filter paper so

that it will fit inside the beaker as shown in the figure below.

2. Draw a pencil line 2 cm from the bottom of this as shown. Mark off this

line at 1 cm intervals. Label each mark with a letter.

600 mL beaker covered

with plastic wrap

Chromatography

paper

Origin

(pencil line)

Solvent

Staples

3. Using a capillary tube, place a small spot of either pure dye or dye mixture

at each location. Do not make the size of this drop too large — it should

be no larger than the letter ‘o’. Note locations to avoid future confusion.

EXPERIMENT 4.1



17

EXPERIMENT 4.1

4. Stand the filter paper inside the beaker, after having added a small amount

of water that will act as the solvent. Make sure that the pencil line is above

the level of this water.

5. Cover the top of the beaker with Gladwrap, and allow the solvent to rise up

the filter paper until it is close to the top.

6. Remove the filter paper, and, using a pencil, mark in the positions reached

by the solvent and each dye component. Note: If the dye components

appear as bands, mark the centre of the band.

7. Calculate Rf values for each dye and comment on your findings.

EXTENSIONDesign an experiment to identify the dyes used in Smarties. The dye from

each Smartie could be spotted onto the pencil line. Samples of the following

dyes will need to be on hand so that identification can occur: azorubine, sunset

yellow FCF, ponceau 4R, tartrazine, green S, brilliant blue FCF.

A mixture of butanol, ethanol and 2 mol L–1 ammonia in the proportions

3:1:1 reportedly gives good separation.



18


Separating mixtures using column chromatography

Both the gas chromatograph (GC) and the high-performance liquid chromato-

graph (HPLC) use a column to separate the components of a mixture.

In this experiment, a mixture of dyes will be separated using the more

traditional method of column chromatography. This should help you to visualise

what happens inside the column of a GC or HPLC instrument.

AIMTo separate a mixture of dyes using column chromatography

APPARATUSchromatography column (or a suitably modifi ed burette)

dropping pipette

small funnel

long piece of thin wire

mixture of commercially available food dyes

methanol

alumina

METHOD

1. Fill a chromatography column about half full with methanol.

2. Using the small funnel, add alumina carefully to the column. The alumina

should be evenly packed, with no bubbles present. If the packing becomes

uneven, carefully use a long piece of wire to pack it down. Fill the alumina

to about 5 cm from the top, making sure that it is covered with a further

2–3 cm of methanol.

3. Prepare a ‘developing solvent’ by mixing together 80 mL of methanol and

20 mL of water.

4. Allow the methanol to begin draining from the column. When the level of

liquid in the top reaches the surface of the alumina, add about 2 mL of the

dye mixture to the top of the column. Take care not to disturb the surface

of the alumina at this stage. Add a little more methanol.

5. When this mixture has just reached the surface of the alumina, very carefully

add some developing solvent. For as long as time permits, keep adding

small portions of this solvent, making sure that the top of the column never

dries out.

QUESTIONDescribe the appearance of the column at various times.

Methanol is poisonous and

fl ammable. Avoid contact with

eyes and naked fl ame.

Methanol is poisonous and

fl ammable. Avoid contact with

eyes and naked fl ame.

EXPERIMENT 4.2


19


A model GC detector

A common type of detector used in gas chromatographs is a flame ionisation

detector (FID). This was invented by two Australians, McWilliam and Dewar,

in 1957–58.

The basis of the FID is that ions will be produced when organic molecules

are burnt in a flame. The flame will therefore conduct an electric current, with

the amount of current being proportional to the number of organic molecules

in the flame.

AIMTo simulate the operation of a flame ionisation detector

APPARATUSmultimeter

two sewing needles (to act as electrodes)

Bunsen burner

candle

cotton wool

1,1,1-trichloroethane

METHOD

1. Set up two sewing needles close enough together so that they are both in

the flame.

2. Connect the electrodes to the multimeter via the probes of the multimeter.

3. Switch the multimeter to the ohms scale. (This provides a constant voltage

across the electrodes.)

4. Use various types of flame and record the readings from the multimeter.

Flames that might be tested include a Bunsen flame, a candle, ethanol (soak

a cotton wool ball in ethanol). Take care that the probes and wires from the

multimeter do not get too hot.

5. Soak a cotton wool ball in 1,1,1-trichloroethane and place beside the airhole

of the burner. Observe the effect on the detected current.

QUESTIONDid each type of flame produce the same reading and therefore the same number

of ions in its flame?



20


Flame colours

In the qualitative analysis of an unknown ionic compound, some idea of the

cation present may be obtained from fl ame tests as described below. A number

of cations do not produce fl ame colours; however, for those that do, the colour

is often distinctive and can point to the presence of a particular cation, which

may be subsequently identifi ed using a follow-up chemical test.

AIMTo observe the fl ame colours produced by a number of cations

APPARATUSplatinum or Nichrome wire loops

2 watchglasses

Bunsen burner

hydrochloric acid

salts to be tested

LABORATORY PREPARATION NOTESThe test loops can be made from either platinum or Nichrome wire. A small

loop is formed in one end of a 10 cm length and the other end is pushed into

a small piece of balsa wood which then acts as the holder. (Paperclips may be

used as an alternative, but will need to be disposed of after each test as they

are diffi cult to clean.)

Samples of solid sodium chloride, potassium chloride, calcium chloride,

strontium chloride and barium chloride will need to be provided for testing.

METHOD

1. Place a small amount of concentrated hydrochloric acid into 2 watchglasses.

One of these will be used for cleaning the wire loop between tests and will

probably quickly become contaminated. The other will be used to dip the

clean loop into, just before doing each test.

2. Clean the wire loop by dipping it into the cleaning acid and heating it in

the fl ame. This procedure will need to be repeated until the wire imparts no

visible colour of its own to the fl ame.

3. Dip the now clean wire into the clean acid and then into a solid sample of

the salt to be tested.

4. Place the loop into the fl ame, and note the colour produced.

5. Repeat steps 1 to 4 for the other salts provided.

QUESTIONS

1. Make a list of as many applications as you can think of that use this principle,

that is, the production of coloured light from a heated substance.

2. How would you modify this experiment to convince yourself that the

observed colours are coming from the cations in each salt, and not from

the anions?

EXTENSIONYou might like to use a spectroscope to try and note the spectrum produced by

each of these cations. This will need to be done in a darkened room to avoid

stray light interfering with the spectrum.

EXPERIMENT 5.1

Beware of the wire used for

testing — it gets very hot!

Concentrated hydrochloric acid

must be treated with care. If you

spill any, immediately wash with

large amounts of water. If you

get it in your eyes, fl ush them for

15 minutes with an eye wash and

seek medical attention.


21


Zinc content of cornflakes using atomic absorption spectroscopy

A popular brand of cornflakes claims that its zinc content is 6.0 mg per 100 g

of cornflakes.

This claim is ideally suited for checking by atomic absorption spectroscopy.

Listed below are the details of an experiment that checked this claim, together

with the results obtained.

AIMTo interpret some typical data obtained using atomic absorption spectroscopy

APPARATUSpencil and graph paper

METHODA 120.0 g sample of cornflakes was treated to extract all the zinc into a solution

with a volume of 500 mL. Exactly 125 mL of this solution was then diluted to

a volume of 1.00 L.

This solution, along with a number of standard zinc solutions, was then

aspirated into an atomic absorption spectrometer and the absorbance recorded.

The results are shown below.

RESULTS

Solution concentration (ppm) Absorbance

0.00 0.010

1.00 0.100

2.00 0.190

3.00 0.280

4.00 0.370

Sample 0.172

Note: ppm is the same as mg L–1

QUESTIONS

1. Cornflakes also contain iron. How would you make sure that the AAS is

measuring zinc content and not iron content?

2. Why is it not necessary to remove all the iron from the samples that you

are analysing in these experiment?

3. Plot the calibration curve (i.e. the graph of absorbance vs concentration)

produced from these data.

4. From your graph, determine the concentration of zinc in the diluted sample

that was analysed.

5. Hence, calculate the concentration of zinc in the undiluted 500 mL sample

of extract.

6. From your answer to question 5, calculate the amount of zinc that is

present in 100 g of the corn flakes. How does your answer compare to the

manufacturer’s claim?

7. A typical serving size is one cup (30 g). What will be the mass of zinc per

serve?

EXPERIMENT 5.2



22


Iron content of wine using UV–visible spectroscopy

The iron content of wine can be determined by either atomic absorption

spectroscopy or by UV–visible spectroscopy. The data below are obtained from

the latter.

To obtain these data, the iron present in the wine was treated to produce the

intensely coloured iron(II) 1,10–phenanthroline complex ion, Fe(C12H8N2)32+.

AIMTo interpret some typical data obtained using UV–visible spectroscopy

APPARATUSpencil and graph paper

METHODTwo samples of wine were prepared for analysis by taking a certain volume and

diluting it to 100 mL. In the first case (unknown 1) a 10.0 mL sample was used

for dilution, while 20.0 mL was used for the second sample (unknown 2).

The absorbance of each of these samples was measured, along with a set of

standard solutions of known iron concentration.

RESULTS

Concentration of iron (ppm) Absorbance

0.0 0.000

0.1 0.015

0.5 0.095

1.0 0.160

2.5 0.480

unknown 1 0.005

unknown 2 0.040

Note: ppm is the same as mg L–1

QUESTIONS

1. Prepare a calibration curve from the above results.

2. From this curve, determine the level of iron in both the diluted samples.

3. Calculate the amount of iron present in each of the undiluted wine

samples.

EXPERIMENT 5.3


23


Phosphates in detergents

To estimate the phosphate levels in detergents, a procedure called colorimetric

analysis may be used. This involves the addition of ammonium molybdate,

a chemical that forms a blue compound if phosphate is present. The more

phosphate, the more intense the blue colour of the solution. If a set of standards

containing known phosphate levels is produced in the same way, the level in

the unknown sample may then be determined by colour matching. Although

we will do this visually, instruments may also be used to measure the light

transmitted through such solutions to obtain more accurate results.

AIMTo determine the amount of phosphate in various detergents by colorimetric

analysis

APPARATUS10 clean conical fl asks

100 mL measuring cylinder

ascorbic acid

stock phosphate solution

ammonium molybdate – sulfuric acid solution

heating equipment

selection of detergents

2 × 1 L beakers

burettes

LABORATORY PREPARATION NOTESThe following steps should be performed by a teacher or a laboratory

technician.

Ammonium molybdate – sulfuric acid solution

1. Add 15.0 g of ammonium molybdate to 150 mL of distilled water, cooling

in an ice bath as it is dissolved. Leave in ice bath.

2. Slowly add 250 mL of concentrated sulfuric acid to 250 mL of water.

(Remember to add the acid to the water, never the reverse!) If the solution

gets too hot, stop and wait for it to cool before proceeding. Chill the solution

in another ice bath when fi nished.

3. Slowly and carefully add the cold ammonium molybdate solution to the

cold sulfuric acid solution.

4. The resulting solution can be stored in plastic bottles. Immediately prior

to performing this experiment, the solution should be placed in a burette,

from which it can be dispensed. The burette should be labelled with the

appropriate safety information.

Standard phosphate solution

Using disodium monohydrogen phosphate, Na2HPO4.12H2O weigh out

3.0265 g and dissolve it in 1 L of distilled water. If the salt you are using has

a different degree of hydration, the mass of salt required will of course be

different. Check the label fi rst!

Immediately prior to the lesson, take a 10 mL aliquot of this stock solution

and dilute it to 1 L. This will give a solution that is 8.0 mg L–1 with respect to

phosphate.

EXPERIMENT 5.4

The ammonium molybdate

– sulfuric acid solution in the

burette contains sulfuric acid

of approximately 7 mol L–1

concentration. Be extremely

careful with this solution. Any

spills on skin or clothing should

be immediately fl ushed with large

volumes of water.




24

METHOD

Part A: Preparation of phosphate standards

1. Into 10 conical flasks, measure respectively 70, 60, 50, 40, 30, 20, 15, 10,

5 and 0 mL of the phosphate solution supplied.

2. Add sufficient distilled water to the flasks so that the total volume becomes

100 mL.

3. Very carefully add 10 mL of the ammonium molybdate – sulfuric acid

solution from the burette to each flask. Swirl to mix.

4. Add a few crystals of ascorbic acid to each flask, mix and bring each

solution to the boil.

5. Line up the flasks in order of decreasing colour intensity. You will now

have colour standards corresponding to 5.6, 4.8, 4.0, 3.2, 2.4, 1.6, 1.2, 0.8,

0.4 and 0.0 mg L–1 of phosphate.

Part B: Analysing the detergentsLiquids

1. Accurately weigh out approximately a 1 g sample of the selected detergent.

Add 200–300 mL of water, mix well and dilute to 1 L.

2. Place 100 mL of this diluted solution into a conical flask and repeat steps

3 and 4 from part A above.

3. Compare the colour of the sample with that of the standard phosphate

solutions and estimate the sample’s concentration.

Powders

1. Accurately weigh out 1 g of powder and add this to 200–300 mL of water.

If the powder does not dissolve, heat gently in a beaker, but do not boil. Mix

well and then transfer the solution to a volumetric flask. Dilute to 1 L.

2. Take a 25 mL aliquot of the solution from step 1 and dilute this to 1 L.

3. Place 100 mL of the solution from step 2 into a conical flask and repeat

steps 3 and 4 from part A.

4. Compare the colour of the sample with that of the standard phosphate

solutions and estimate the sample’s concentration.

QUESTIONCalculate the mass of phosphate in each detergent sample and hence the

percentage of phosphate in each sample. Where possible, compare your results

to the manufacturer’s claims.

EXTENSIONIf a light meter is available, the amount of light transmitted through each of the

reference solutions could be measured, and a calibration graph drawn from the

results. The transmittance through the test sample could then be measured in

the same way, and the result read off from the graph.

EXPERIMENT 5.4


25


Constructing models of alkanes and alkenes

AIMTo construct models of alkanes, alkenes and their isomers

APPARATUSMolecular building kit. If this is unavailable, coloured plasticine and straws

or satay sticks will do. Roll the plasticine into small balls (the size of a small

marble), making sure that a different colour is assigned to each element. Cut

the straws or satay sticks into 5 cm lengths.

METHOD

1. In your record book, draw the structural arrangement for the first five

members of the alkane and alkene series.

2. Construct the three-dimensional model for these hydrocarbons.

3. Construct the three-dimensional models for the different isomers of butane,

pentane, butene and pentene. How many isomers of each did you get?

4. Draw the structures of the isomers that you have constructed.

5. Name the isomers.

QUESTIONS

1. What are the names of the first five members of the alkane and alkene

series?

2. Explain the differences between alkanes and alkenes.

3. What are isomers?



26


Constructing models of organic compounds

AIMTo construct models of a variety of organic compounds

APPARATUSMolecular building kit. If this is unavailable, coloured plasticine and straws

or satay sticks will do. Roll the plasticine into small balls (the size of a small

marble), making sure that a different colour is assigned to each element. Cut

the straws or satay sticks into 5 cm lengths.

METHOD

1. In your record book, draw the structural arrangement for:

(a) tetrachloromethane

(b) 1,2-dibromopropane

(c) ethanol

(d) polyethene (from three monomer units)

(e) 1,1,2,2-tetrachloroethane

(f) butene

(g) 1-chloro-2-methylprop-2-ene

(h) pentabromoethane

(i) methanoic acid

(j) ethanoic acid.

2. Construct the three-dimensional molecular model for each of the compounds

in step 1.

QUESTIONS

1. Is there another way of arranging the two bromine atoms in 1,2-

dibromopropane?

2. What are functional groups? What are the functional groups of ethanol and

ethanoic acid?

3. Are there other isomers for tetrachloroethane apart from the one you

constructed: 1,1,2,2-tetrachloroethane? Draw their structural formulae.

EXPERIMENT 6.2


27


Identifying alkanes and alkenes

AIMTo distinguish between alkanes and alkenes by a number of simple tests

APPARATUS2 watchglasses

4 test tubes with stoppers

cyclohexane

cyclohexene

bromine dissolved in trichloroethane

1 mol L–1 sulfuric acid

0.02 mol L–1 potassium permanganate solution

METHOD

1. Perform this step in a fume hood. Place a few drops of each hydrocarbon

on separate watchglasses. Ignite them, and compare the sootiness and

luminosity of each fl ame.

2. Place about 1 mL of each hydrocarbon into separate test tubes. Add 1 mL

of the bromine solution to each. Stopper and shake, noting any colour

changes.

3. Add 5 drops of each hydrocarbon to two further test tubes. Add about

10 drops of the potassium permanganate solution and a few drops of the

sulfuric acid solution to each. Stopper and shake them, observing any

changes.

QUESTIONS

1. Write an equation for the reaction that occurs in step 2.

2. Which class of compounds, alkanes or alkenes, seems to be the more

reactive? Explain this in terms of their structure.

Take care with the sulfuric acid

and bromine solutions.

Cyclohexane and cyclohexene

are fl ammable.

Take care with the sulfuric acid

and bromine solutions.

Cyclohexane and cyclohexene

are fl ammable.

EXPERIMENT 6.3



28


Carbohydrate models

Carbohydrates are molecular substances made up of carbon, hydrogen and

oxygen.

AIM

1. To construct both the linear and cyclic forms of a glucose and fructose

molecule

2. To use the cyclic structures to form a disaccharide and polysaccharide

APPARATUSplasticine, toothpicks, or a molecular modelling kit

METHODUsing plasticine and toothpicks (or a molecular modelling kit, if available)

construct both the linear and cyclic forms of a glucose and a fructose molecule

(see page 154 of textbook).

QUESTIONS

1. Glucose and fructose are said to be isomers (i.e. same molecular formula,

but a different arrangement of atoms).

(a) Use the models you have constructed to describe the similarities and

differences between the two molecules.

(b) Describe how the physical properties of each molecule can be

determined by its structure.

2. (a) Which groups on the two molecules participate in the condensation

polymerisation reaction that forms a disaccharide?

(b) Can more than one disaccharide be produced from the condensation

polymerisation between glucose and fructose? Explain your answer.

(c) If the disaccharide is to be split into each monosaccharide, what would

be needed for the reaction to occur?

3. (a) Use models of glucose to show how a part of a starch molecule would

form.

(b) Why is the formation of starch from glucose classed as a condensation

polymerisation reaction?

(c) Describe the similarities and differences between the formation of

starch and cellulose from glucose.

4. Write equations to represent the formation of starch and cellulose from

glucose.


29


Modelling the structure and denaturation of proteins

Proteins contain carbon, hydrogen, oxygen and nitrogen. Proteins can have

four structural levels distinguishable. Unfolding of the polypeptide chains,

denaturation, leads to loss of biological activity.

AIMTo model the four structural levels distinguishable in proteins and the

denaturation of a protein

METHODUsing a streamer, demonstrate:

1. the primary structure of a protein by drawing a tripeptide along the length

of the streamer

2. the secondary structure of a protein by twisting the streamer to form a

spiral, while holding the internal part with your fingers

3. the tertiary structure of a protein by gathering in the spirals

4. the quaternary structure of a protein by combining your spiral with that/

those of at least one other student

5. the denaturation of a protein by crumpling up or changing the shape of

your spiral.

QUESTIONS

1. Describe the four levels of bonding that can exist in proteins.

2. Which level of bonding is affected when denaturation occurs?



30


A stomach enzyme

Many of the early experiments on the workings of the stomach were the result of

an accident. In 1822, a man was badly injured by a shotgun wound to the chest.

The blast had torn away ribs and muscles as well as part of the stomach wall.

Although he survived the accident, the wound healed in an unusual way. The

skin and stomach wall grew together and failed to join up completely. This meant

that he had a hole in his side that led straight to his stomach. Bandages had to

be kept over the hole to prevent food from falling out. His doctor could see what

was happening in his stomach and realised that this was a good opportunity to

try some experiments. One of the experiments he tried was to see what happened

when chemicals made by the stomach were placed on food.

AIMTo investigate the effect of two chemicals produced by the stomach, hydrochloric

acid and pepsin, on digestion of food

APPARATUS4 test tubes

test-tube rack

marking pen

mortar and pestle

1 teaspoon of minced beef

dilute hydrochloric acid

pepsin solution

METHOD 1. Take a teaspoon of minced beef and grind it up with a mortar and pestle.

2. Place a small sample in each of four test tubes.

3. Cover the beef in one tube with water.

4. In the next tube, cover the meat with hydrochloric acid.

5. In the third tube, cover the meat with pepsin solution.

6. Cover the beef in the fourth test tube with an equal mixture of hydrochloric

acid and pepsin solution.

7. Label each tube and store overnight in an incubator or oven at about 37°C.

Observe and record your results the next day in the form of a table.

Test-tube contents Extent of digestion

water

hydrochloric acid

pepsin solution

acid and pepsin mixture

QUESTIONS

1. What part of the digestive system would normally grind meat?

2. Why are the tubes incubated at 37°C?

3. What was the purpose of the tube with water and meat?

4. Use your results to explain how protein may be digested in the stomach.

EXTENSIONDevise and conduct an experiment to test the ideal proportion of hydrochloric

acid to pepsin solution needed for the most efficient digestion of protein.

EXPERIMENT 7.3


31


The action of bile

Bile is produced by the liver and stored in the gall bladder. When food enters

the duodenum, the muscular wall of the gall bladder is stimulated to contract

and the bile is forced down the gall duct into the duodenum. Bile contains

substances called bile salts. These act on fat, breaking it up into small droplets,

thereby acting as emulsifying agents.

AIMTo investigate the action of bile salts on fat


marking pen

15 mL of oil

1 microspatula of bile salts

5 drops of liquid detergent

METHOD

1. Label three test tubes A, B and C.

2. Pour some oil into each test tube to a depth of 5 cm.

3. Add a few drops of water to A and shake well. Record your observations.

4. Add a microspatula of powdered bile salts to B and shake well. Record

your observations.

5. Add a few drops of liquid detergent to C and shake well. Record your

observations.

QUESTIONS

1. What is the purpose of test tube A?

2. How do bile salts help digestion?

3. In what way are bile salts similar to liquid detergent in action?

EXPERIMENT 7.4



32


Testing for nitrates in soils

When concentrated sulfuric acid reacts with a mixture of iron(II) ions and a

nitrate, a brown colour results:

3Fe2+(aq) + NO3–(aq) + 4H+(aq) 3Fe3+(aq) + NO(g) + 2H2O(l)

and Fe2+(aq) + NO(g) FeNO2+(aq) brown

This fact is used as a test for the presence of nitrates. To do the test, equal

volumes of a nitrate solution and fresh iron(II) sulfate are mixed in a test tube.

Concentrated sulfuric acid is then poured slowly down the side of the sloping

test tube so that it forms a layer on the bottom. A brown ring forms where the

layers meet. (See the fi gure below.)

Nitrate +

Iron(II)

sulfate

Concentrated

sulfuric acid

Brown

ring

AIMTo test for the presence of soluble nitrates in soil samples

METHOD

1. Collect 5 g each of at least fi ve different soil samples. Include at least one

sample in which plants grow well and another sample in which little plant

growth is observed. Record the extent of plant growth (fertility) in each

sample.

2. Since nitrates are soluble in water, extract the nitrate in each soil sample

by mixing the soil with 10 mL of water. Filter the mixture and retain the

fi ltrate. Keep each sample separate.

3. Test for the presence of nitrates in each sample by carefully pouring some

concentrated sulfuric acid down the side of a test tube containing a portion

of the fi ltrate collected from each soil sample. Record your results.

QUESTIONS

1. How did your results for each sample compare? Try to explain your results

in terms of the observed fertility of the soil.

2. Is this test a good measure of soil fertility? Justify your answer.

3. How could this test be improved to give a better indication of soil

fertility?

4. What other factors are important in determining soil fertility? Outline an

experiment that could test for at least one of these factors.

5. Construct a fl ow chart to describe how the nitrates in the soil impact on the

food supply.

Care must be taken when

handling concentrated sulfuric

acid. Wash spills immediately

with copious amounts of water.


33


Electrophoresis simulations

AIMTo investigate protein separation using electrophoresis activities

METHOD

Part A: Interpreting gel electrophoresisGel electrophoresis is a common method for identifying proteins or polypeptides

and determining their molecular weights. The proteins and reference polypeptides

are run in an acrylamide gel of known concentration. The position of the protein

bands is found using appropriate staining techniques.

1. The figure below shows the result of an electrophoresis experiment on an

unknown mixed protein sample (X) run in a 10% acrylamide gel.

The unknown (X) is run in Lane 1. High molecular weight range

reference proteins are run in Lane 2. Low molecular weight range proteins

are run in Lane 3.

Lane 1

Unknown

X

High molecular

weight range

reference proteins

Low molecular

weight range

reference proteins

Lane 2

Myosin

Start

End

Betagalactosidase

Carbonicanhydrase

Trypsinogen

Lane 3

Bovine serumalbumin

Egg albumin

The reference samples can be used to determine the proteins present in sample X.

2. Examine the data in the above figure and use this information to determine

which of the reference standard proteins (or polypeptides) are present in

sample X.

3. What evidence is there that other proteins (or polypeptides) are also present

in the sample?

4. Explain the relationship between the molecular weight of the proteins and

their relative migration position in the gel.

Part B: Determining the molecular weight of a protein

1. The molecular weight of a protein and its relative migration distance (or

mobility) in the gel are related. The table on the following page provides

data for some reference standard proteins and an unknown protein (Y) for

an experiment using a 10% acrylamide gel.

EXPERIMENT 8.1




34

Protein/polypeptide

reference marker

Relative migration

distance Molecular weight

trypsin inhibitor (soybean) 0.85 20 100

trypsinogen (bovine) 0.75 23 700

carbonic anhydrase 0.63 29 000

G-3-P-dehydrogenase

(rabbit)

0.54 36 000

egg albumin 0.45 45 000

bovine albumin 0.33 66 000

phosphorylase B (rabbit) 0.23 97 400

unknown (Y) 0.70 ?

2. Use a spreadsheet application (such as Excel) to plot a line graph of

molecular weight (y-axis) versus relative migration distance (x-axis). Draw

the line of best fit.

3. Interpolate from the graph the molecular weight of protein Y.

4. An enzyme (Z) has a relative migration distance of 0.95 under the same

conditions. Explain why its molecular weight could not be reliably

determined from these data.

EXPERIMENT 8.1


35


Preparation of aspirin

Aspirin, otherwise known as acetylsalicylic acid, is a well-known drug used to

relieve aches and pains, reduce inflammation and lower fever. It also has blood

thinning properties. Aspirin was originally obtained from the bark of the white

willow tree. Acetylsalicylic acid can be produced by reacting salicylic acid with

acetic anhydride.

AIM To prepare and test a sample of aspirin (acetylsalicylic acid)

APPARATUS250 mL conical flask

retort stand and clamp

filter flask

Buchner funnel

filter paper

500 mL beaker

3.0 g salicylic acid

8 mL acetic anhydride

1 mL concentrated sulfuric acid

0.1 mol L–1 iron(III) nitrate

ice water

aspirin tablet

METHOD

1. Weigh a piece of filter paper and record the mass.

2. Accurately weigh about 3.0 g salicylic acid into a clean, dry 250 mL conical

flask.

3. In a fume cupboard, carefully add 8 mL of acetic anhydride and 1 mL of

concentrated sulfuric acid.

4. Swirl to mix the solution and place the flask in a beaker of warm water for

about 10 minutes.

5. In a fume hood, carefully add 50 mL of ice water.

6. Scrape the inside of the flask with a glass stirring rod to hasten the

appearance of crystals.

7. Leave in an ice bath for 5 minutes.

8. Filter off the solid using a Buchner funnel and wash with ice water.

9. Allow to dry overnight and weigh. Record the mass.

10. To test if the reaction was complete, test a small sample of the product with

Fe(NO3)3. A purple colour indicates the presence of unreacted salicylic

acid. Crush an aspirin tablet and test with 1 mol L–1 Fe(NO3)3 to compare

the colour.

EXPERIMENT 8.2




36

QUESTIONS

1. The equation for the reaction is shown below. State the functional groups

present in salicylic acid.

O

OH

OH

C

O

O

C

O

C

H3C CH

3

acetic anhydride

(C4H

6O

3)

salicylic acid

(C7H

6O

3)

O

OH

C

H3C

+ +

O

O

C

CH3

O

OH

C

acetylsalicylic acid

(C9H

8O

4)

2. Name the other product formed apart from the aspirin (acetylsalicylic

acid).

3. What is the purpose of the sulfuric acid?

4. The density of acetic anhydride is 1.08 g mL–1.What is the limiting reactant

in this reaction?

5. Why was the water added at the end of the reaction?

6. Calculate the theoretical yield of aspirin.

7. Calculate the percentage yield of aspirin.

8. From your test with Fe(NO3)3, comment on the purity of your sample.

9. What would be the effect on the calculated result if the filter paper were

still damp?

EXPERIMENT 8.2


37


Reaction rates

In this experiment, the effect of concentration, temperature and catalysts on

reaction rates will be investigated.

The reaction to be studied is the reaction between iodide ions and

peroxydisulfate ions, which proceeds according to the following equation:

2I–(aq) + S2O82–(aq) I2(aq) + 2SO4

2–(aq)

Besides the above chemicals, the reaction mixture will also contain a certain

amount of thiosulfate ions (S2O32–) and some starch.

As the iodine is produced by the above reaction, it is immediately removed

by reaction with the thiosulfate that has been added. However, there will come

a stage where enough iodine has been produced to use up all the thiosulfate that

is present. Further production of iodine, from the equation above, then produces

a blue colour due to this new iodine reacting with the starch. Therefore, by

timing the appearance of this blue colour, we are effectively measuring the time

taken to produce a certain amount of iodine.

AIMTo investigate the effect of concentration, temperature and a catalyst on the rate

of a chemical reaction

APPARATUSFour solutions made up and labelled as follows:

solution 1 — 0.10 mol L–1 potassium iodide solution with some added starch

solution 2 — 0.0025 mol L–1 sodium thiosulfate solution

solution 3 — 0.10 mol L–1 ammonium peroxydisulfate solution

solution 4 — catalyst solution containing 0.01 mol L–1 copper(II) sulfate

10 mL and 100 mL measuring cylinders

250 mL beaker

100 mL volumetric flask

glass stirring rod

thermometer

timing device

METHOD

Part A: Investigating concentrationIn the following steps, it is important to measure all volumes accurately, and to

make certain that all glassware is rinsed between uses.

1. Place 20 mL of solution 1 into a beaker.

2. Add 10 mL of solution 2 to the same beaker.

3. Prepare 20 mL of solution 3, but do not add it to the beaker at this stage.

4. When ready, add solution 3 to the beaker and immediately start stirring and

timing.

5. Record the time taken for the first sign of blue colour to appear.

Change step 1 to 10 mL of solution 1 and 10 mL of water. Repeat steps 2–5

above.

Part B: Investigating temperature

1. Place 20 mL of solution 1 into a conical flask and then warm it until it is

about 30°C above room temperature.

2. Perform steps 2–5 as described in part A.

EXPERIMENT 9.1




38

Part C: Investigating the effect of catalysts

1. Perform steps 1–3 as described in part A.

2. Before doing step 4, add four drops of the copper(II) sulfate solution to the

measured amount of solution 3.

3. Carry out steps 4 and 5 as described in part A.

QUESTIONS

1. Which chemical has its concentration altered during the second and

third repetitions of part A? What can you conclude about the effect of

concentration on the rate of this reaction?

2. Compare your results for the first reaction of part A with your results for

part B. What can you conclude about the effect of temperature on the rate

of this reaction?

3. Compare your results for the first reaction of part A with your results for

part C. What can you conclude about the effect of a catalyst on the rate of

this reaction?

4. Make a summary table of your results.

EXTENSIONIn heterogeneous reactions, another factor that can affect the rate of the reaction

is surface area.

Design and carry out an experiment to investigate the effect of surface area

on reaction rate. You will need to use different chemicals to those used in this

experiment.

EXPERIMENT 9.1


39


Investigating changes to the position of an equilibrium

This experiment investigates the effect of making changes to a mixture that is

at equilibrium. The progress can be monitored by observing the colour of the

mixture, which, under the conditions of the experiment, can be thought of as

being entirely due to the Fe(SCN)2+ ions present.

By viewing down the test tubes to make the comparisons, the relative amounts

of the coloured Fe(SCN)2+ ions may be inferred.

AIMTo observe the effect of a number of changes on the position of the equilibrium

represented by the equation:

Fe3+(aq) + SCN–(aq) s Fe(SCN)2+(aq)

APPARATUSsemi-micro test tubes and holder

a solution of 5 × 10 – 4 mol L–1 Fe(SCN)2+

0.1 mol L–1 solutions of Fe(NO3)3, KSCN, NaF and AgNO3, preferably in

small dropping bottles

white tile

LABORATORY PREPARATION NOTESThe Fe(SCN)2+ solution may be made by adding 20 mL of 0.1 mol L–1 Fe(NO3)3

to 20 mL of 0.1 mol L–1 KSCN and making up to 1 L.

METHOD

1. Obtain six semi-micro test tubes and a test-tube rack to place them in.

2. Place 20 drops of the Fe(SCN)2+ solution into each test tube.

3. Keeping the fi rst test tube as a reference:

(a) place two drops of Fe3+ solution in the second test tube

(b) place two drops of SCN– solution in the third test tube

(c) place two drops of Ag+ solution in the fourth test tube. This reacts with

the SCN– ions to produce a white precipitate of AgSCN.

(d) place two drops of F– solution in the fi fth test tube. The fl uoride ions

react with the Fe3+ ions and convert them into a colourless product.

(e) add enough water to the remaining test tube to double the volume of

its contents.

4. By viewing down each test tube, using a white tile for the background,

compare the colour in each of the fi ve altered mixtures to that of the

reference mixture. In particular, note whether each mixture is darker or

lighter than the original.

QUESTIONS

1. For each of the fi ve altered test tubes, state:

(a) which component of the equilibrium mixture has been initially affected

by the addition of the extra material

(b) how its amount is thus initially altered

(c) how the amount of Fe(SCN)2+ changes as a result of this addition, and

(d) the direction (left or right) in which the reaction must have moved in

order to restore equilibrium.

NaF solution is extremely

poisonous, handle it with care

and wash hands thoroughly at

the end of the lesson.

Other reagents can stain skin

and clothes.




40

2. Explain each of the observations for each test tube in terms of Le Chatelier’s

principle.

3. Draw concentration/time graphs that show the relative changes in

concentration of each species during each of the performed tests.

EXPERIMENT 9.2


41


Rate of hydrogen production — a problem-solving exercise

In this experiment, you are to design a procedure that will investigate the rate

at which hydrogen is produced from the following reaction:

Zn(s) + 2HCl(aq) ZnCl2(aq) + H2(g)

You should be able to produce a design that will investigate the effect of

temperature, surface area and concentration.

CONSIDERATIONS If this experiment is carried out in an open beaker, the mass of its contents

will decrease. Why?

The rate at which the hydrogen is evolved can therefore be monitored by

simply carrying out the reaction in a beaker that is placed on an electronic

balance.

•

•

EXPERIMENT 9.3



42


Temperature and the equilibrium constant

Nitrogen dioxide gas dimerises according to the equation:

2NO2(g) s N2O4(g)

This reaction is exothermic.

The proportion of NO2 in this equilibrium is easily monitored by observing

the colour of the mixture. The NO2 is brown, whereas N2O4 is colourless.

AIMTo observe the effect of temperature on the value of an equilibrium constant

APPARATUS3 test tubes with stoppers

hot water

ice water

500 mL beakers

NO2/N2O4 mixture in stoppered test tubes

LABORATORY PREPARATION NOTESPrior to the experiment, the test tubes should be fi lled with NO2/N2O4 mixture

to the same shade and then stoppered tightly to avoid any leakage. The NO2

may be generated in the apparatus shown below and the test tubes subsequently

fi lled by upward displacement of air.

14 mol L–1 nitric acid

Glass deliverytube (near base

of gas jar)

Copper metal

METHOD

1. Keeping one test tube as a reference, place a second test tube of the gas

mixture in a beaker of ice water and place the third test tube in a beaker of

hot water.

2. After several minutes, compare the colours in the three test tubes.

QUESTIONFrom your observations, what can you infer about the value of the equilibrium

constant at the higher temperature, compared to its value at the lower

temperature?

EXPERIMENT 9.4

NO2 is a poisonous gas. Check

equipment to avoid any leaks,

and perform the experiment in

either a well-ventilated room or in

a fume cupboard.


43


Modelling an equilibrium

This activity uses reaction rate data to model the approach to equilibrium for a

simple reaction. The reaction is:

A(aq) + B(aq) s C(aq) + D(aq)

In each investigation, the initial quantities of A and B (or C and D) are equal

and set at 1.00 mol L–1. Consequently, the rate data are presented in terms of A

and C only; the concentration of B always equals the concentration of A, and

the concentration of D always equals the concentration of C.

AIMTo use a computer spreadsheet application and reaction rate data to model a

chemical reaction and to determine the effect of temperature on the equilibrium

constant for the reaction

METHODUse a spreadsheet application such as Excel for this activity. There are four

spreadsheeting tasks to perform. Go to www.jaconline.com.au/ict-me to see

instructions for tabulating the data in the spreadsheet and to create charts for

analysis.

Tasks 1 and 2 represent the system at the same temperature (25°C). In task

1, A and B are added to the vessel and no C and D are present initially. In task

2, C and D are added and no A and B are present initially. These two tasks will

allow you to examine the approach to equilibrium from opposite directions.

Tasks 3 and 4 involve A and B reacting at higher temperatures.

TASK 1 (A REACTS TO FORM C AT 25°C.)

1. Copy the data from the table below into the spreadsheet.

TASK 1 Concentration data (mol L–1) at 25°C

t (min) 0 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20

[A] 1.00 0.92 0.84 0.77 0.71 0.65 0.59 0.54 0.50 0.46 0.43 0.41 0.40 0.40 0.40 0.40 0.40 0.40 0.40 0.40 0.40

[C] 0.00 0.08 0.16 0.23 0.29 0.35 0.41 0.46 0.50 0.54 0.57 0.59 0.60 0.60 0.60 0.60 0.60 0.60 0.60 0.60 0.60

2. Use the instructions at www.jaconline.com.au/ict-me to produce an XY

(scatter) graph of the data in the table above, showing the data points and

a line of best fit. Make sure your graph has a title, an appropriate scale on

each axis and the appropriate units in the axis titles. Incorporate a legend to

show which of the two curves represents [A] and which represents [C ]. Save

this file as ‘chart 1’.

TASKS 2, 3 AND 4The tables below and over the page provide the data for tasks 2, 3 and 4.

Proceed as for task 1 and generate three more charts. Save these files as ‘chart 2’,

‘chart 3’ and ‘chart 4’.


t (min) 0 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20

[C] 1.00 0.94 0.88 0.82 0.78 0.73 0.70 0.67 0.65 0.63 0.62 0.61 0.60 0.60 0.60 0.60 0.60 0.60 0.60 0.60 0.60

[A] 0.00 0.06 0.12 0.18 0.22 0.27 0.30 0.33 0.35 0.37 0.38 0.39 0.40 0.40 0.40 0.40 0.40 0.40 0.40 0.40 0.40

EXPERIMENT 9.5




44


t (min) 0 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20

[A] 1.00 0.85 0.72 0.62 0.53 0.45 0.39 0.34 0.31 0.30 0.30 0.30 0.30 0.30 0.30 0.30 0.30 0.30 0.30 0.30 0.30

[C] 0.00 0.15 0.28 0.38 0.47 0.54 0.61 0.66 0.69 0.70 0.70 0.70 0.70 0.70 0.70 0.70 0.70 0.70 0.70 0.70 0.70


t (min) 0 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20

[A] 1.00 0.80 0.63 0.48 0.34 0.24 0.20 0.20 0.20 0.20 0.20 0.20 0.20 0.20 0.20 0.20 0.20 0.20 0.20 0.20 0.20

[C] 0.00 0.20 0.37 0.52 0.66 0.76 0.80 0.80 0.80 0.80 0.80 0.80 0.80 0.80 0.80 0.80 0.80 0.80 0.80 0.80 0.80

EXPERIMENT 9.5

QUESTIONS

1. Compare the equilibrium concentrations of A and C in charts 1 and 2.

2. State the equilibrium concentrations of A, B, C and D at 25°C.

3. Write the equilibrium constant expression for this reaction.

4. Calculate the value of the equilibrium constant at 25°C.

5. Use charts 1, 3 and 4 to compare the equilibrium concentrations of A and

C as the temperature increases.

6. Describe how the position of the equilibrium changes as the temperature

increases.

7. Classify the forward reaction as endothermic or exothermic.

8. Calculate the equilibrium constant, K, for each temperature investigated.

9. Explain how the charts show that equilibrium is attained faster as the

temperature increases.


45


The acidity constant of acetic acid

The strength of an acid is a measure of the extent to which it reacts with water

to produce H3O+ ions. Quantitatively, the degree of such a reaction is described

by an acid’s ‘acidity constant’, Ka, which is just a special kind of equilibrium

constant.

In this experiment, you will determine the value of the acidity constant for

acetic acid. The reaction of acetic acid with water is given by the equation:

CH3COOH(l) + H2O(l) s CH3COO–(aq) + H3O+(aq)

The acidity constant is defined by the expression:

Ka3

–3

+

3

=[CH COO ][H O ]

[CH COOH]

AIMTo determine the acidity constant for acetic acid

APPARATUSelectronic pH meter*

distilled water**

0.1 mol L–1 acetic acid solution

10 mL and 100 mL measuring cylinders

dropping pipettes

two 10 mL beakers

* The pH meter may be replaced by universal indicator solution.

** Distilled water may have a pH as low as 4, owing to dissolved CO2. This

may be removed by boiling the distilled water (with a few boiling chips)

and then allowing it to cool while covered with plastic film.

LABORATORY PREPARATION NOTESThe accuracy of the pH meter will need to be checked before the class and

recalibrated against standard buffers if necessary.

It is important that the glassware in this experiment be as clean as possible.

METHOD

1. Place about 50 mL of the supplied 0.1 mol L–1 acetic acid solution into one

of the 100 mL beakers. Measure and record the pH of this solution.

2. Using the 10 mL measuring cylinder, measure out 1 mL of the acetic acid

solution from step 1.

3. Transfer this to a 100 mL measuring cylinder. Rinse the 10 mL measuring

cylinder with distilled water at least twice, adding these washings to the

larger cylinder. Make up the volume in the larger cylinder to 100 mL with

more distilled water. Mix thoroughly.

4. Transfer about 50 mL of this diluted solution into the second 100 mL

beaker and measure the pH as in step 1.

CALCULATIONSWe can assume that the contribution to the total [H3O

+] from the ionisation of

water is so small as to be negligible. It follows that the acid will be the only

source of H3O+. If we call this concentration x, it follows from the stoichiometry

of the above reaction that the concentration of CH3COO– ions will also be x.

EXPERIMENT 10.1




46

If we further assume that the amount of acetic acid undergoing hydrolysis is

small (a reasonable assumption for a weak acid), the equilibrium concentration

may be approximated to the initial concentration (i.e. that stated ‘on the bottle’).

Suppose we call this value c.

The above equilibrium expression thus becomes:

Ka = x

c

2

The concentration of the H3O+ ions at equilibrium is calculated from the

measured pH value. (Ask your teacher how to do this if you are uncertain.)

Calculate the Ka value for both acetic acid solutions that you used in this

experiment.

QUESTIONS

1. The accepted Ka value for acetic acid (at 25°C) is 1.74 × 10–5. Use this to

calculate the percentage error in your average result.

2. Would this method of calculation be suitable for a strong acid? Explain.

EXPERIMENT 10.1


47


Buffers

A buffer is a solution made from a weak acid and the salt of its conjugate base.

Alternatively, a weak base and the salt of its conjugate acid may be used.

Buffers can maintain a reasonably constant pH despite additions of small

amounts of acid or base. It is this effect that will be demonstrated in this

experiment.

AIMTo prepare a number of buffer solutions and to demonstrate their ability to

maintain a relatively constant pH when compared to unbuffered solutions


test-tube rack

universal indicator solution


2 × 100 mL beakers

samples of the following solutions (all of 0.1 mol L–1 concentration): Na2CO3,

NaHCO3, Na2HPO4, NaH2PO4

freshly boiled and cooled distilled water

solutions of 1 mol L–1 HCl and 1 mol L–1 NaOH

METHOD

1. Label the seven test tubes with the numbers 1–7.

2. Measure 10 mL of the 0.1 mol L–1 Na2CO3 solution and 10 mL of the

0.1 mol L–1 NaHCO3 solution into a 100 mL beaker. Mix thoroughly and

divide this evenly into test tubes 1 and 2.

3. Repeat step 2 using the 0.1 mol L–1 Na2HPO4 solution and the NaH2PO4

solution. Divide this evenly into test tubes 3 and 4.

4. Place 10 mL of 0.1 mol L–1 NaHCO3 solution in test tube number 5, 10 mL

of 0.1 mol L–1 NaH2PO4 solution in test tube number 6, and 10 mL of

distilled water into test tube number 7.

5. Add a few drops of universal indicator to each test tube and estimate the

pH.

6. Add 2 drops of the 0.1 mol L–1 HCl to test tubes 1, 3, 5, 6 and 7. Mix

thoroughly and re-estimate the pH values for each test tube.

7. Add 4 drops of the 0.1 mol L–1 NaOH solution to the same test tubes as in

step 6. Estimate the new pH values.

QUESTIONS

1. Write the ionic equation for what happens when an acid is added to the

buffer system present in test tubes 1 and 2.

2. Write the ionic equation for the reaction that occurs when a base is added

to the buffer system present in test tubes 1 and 2.

The hydrochloric acid and the

sodium hydroxide solutions

should be diluted with water if

you spill them on your skin or

clothes.

Wear safety glasses to avoid

splashing these chemicals into

your eyes. If this happens, the

eye should be fl ushed with water

for at least 15 minutes.

The hydrochloric acid and the

sodium hydroxide solutions

should be diluted with water if

you spill them on your skin or

clothes.

Wear safety glasses to avoid

splashing these chemicals into

your eyes. If this happens, the

eye should be fl ushed with water

for at least 15 minutes.

EXPERIMENT 10.2



48


Catalytic oxidation of ammonia

This experiment demonstrates the catalytic oxidation of ammonia. This is the

important fi rst step of the Ostwald process.

AIMTo observe the catalytic oxidation of ammonia using a number of metals as

catalysts

APPARATUS300 mL conical fl ask

spatula

concentrated ammonia solution (fresh)

samples of wire made from a number of metals including copper and platinum

(if available)

METHOD

1. Place concentrated ammonia solution to a depth of 1 cm in the conical

fl ask.

2. Choosing one of the wires, wind the end of it into a tight spiral. Adjust the

length of this wire so that, when suspended from a spatula laid across the

top of the fl ask, the tip is just above the surface of the ammonia, as shown

below.

Spatula

Copper wire formed

into a spiral at its end

Concentratedammonia solution

°

°

°

3. Heat the wire until it is red hot and then quickly suspend it from the

spatula.

4. Note what happens and then repeat the experiment for wires made from

other metals.

QUESTIONS

1. What is the equation for the catalytic oxidation of ammonia?

2. Which metal worked best at catalysing this reaction?

3. Is this process exothermic or endothermic? Discuss the evidence for your

answer.

Concentrated ammonia is a

respiratory and eye irritant.

Ensure adequate ventilation

in the room, or alternatively,

perform the experiment in a fume

cupboard.

Safety glasses should be worn.


49


Cracking of paraffi n oil

Cracking is the process of converting large alkane molecules into a mixture of

smaller alkanes and alkenes. It can be done using either high temperatures or

a catalyst. In this experiment, paraffi n oil, which is a mixture of large alkane

molecules, is heated to bring about cracking. The products are then tested to

establish their chemical nature.

AIMTo convert a sample of paraffi n oil into a mixture of smaller saturated and

unsaturated hydrocarbons

APPARATUSPyrex test tube fi tted with a stopper and delivery tube

4 further test tubes with stoppers

pneumatic trough

stand and clamp

small pieces of broken porcelain

glass wool

wax tapers

paraffi n oil

bromine solution

Bunsen burner

METHOD

1. Place glass wool in the bottom of the Pyrex test tube to a depth of about

3 cm. Add about 5 mL of the paraffi n oil.

2. Fill the rest of the test tube with the pieces of broken porcelain.

3. Assemble the equipment as shown below. Make sure that the test tube to

be heated is almost horizontal. The remaining test tubes can be laid down,

fi lled with water, in the pneumatic trough.

Glass wool

and paraffin oilPorcelain chips

Test tube

Further test tube

Delivery tube

Pneumatic

trough

Bunsen

burner

4. Heat the centre of the test tube until it begins to glow red hot. Then move

the burner fl ame down to the glass wool for a few seconds, then back up to

the centre. Continue to heat in this way.

5. Collect the gas evolved in the test tubes, changing them as each one fi lls.

Number these test tubes in the order in which they are fi lled.

6. When fi nished, turn off the Bunsen burner and remove the delivery tube

immediately from the water.

Wear safety goggles. The

equipment gets extremely hot.

Flammable gases are produced.

Be careful when using bromine

solution.

EXPERIMENT 11.2

7. Add a little bromine solution to one of the test tubes. Stopper and shake.

Note any colour change.

8. Remove the stopper from another test tube and quickly put a lighted taper

to the mouth of the tube. Describe your observations.

QUESTIONS

1. Is there any evidence that alkenes have been produced? Explain.

2. What evidence is there to suggest that smaller molecules have been formed

in this experiment?




50

EXPERIMENT 11.2


51


Fractional distillation of ethanol/water mixture

When distillation is used to separate two or more types of liquids, fractional

distillation is used. As the name implies, this type of distillation splits the

mixture into different parts (fractions). The method works because the liquids

boil at different temperatures. For example, ethanol boils at 78°C so it vaporises

more easily than water. In this experiment, you will carry out a simple fractional

distillation experiment to separate a mixture of ethanol and water.

AIMTo separate a mixture of ethanol and water by fractional distillation

APPARATUShot plate

round-bottomed fl ask

fractionating column

thermometer

condenser

glass adapters

250 mL conical fl ask

anti-bumping granules

absolute ethanol

water

Fractionatingcolumn

Anti-bumpinggranules

ClampCold water

Hot plateset at 80oC

METHOD

1. Make up a mixture of 50 mL of absolute ethanol and 50 mL of water in the

round-bottomed fl ask.

2. Add a spatula full of anti-bumping granules.

3. Set up the apparatus as shown in the fi gure above.

4. Set the thermostat on the hot plate to around 80°C.

5. Start the heating.

6. Observe the temperature when the fi rst drop of ethanol drips into the conical

fl ask.

EXPERIMENT 11.3

Ethanol is a volatile liquid and

is fl ammable.




52

7. Collect the ethanol (about 40–50 mL).

8. Identify that the liquid you have collected is mainly made up of ethanol.

9. Raise the temperature of the hot plate to 100°C and continue heating.

QUESTIONS

1. At what temperature did the ethanol vapour start condensing into liquid?

2. How would you identify that the first fraction of liquid that is collected in

the flask is ethanol?

3. Was there a period of time where no liquid dripped into the conical flask

(i.e. between the alcohol and water fractions)?

4. Name two uses for ethanol.

EXPERIMENT 11.3


53


Catalytic dehydration (cracking) of ethanol

Cracking is a process used in oil refi neries to split large hydrocarbon molecules

into smaller ones. The reaction takes place in the gas phase and involves heat

and catalysts. In this experiment, you will ‘crack’ ethanol to produce ethene

gas. The reverse reaction, the conversion of ethene to ethanol, is the industrial

method of producing ethanol on a large scale. This process involves mixing

ethene with steam and passing it over a phosphoric acid catalyst at about

300°C.

300°Cethene + steam ethanol

phosphoric acid

AIMTo produce ethene gas by the process of catalytic dehydration

APPARATUS2 large test tubes

rubber bung with hole in the centre

glass tubing

mineral wool

ethanol

aluminium oxide

Bunsen burner

METHOD

1. Set up the apparatus as shown in the fi gure below.

Heat

Water

Aluminium oxide

Mineral

wool soaked

in ethanol

2. Heat the test tube containing ethanol and aluminium oxide, using the blue

fl ame of the Bunsen burner.

3. Collect a tube full of ethene gas.

4. Cover the mouth of the test tube and place in a rack in a fume cupboard.

5. Test that the gas burns by quickly dropping a lighted splinter into the test

tube.

QUESTIONS

1. What is the function of aluminium oxide in this experiment?

2. Write a balanced equation for the catalytic cracking of ethanol.

3. What other tests for ethene do you know of ?

4. Why is ethene an important compound to chemists?

EXPERIMENT 11.4

Ethene and ethanol are both

fl ammable substances. This

experiment should be carried out

in a fume cupboard.



54


Constructing models of esters

AIMTo construct molecular models of alcohols, carboxylic acids and esters

APPARATUSmolecular building kit or coloured plasticine and satay sticks or straws cut into

5 cm lengths

METHOD

1. In your record book, draw the structural arrangement for methanol, ethanol,

methanoic acid, ethanoic acid and propanoic acid.

2. Construct the three-dimensional models for these hydrocarbons.

3. In your record book, write the general equation for the reaction between an

alcohol and a carboxylic acid to form an ester. Use R as the alkyl group for

carboxylic acid and R´ for alcohol. Circle the ester linkage.

4. Draw the structural arrangement of the esters formed from:

(a) methanol + propanoic acid

(b) ethanol + ethanoic acid

(c) ethanol + methanoic acid.

In each case circle the ester linkage.

5. Use the models in step 2 to construct the esters that you have drawn up in

step 4. In each case, what molecule was eliminated?

6. Name the esters that you have constructed.

7. What type of reaction is the formation of an ester from an alcohol and a

carboxylic acid?

8. Name two uses of esters.

EXPERIMENT 11.5


55


An esterifi cation experiment

Esters are formed from alcohols and carboxylic acids. In this experiment, you

will carry out an esterifi cation process to produce ethyl ethanoate. The ester

formed is insoluble in water.

AIMTo prepare a small amount of the ester ethyl ethanoate from ethanol and

ethanoic acid

APPARATUStest tube

50 mL and 600 mL beakers

Bunsen burner

hot plate

dropping pipette

retort stand and clamp

glacial acetic acid (ethanoic acid)

ethanol

concentrated sulfuric acid

METHOD

1. Half fi ll the 600 mL beaker with water and place it over a hot plate. Heat

the water to boiling temperature. This will act as a water bath for the

esterifi cation reaction.

2. For safety, do this step in the fume cupboard. To a test tube, add 1 mL

of glacial acetic acid (glacial ethanoic acid) and 2 mL of ethanol. (It is

necessary to use a fume cupboard because glacial acetic acid is a respiratory

irritant.)

3. Still in the fume cupboard, add 3 drops of concentrated sulfuric acid to the

mixture in step 2. Mix well.

4. Heat the mixture in the water bath.

5. After about 5 minutes, remove the test tube from the water bath and pour

the mixture into a beaker of about 20 mL of cold water.

6. Using a dropper, carefully collect the fl oating layer into a small vial or

another test tube.

7. Smell the substance that you have just collected. This is the ester ethyl

ethanoate.

QUESTIONS

1. Do ethanol, ethanoic acid and sulfuric acid mix with water?

2. Does the mixture in question 1 have a sweet, gluey smell?

3. What can you say about the reaction between ethanol and ethanoic acid?

4. Why do you think the reaction mixture is poured into the beaker of cold

water?

5. Describe the smell of the ester that you have made.

6. Draw the structural equation for the reaction. Make sure you circle the ester

linkage in the product.

EXPERIMENT 11.6

Concentrated sulfuric acid is

very corrosive. Handle this

substance with gloves and carry

out the mixing of sulfuric acid in

the fume cupboard.

Ethanol is a volatile liquid and is

fl ammable.

Glacial ethanoic acid is a

respiratory irritant.




56

EXTENSION: PRODUCTION OF OTHER ESTERSUsing the same procedure, make other esters from:

(a) ethanol and butanoic acid

(b) octan–1–ol and ethanoic acid

(c) methanol and butanoic acid.

Identify the odour of the ester in each case.

EXPERIMENT 11.6


57


Investigating heat changes in reactions

AIMTo investigate and draw energy diagrams for some exothermic and endothermic

reactions

APPARATUSevaporating dish

100 mL beaker

test tube

gauze mat

heating apparatus

spatula

thermometer

5 drops concentrated sulfuric acid

3 pellets of solid sodium hydroxide

2 pieces of calcium

2 spatulas of ammonium nitrate

water

1 spatula of solid anhydrous copper(II) sulfate

10 mL hydrogen peroxide1

2spatula manganese dioxide

PART A: DILUTION OF ACIDMETHOD

1. Place 10 mL of water into a test tube, then measure and record the initial

temperature.

2. Carefully add 5 drops of concentrated sulfuric acid into the test tube.

3. Record the change in temperature of the reaction.

QUESTIONS

1. Is the dilution of sulfuric acid exothermic or endothermic? Justify your

decision.

2. Draw an energy diagram for the reaction.

3. Explain why, when diluting acids, the rule is to ‘always add acid to the

water rather than water to the acid’.

PART B: DISSOLVING SODIUM HYDROXIDE IN WATERMETHOD


temperature.

2. Carefully add a small pellet of solid sodium hydroxide to the water. Stir

gently with the thermometer so that the solid dissolves.

3. Feel the outside of the test tube and record the change in temperature of the

reaction.

QUESTIONS

1. Is the reaction NaOH(s) NaOH(aq) endothermic or exothermic? Justify

your decision.

2. Do the reactants or the products have the greatest enthalpy? Explain.

3. Draw an energy diagram for this reaction.

EXPERIMENT 13.1

Wear safety goggles for this

experiment.

Sodium hydroxide is caustic.

Avoid contact with skin.

Concentrated sulfuric acid is

corrosive. Avoid contact with

skin, eyes and clothing.

Copper(II) sulfate is toxic. Handle

with care.




58

PART C: DISSOLVING AMMONIUM NITRATE IN WATERMETHOD


temperature.

2. Carefully add a spatula of solid ammonium nitrate to the water. Stir gently

with the thermometer so that the solid dissolves.


reaction.

QUESTIONS

1. Is the reaction NH4NO3(s) NH4NO3(aq) endothermic or exothermic?

Justify your decision.


PART D: REACTING CALCIUM WITH WATERMETHOD

1. Place two pieces of calcium in a dry test tube.

2. Insert a thermometer so that it touches the pieces of metal and record the

initial temperature.

3. Add eight drops of water and note the change in temperature.

QUESTIONS

1. Is the reaction endothermic or exothermic? Justify your decision.

2. Write an equation for the reaction.


PART E: DISSOLVING ANHYDROUS COPPER(II) SULFATE IN WATERMETHOD


temperature.

2. Add a spatula of solid anhydrous copper(II) sulfate to the water. Stir gently

with the thermometer so that the solid dissolves.


reaction.

QUESTIONS

1. Is the reaction endothermic or exothermic? Justify your decision.

2. Predict the temperature change for the reverse reaction.


PART F: EFFECT OF A CATALYST ON THE DECOMPOSITION OF HYDROGEN PEROXIDEMETHOD

1. Place 10 mL of hydrogen peroxide into each of two test tubes, then measure

and record the initial temperature of the liquid in both test tubes.

2. Add half a spatula of manganese(IV) oxide powder, MnO2, to one of the

test tubes.

3. Record the temperature changes in each test tube every 30 seconds for

8 minutes.

EXPERIMENT 13.1



59

QUESTIONS

1. Is the reaction H2O2(l) H2(g) + O2(g) endothermic or exothermic?

Justify your decision.

2. Manganese dioxide acts as a catalyst for the decomposition of hydrogen

peroxide. What experimental evidence indicates that the catalyst speeds up

the reaction?

3. Draw an energy diagram for this reaction, showing the catalysed and

uncatalysed reactions on the same set of axes.

EXPERIMENT 13.1



60


Calorimetry: Determining the heat of reaction of a simple redox reaction

Zinc displaces copper from copper(II) sulfate solution:

Zn(s) + CuSO4(aq) Cu(s) + ZnSO4(aq)

or

Zn(s) + Cu2+(aq) Cu(s) + Zn2+(aq)

The heat of reaction may be calculated by using a calibrated calorimeter. In the

laboratory, a solution calorimeter can be made from a disposable foam cup (see

the figure below).

AIMTo determine the heat of reaction for zinc metal and copper ions

APPARATUS2 disposable polystyrene cups

10 cm × 10 cm sheet of cardboard

thermometer

6 V DC supply

voltmeter

ammeter

switch

heating coil of 2 ohms resistance

filter paper

0.5 g of zinc powder

100 mL of 0.2 mol L–1 copper(II) sulfate solution

weighing scales

stopwatch

PART A: CALIBRATING THE CALORIMETERMETHOD

1. Add 100 mL of water to the calorimeter and record its temperature.

2. Set up the apparatus for calorimeter calibration as shown in section (a) of

the figure on the next page, which also shows the circuit diagram of the

apparatus in section (b). Use a heating coil of 2 ohms resistance and a 4–6

volt DC power supply.

3. Switch on the current and commence timing.

4. Record the current and the potential difference while the water is heating.

EXPERIMENT 13.2



61

5. Switch off the current after exactly three minutes, stir the water and record

the highest temperature reached.

QUESTIONCalculate the calibration factor for your calorimeter.

CalorimeterSwitch

Thermometer

VoltmeterAmmeter6 V DCsupply

6 V

A

X YV

(a)

(b)

+–

+

+

–

–

PART B: DETERMINING THE HEAT OF REACTION OF A SIMPLE REDOX REACTIONMETHOD

1. Place 100 mL of 0.2 mol L–1 copper(II) sulfate solution in the calibrated

polystyrene cup calorimeter.

2. Weigh out approximately 2 g of zinc powder on a piece of filter paper, then

tip the powder into the calorimeter.

3. Immediately record the initial temperature of the solution.

4. Record the temperature of the solution when the reaction is complete (at

this point, the blue colour of the solution will completely fade).

5. Calculate the heat of reaction, then write a thermochemical equation for the

reaction between zinc metal and copper ions.

QUESTIONS

1. Is the reaction exothermic or endothermic?

2. Why does the blue colour of the copper(II) sulfate solution disappear?

3. In the experiment, an excess of zinc powder was used. Why was this

done?

4. Draw an energy-level diagram for the reaction.

5. List the sources of error in this experiment.

6. How would you expect these sources of error to affect the value you obtain

for the heat of reaction for the solution?

7. Outline modifications to the experimental procedure that would allow you

to determine the heat of reaction more accurately.

EXPERIMENT 13.2



62


The energy content of a peanut

Biochemists often measure the energy value of a food by burning a sample of

the food under controlled conditions in a bomb calorimeter. This process may

be simulated in the laboratory using simpler, though less accurate, equipment.

The energy of the food being tested may be calculated by assuming that the

specific heat of water is 4.184 J g –1 °C–1.

Sample calculation

A 5 g sample of breakfast cereal was burnt in a laboratory calorimeter. The

complete combustion of the cereal raised the temperature of 100 mL of water

from 19°C to 28°C. Calculate the energy content of the cereal.

Solution

100 mL of water corresponds to 100 g of water. If 4.184 J of energy will raise

1 g of water by 1°C, then 4.148 × 100 × (28°C – 19°C) joules of energy is

provided by the breakfast cereal.

AIMTo measure the energy content of a peanut by calorimetry, and to compare the

experimental value with published energy values

APPARATUSpeanut

10 mL water

test tube

retort stand

utility clamp

cork

needle

thermometer

measuring cylinder

matches

METHOD

1. Pour 10 mL of water into a test tube. Measure and record the temperature

of the water.

2. Accurately weigh a peanut. Record its mass.

3. Set up the apparatus shown in the figure on the following page.

4. Set the peanut alight and allow it to burn completely.

5. Record the final temperature of the water when the peanut has been

completely burnt.

6. Assuming that the specific heat of water is 4.184 J g –1 °C–1, calculate the

heat content of the peanut.

EXPERIMENT 13.3



63

Test tube

10 mL water

Needle

Peanut

Cork

Retort

stand

Bosshead

and

clamp

QUESTIONS

1. Compare the value you obtained for the heat content of a peanut in your

experiment with the expected heat of 8 kJ. How accurate are your results?

2. List the sources of error in your experiment.

EXPERIMENT 13.3



64


Investigating the Daniell cell

If a reductant and oxidant are reacted directly, the energy is released as heat.

Although some of this heat can be converted into electrical energy, a more

efficient conversion of chemical energy to electrical energy may be achieved if

the reaction is set up as a galvanic cell.

In a galvanic cell, the oxidant and reductant are separated. The cells are

connected through an external circuit via electrodes and connecting wires, and

through an internal circuit via a salt bridge. Electrons flow from the reductant to

the oxidant through the external circuit and there is a complementary movement

of ions in the salt bridge (see figure below).

The Daniell cell is one of the earliest types of galvanic cell constructed in

the laboratory. In this cell, zinc metal is the reductant and copper(II) ions act

as oxidants.

+–

Cathode

Electron flow

Salt bridgeAnions Cations

Anode

e– e–

V

Reductant is oxidisedin this compartment

Oxidant is reduced inthis compartment

AIMTo set up, and observe the operation of, a Daniell cell

APPARATUS2 × 100 mL beakers

Petri dish

emery paper

a 10 cm × 1 cm strip of filter paper

a voltmeter

connecting wires

60 mL 1 mol L–1 copper(II) sulfate solution

60 mL 1 mol L–1 zinc sulfate solution

20 mL potassium chloride solution

10 cm strip of copper

10 cm strip of zinc

EXPERIMENT 14.1



65

METHOD

Voltmeter

Filter paper strip soaked in KCI(aq)

Electrical wire

Alligator clip

CuSO4(aq) ZnSO

4(aq)

Copper

strip

Zinc strip

1. Clean the metal strips with emery paper.

2. Add the copper(II) sulfate solution to one beaker and the zinc sulfate

solution to the other beaker (see figure above).

3. Soak the filter paper in a Petri dish containing potassium chloride solution,

then form a ‘bridge’ into both solutions.

4. Place the zinc strip in the beaker containing zinc sulfate and the copper

strip in the beaker containing copper(II) sulfate. Connect with leads to a

voltmeter.

5. Allow the reaction to proceed for 20 minutes and record all observations,

including the voltmeter readings at 5-minute intervals.

QUESTIONS

1. Why should the metal strips be cleaned with emery paper?

2. Why can’t steel wool be used to clean the metal strips?

3. What purpose does the filter paper soaked in potassium chloride solution

serve?

4. What happens to the voltmeter reading when the filter paper is taken out?

5. Describe the reactions that occurred in each half-cell and then write the

half-cell reaction.

6. Write the overall ionic equation for the Daniell cell.

7. Write the shorthand symbol for the Daniell cell.

8. Draw a diagram of the Daniell cell, labelling the anode, cathode and the

direction of electron flow.

EXTENSIONDraw a graph of the voltmeter readings obtained versus time. Extend the time

axis to 24 hours, and predict expected voltage.

EXPERIMENT 14.1



66


Galvanic cells and redox potentials

Simple galvanic cells can be set up in the laboratory by preparing half-cells

in test tubes and using a salt bridge that is constructed from a sheet of filter

paper soaked in electrolyte solution. The half-cells may then be connected to

a voltmeter.

Connecting

wires

Test

tube

Test-tube rack

Voltmeter

Filter paper

soaked in

KCI

Alligator

clip

Metal

strip

V

AIMTo set up galvanic cells, measure the cell voltages and predict the relative

oxidising–reducing strength of four redox pairs


test-tube rack

filter paper

Petri dish half-filled with 1 mol L–1 potassium chloride, KCl, solution

electrical leads

voltmeter

20 mL each of 1 mol L–1 ZnCl2, CuSO4, FeSO4 and Pb(NO3)2

10 cm strips of zinc, copper, iron and lead

METHOD

1. Copy the table.

1st half-cell

components

2nd half-cell

components

Predicted

observation

Actual

observation

Zn2+/Zn Cu2+/Cu

Zn2+/Zn Fe2+/Fe

Zn2+/Zn Pb2+/Pb

Cu2+/Cu Fe2+/Fe

Cu2+/Cu Pb2+/Pb

Fe2+/Fe Pb2+/Pb

EXPERIMENT 14.2



67

2. Use a table of standard electrode potentials to predict whether a spontaneous

redox reaction will occur, and determine the expected cell voltage for each

of the six reactions in the table in step 1. Fill in the third column of the

table.

3. Place the solutions of ZnCl2, CuSO4, FeSO4 and Pb(NO3)2 in separate test

tubes and add the corresponding metal strips to the solution. Each of these

four test tubes forms a half-cell.

4. Soak a strip of filter paper in the Petri dish containing the potassium

chloride solution. This will form your salt bridge.

5. Connect the first two half-cells listed in the table with the salt bridge.

6. Connect the metal strips from each half-cell to the voltmeter, ensuring that

the strips maintain contact with the solution in the test tubes.

7. Allow the reaction to occur for a few minutes, then note the potential

difference reading and record it in the fourth column of the table.

8. Observe what happens in each half-cell and record the results in the fourth

column of the table.

9. Repeat steps 5–8 for the remaining five redox reactions, using a new salt

bridge for each different reaction.

QUESTIONS

1. What would happen if two Cu2+/Cu half-cells were connected to form a

cell? Explain your answer.

2. Write equations for any reactions that occurred.

3. How did your predicted results compare with your actual results? Explain

any differences.

4. From your observations, list the oxidants tested in decreasing order of

strength, explaining the reasons for your selected order. How does your

order compare with that in a standard electrode potential table?

EXPERIMENT 14.2



68


Looking at a dry cell

The Leclanché dry cell is the most common source of small-scale, portable

electrical energy. The electrical potential of this cell usually begins at 1.5 V

when new, but falls steadily during use to about 0.8 V. Dry cells of this

construction are not rechargeable.

AIMTo examine the contents of a dry cell and to investigate the redox reactions

occurring in the cell

APPARATUSdry cell (discharged if possible) with a cut through the zinc casing around

the circumference so that the two halves can be pulled apart (see the figure

below)

spatula

sheet of newspaper

graduated 100 mL beaker

zinc strip (5 cm × 2 cm)

carbon electrode (may be removed from cell)

2 leads with alligator clips

0–3 V voltmeter

glass stirring rod

emery paper

25 mL saturated NH4Cl solution

25 mL 1 mol L–1 ZnCl2 solution

universal indicator solution

– M

AD

E I

N A

US

TR

AL

IA +

PART A: EXAMINATION OF A COMMERCIAL DRY CELLMETHOD

1. Sketch the outside of a complete dry cell, labelling the positive and negative

ends.

EXPERIMENT 15.1



69

2. Pull apart a pre-cut dry cell (latitudinal section), then sketch and label the internal parts.

3. Place a sheet of newspaper on your bench and carefully scrape out the black powder from the bottom half of the cell with a spatula. Look carefully at the base of the cell and sketch what you see. Suggest an explanation for the arrangement at the base of the cell.

4. Repeat step 3 with the top half of the cell and sketch what you see. Suggest an explanation for the arrangement at the top of the cell.

5. From the diagrams you have already sketched, suggest the structure for a longitudinal section of the cell. Draw your proposed structure, labelling all features. Compare your diagram with those of other members of your class.

6. Remove the carbon electrode from the cell and retain it for part B of the experiment. Keep some of the cell contents aside if you intend to perform the extension exercise. Wrap the remainder of the cell in the newspaper and discard it into a rubbish tin.

PART B: DRY CELL REACTIONSMETHOD

1. Thoroughly clean a strip of zinc with emery paper.

2. Pour 25 mL of saturated ammonium chloride solution and 25 mL of 1 mol L–1 zinc chloride solution into a 100 mL beaker. Add 10 drops of universal indicator solution to the mixture and stir.

3. Place the zinc strip on one side of the beaker and a carbon electrode on the other side of the beaker. Connect the zinc strip and the carbon rod with a voltmeter. Record the voltage generated and the polarity of the two electrodes.

4. Disconnect the voltmeter, then connect the two electrodes using a lead. Allow the cell to discharge for 20 minutes.

5. Examine the cell, paying particular attention to the area around each

electrode. Record all observations.

QUESTIONS

1. Why wasn’t a salt bridge necessary in part B of the experiment?

2. Explain the observations made at each electrode in terms of the reactants being used and the products being formed.

3. Identify the reaction occurring at the anode. Write a half-equation for this reaction.

4. Identify the reaction occurring at the cathode. Write a half-equation for this reaction.

5. Write an overall cell reaction.

EXTENSIONThe electrolytic paste in a dry cell may contain ammonium ions, NH4

+, or chloride ions, Cl–. These may be detected by collecting a small sample of the cell contents and performing the following tests.

1. Test for ammonium ions. Add sufficient water to a sample of the cell contents to form a paste. Stir in a small amount of Ca(OH)2 and then gently warm the mixture. If ammonium ions, NH4

+, are present, ammonia gas will be evolved. This can be identified by its characteristic odour, or by passing red litmus paper through the gas.

2. Test for chloride ions. Shake a sample of the cell contents with water, then allow the mixture to settle. Decant the supernatant liquid into a test tube and add 0.1 mol L–1 AgNO3. Chloride ions will cause a white precipitate,

AgCl, to form.

EXPERIMENT 15.1



70


Lead storage batteries

Lead–acid accumulators are sometimes called storage batteries because they

contain chemicals from which electricity may be produced via a spontaneous

redox reaction.

AIMTo simulate the chemistry of a ‘car battery’

APPARATUS100 mL beaker

DC power source

electrical leads with clips

alligator clips

voltmeter

50 mL of 2 mol L–1 sulfuric acid

2 strips of lead foil (each 1 cm × 8 cm)

METHOD

1. Attach two strips of lead foil to opposite sides of a 100 mL beaker with the

alligator clips, then carefully pour 50 mL of 2 mol L–1 sulfuric acid into the

beaker.

2. Connect strips to a voltmeter and take an initial reading.

3. Connect a 6–12 volt DC power supply into the circuit and allow the current

to fl ow for 20 minutes. Observe the cell, particularly the areas around the

electrodes, and record your results.

4. Disconnect the power source and connect a voltmeter to the cell. Record

the potential difference obtained.

5. Collect other students’ cells and connect each in series (positive to negative,

positive to negative, etc.). Connect the two extreme terminals to a voltmeter.

Record the number of cells used and the potential difference generated.

QUESTIONS

1. In what ways does the cell you constructed resemble a lead–acid cell?

2. In what ways does the cell you constructed differ from a lead–acid cell?

3. What are the requirements for a battery that is to be used in a car?

4. What are the problems associated with using a series of lead–acid cells?

Sulfuric acid is corrosive. Avoid

contact with skin and eyes. Wash

spills with a copious supply of

water.

EXPERIMENT 15.2


71


Investigating a hydrogen–oxygen fuel cell

AIMTo investigate the chemistry of the hydrogen–oxygen fuel cell

APPARATUStransparent plastic cup

2 graphite electrodes

500 mL 1 mol L–1 potassium hydroxide solution

2 semi-micro test tubes

gloves

retort stand

boss head

clamps

DC power supply

voltmeter

METHOD

1. Set up the fuel-cell structure as illustrated in the figure below, using a retort

stand, boss head and clamps to set the plastic cup in a vertical position.

Retortstand

Clamp

Cup

Semi-micro test tubes

Potassiumhydroxide solution

Siliconesealant

Carbon rods

2. Fill the cup with 1 mol L–1 potassium hydroxide solution until the carbon

rods are submerged.

3. Fill a semi-micro test tube with potassium hydroxide solution. Using gloves,

invert the test tube, then lower it over one of the carbon electrodes.

4. Repeat step 3 placing another semi-micro test tube over the second carbon

electrode.

5. Connect the bottom ends of the carbon electrodes to the terminals of a DC

power supply.

6. Adjust the output of the power supply to 6 volts and switch the power

supply on to begin electrolysis. You will notice that a gas begins to form at

each electrode.

EXPERIMENT 15.3




72

7. Switch off the power suppy when one of the test tubes is at least half-filled

with gas.

8. Disconnect the power supply, then connect the electrodes of the fuel cell to

a voltmeter and record the voltage of the cell.

QUESTIONS

1. Identify the cathode and anode in the fuel cell. Write equations for the

reactions occurring at each electrode and the overall cell equation.

2. Account for the difference in gas volumes evolved at the electrodes.

3. Write an equation for the reaction that occurs when the fuel cell is connected

to the voltmeter.

4. How does the voltage of the cell during discharge compare with the

theoretical voltage produced in a standard hydrogen–oxygen fuel cell?

Account for any differences.

5. In what ways is your laboratory fuel cell similar to a commercial hydrogen–

oxygen fuel cell?

6. In what ways is your laboratory fuel cell different from a commercial

hydrogen–oxygen fuel cell?

EXPERIMENT 15.3


73


Electrolysis of aqueous solutions of electrolytes

Electrolysis of molten potassium bromide yields bromine at the anode and potassium at the cathode. However, when a concentrated solution of potassium bromide in water is electrolysed, bromine is still produced at the anode, but hydrogen is evolved at the cathode.

In general, the electrolysis of an aqueous solution produces a metal or

hydrogen at the cathode and oxygen or a halide at the anode.

AIMTo investigate the electrolysis of solutions of salts, acids and alkalis

APPARATUS0.5 mol L–1 solutions of each of the following: nitric acid, potassium hydroxide, sodium bromide, sodium iodide and copper(II) sulfateuniversal indicatorelectrolytic cell with carbon electrodes2 ignition tubes6 V DC power source2 electrical leads with clipsstand, boss head and clampsplints

METHOD

1. Copy tables 16A, B, C (on the next page) and complete your predictions of the products of electrolysis of each solution (table 16C).

2. Set up the apparatus as shown in the figure below.

EXPERIMENT 16.1

Ignition tube

Bosshead

Clamp

Solution

undergoing

electrolysis

Cathode

(carbon

electrode)

Electrical lead

with clip

6 V DC power source

Stand

Anode

(carbon

electrode)

3. Fill the electrolytic cell and ignition tubes with the nitric acid solution.

4. Set the voltage supply to 6 V and pass electricity through the solution until the ignition tube at the cathode is full of gas.

5. Observe, and record, any colour changes around the anode or cathode.

6. Test any gas produced at the anode or cathode with a lighted splint.

7. Carefully smell the product at the anode and record your results.

8. Test the products at each electrode with universal indicator. Record the pH.

9. Wash out the electrolytic cell and ignition tubes.

10. Repeat steps 1 to 9 for the following solutions, in

order: potassium hydroxide, sodium bromide, sodium

iodide, copper(II) sulfate.




74

RESULTSCopy and complete tables 16A and 16B.

TABLE 16A

Solution

Observations at:

Anode Cathode

HNO3(aq)

KOH(aq)

NaBr(aq)

NaI(aq)

CuSO4(aq)

TABLE 16B

Solution

Name of product at:

Anode Cathode

HNO3(aq)

KOH(aq)

NaBr(aq)

NaI(aq)

CuSO4(aq)

TABLE 16C

Solution

Predictions at:

Anode Cathode

HNO3(aq)

KOH(aq)

NaBr(aq)

NaI(aq)

CuSO4(aq)

QUESTIONS

1. Write the overall cell reaction for each solution tested.

2. How does the use of universal indicator help identify reaction products?

3. How do the products of the electrolysis of the NaBr(aq) solution tested

compare with the expected products of electrolysis of the molten

compound?

EXPERIMENT 16.1


75


Factors affecting electrolysis

The products of electrolysis cannot always be predicted with confidence. In

an electrolytic cell, changing the conditions under which the electrolysis is

performed may produce different end products.

Electrolysis products may be affected by the concentration of the electrolyte,

the nature of the electrodes, the distance between electrodes, the area of the

electrodes, temperature, presence of other substances, preparation of the surface

to be plated and voltage.

AIMTo study the effect of changing one of the parameters of a specific electrolytic

cell

METHOD

1. The class should decide on a specific electrolytic cell to investigate, such

as a copper(II) sulfate solution.

2. Choose a parameter you wish to investigate. Each group should choose a

different parameter.

3. Design a controlled experiment to test your idea. Check your design with

your teacher before proceeding with the experiment.

4. Try to predict the results you may find.

5. Prepare a report of your experimental findings that can be presented to

the rest of the class. Include a statement of how your expected findings

compared with your experimental findings.

6. Summarise class results in a table.

QUESTIONS

1. Explain how three different factors may affect the products of electrolysis.

2. Swap experimental results with someone who investigated a different factor

from you. Write a full experimental report using their second-hand data.

EXPERIMENT 16.2



76


Electroplating

In electroplating, a thin layer of one metal is deposited on the surface of another

metal. The article to be plated is made the cathode of an electrolytic cell and

the electrolyte contains ions of the element forming the plate. The anode is

usually made from the same element as the metal plating.

AIMTo plate a piece of copper with nickel

APPARATUSnickel sheet, 5 cm × 3 cm

copper sheet, 5 cm × 3 cm

80 mL 0.2 mol L–1 nickel(II) sulfate solution

25 mL 1 mol L–1 sodium hydroxide solution

distilled water

nail varnish

propanone

steel wool

2 × 100 mL beakers

paper towels

6 V DC power source

2 electrical leads with clips

holder for electrodes

Nickel

sheet

Nickel(II) sulfate

solution

Beaker

Copper

sheetClips

Electrode

holder

METHOD

1. Clean both sides of the copper sheet using the steel wool.

2. Dip the copper sheet into a beaker containing sodium hydroxide solution.

Leave it for a few seconds, then wash the sheet with distilled water.

3. Write your name on the copper sheet using clear nail varnish and allow the

varnish to dry.

4. Place the nickel(II) sulfate solution into the second beaker.

5. Place the nickel sheet (anode) and the copper sheet (cathode) in the electrode

holder.

EXPERIMENT 17.1

Sodium hydroxide is caustic.

Spills on skin should be washed

immediately under running water

from a tap.

Propanone is highly fl ammable.

Keep it away from fl ames.



77

6. Lower the electrodes into the nickel(II) solution and connect the electrodes

into the circuit, as shown in the figure on the previous page.

7. Allow the current to flow for 10 minutes.

8. Remove the copper sheet. Remove the nail varnish by dipping a sheet of

paper towel in propanone and rubbing gently over the surface of the copper

sheet.

RESULTSDraw a fully-labelled sketch of your nickel-plated copper sheet.

QUESTIONS

1. Why isn’t it necessary to clean the nickel sheet for this experiment?

2. Why is it necessary to clean the copper plate with sodium hydroxide?

3. Why does nickel not form on the nail varnish?

4. What would happen in this experiment if the voltage used was:

(a) much higher, or

(b) much lower?

5. Describe the appearance of the solution in the beaker during the

electrolysis.

6. Write appropriate equations for the electrolytic reactions occurring during

this experiment.

EXTENSIONInvestigate the effect of: (a) temperature, (b) concentration, (c) pH values, and

(d) compositions of electrolytes, on electroplating.

EXPERIMENT 17.1



78


Anodising aluminium

The thickness of the layer of aluminium oxide is often deliberately increased to

give extra protection against corrosion through a process called anodising. This

is a way of increasing the thickness of the aluminium oxide layer by electrolysis.

A piece of aluminium is made into a positive electrode by connecting it to

the positive terminal of a power supply. It is dipped into an electrolysis cell

containing dilute sulfuric acid and an electric current is passed through it. During

electrolysis, oxygen is made at the aluminium electrode. The oxygen reacts with

the aluminium, increasing the layer of aluminium oxide. The anodised layer can

be dyed for a more attractive fi nish.

AIMTo anodise aluminium

APPARATUS2 sheets (10 cm × 20 cm and 5 cm × 8 cm) of aluminium (alloy 1145-0, available

from Alcoa of Australia or the Science Teachers’ Association of Victoria)

400 mL of 2 mol L–1 H2SO4

detergent

3 × 500 mL beakers

2 alligator clips and wires

forceps

paperclip

wooden icy-pole stick

heating apparatus

tripod

gauze mat

DC power source

tissue

conductivity testing apparatus

400 mL water-soluble dye solution

PART A: ANODISING ALUMINIUMMETHOD

1. Line a 500 mL beaker with a 10 cm × 20 cm piece of aluminium and fi ll

the beaker with 400 mL of 2 mol L–1 sulfuric acid. Use an alligator clip to

attach a wire to the cylinder (without allowing the alligator clip to dip into

the sulfuric acid), then connect the wire to the negative terminal of a DC

power source. The aluminium cylinder will become the cathode (see the

fi gure below).

Cylindrical

cathode

2. Thoroughly clean and degrease a 5 cm × 8 cm piece of aluminium. Scrub

with warm water and detergent, then rinse thoroughly and dry with a tissue.

Once cleaned, the aluminium sheet should be handled only with forceps.

EXPERIMENT 17.2

Safety goggles should be worn

during the electrolysis.



79

3. Suspend the aluminium sheet in the centre of the beaker so that it does not

touch the cylinder. A bent paperclip can be used to suspend the sheet from

a wooden icy-pole stick lying across the top of the beaker (see the figure

below).

4. Connect the aluminium sheet in the centre of the beaker to the positive

terminal of the DC power source. This central sheet of aluminium will

become the anode.

Cathode

Bent

paperclip

Paddlepop

stick

Beaker Anode

5. Turn on the power supply and increase the voltage slowly to 12 volts. Bubbles

should appear in the solution indicating that the electrolysis reaction is well

under way. Leave for 15 minutes, then remove the aluminium anode.

QUESTIONS

1. Describe the appearance of the anodised part of the aluminium anode

compared to the part that had not been in the solution.

2. Using conductivity apparatus, compare the conductivity of the anodised

part of the aluminium anode with that of the part which had not been in the

solution. Explain your results.

PART B: COLOURING THE ANODIC OXIDE LAYERMETHOD

1. Heat a prepared dye solution until it boils.

2. Completely immerse the anodised aluminium in the dye solution and leave

for ten minutes, or until the aluminium has acquired a permanent colour.

3. Rinse in water and allow to cool.

4. Seal the coloured oxide layer by totally immersing the aluminium in boiling

water for ten minutes.

5. Rinse in water and dry.

QUESTIONS

1. Using conductivity apparatus, compare the conductivity of your coloured,

anodised aluminium strip when:

(a) the aluminium is touched with the electrodes

(b) the aluminium is scratched with the electrodes.

Explain your results.

2. In what ways does an electrolytically-produced anodic coating differ from

aluminium’s normal oxide layer?

EXPERIMENT 17.2



80


Vitamins and minerals

Vitamins and minerals are obtained from food that we eat and are essential

for our bodies to function properly. They keep people healthy and assist in

growth and development. Vitamins are organic chemicals whereas minerals are

inorganic substances.

PART APrepare a short presentation on the importance of vitamins and minerals in the

diet. In particular state the role of iron, salt and vitamin C in the body. State

the sources of these substances and the effects of a deficiency of iron, salt or

vitamin C.

PART BUse a redox volumetric titration using potassium permanganate to find the

percentage iron in a liquid iron supplement and compare your findings with the

manufacturer’s stated value.

PART CUse an acid–base volumetric titration to analyse for vitamin C in tablets and

compare your findings with the manufacturer’s stated value.

PART DUse a gravimetric analysis to measure the percentage mass of salt in salt tablets

and compare your findings with the manufacturer’s stated value.

PART EUse data obtained in atomic absorption spectroscopy to determine the percentage

by mass of iron in a liquid iron supplement.

PART FQUESTIONS

1. (a) In volumetric analysis it is important to wash the burette and the pipette

with the solution that is being put into them. Explain why this is the

case.

(b) Should conical flasks and volumetric flasks also be washed with the

solution that is put into them? Explain your answer.

2. Why is an indicator NOT required for the reaction in part C?

3. If there was citric acid present in the vitamin C tablets, how would this

affect the result for part C?

4. Discuss why it is important that the sodium hydroxide solution be recently

standardised for part C.

5. A back titration is sometimes performed in analysis of a commercial

product. Explain what a back titration is and why it would be required.

6. What are the main sources of error associated with:

(a) volumetric analysis

(b) gravimetric analysis?

7. For each of the analyses above:

(a) describe an alternative method that might be used

(b) outline the disadvantages and advantages of each method.

ALTERNATIVE INVESTIGATIONSTitrate the vitamin C with standardised iodine solution and compare the results.

Measure the amount of vitamin C in a variety of tablets or fruit juices.



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An investigation of iron in a liquid iron supplement

AIM To find the amount of iron(II) sulfate in a liquid iron supplement

APPARATUSburette

20.00 mL pipette

25.00 mL pipette

retort stand

measuring cylinder

beaker

4 × 100 mL conical flasks


liquid iron supplement

1 mol L–1 sulfuric acid

standardised 0.005 00 mol L–1 potassium permanganate

METHOD

1. Record the brand and iron content of the liquid iron supplement.

2. Record the concentration of the potassium permanganate solution.

3. Pipette a 25.00 mL aliquot of the iron supplement into a volumetric flask.

4. Add 40 mL of 1.0 mol L–1 sulfuric acid.

5. Make up to 250 mL using deionised water.

6. Transfer a 20.00 mL aliquot into a conical flask.

7. Titrate with 0.005 00 mol L–1 potassium(IV) permanganate. The end point

of the titration is reached when the purple colour of the permanganate

solution remains.

8. Repeat steps 6 and 7 until three concordant results are obtained.

QUESTIONS

1. Write the half equations for the reactions showing the oxidation of Fe2+

to Fe3+ and the reduction of MnO4– to Mn2+. From these, write the overall

equation.

2. Why is acid added to the volumetric flask?

3. Calculate the average volume of the three concordant titres.

4. Calculate the number of moles of MnO4– used to react with the Fe2+

aliquot.

5. Determine the number of moles of Fe2+ that reacted.

6. Calculate the number of moles of Fe2+ present in the original solution.

7. Calculate the mass of Fe present in the sample.

8. Calculate the mass of Fe present per mL.

9. Calculate the mass of FeSO4.7H2O in the sample.

10. Compare your result with that of the manufacturer and comment on your

findings.




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Analysis of vitamin C

Vitamin C is important for normal health and development. It is a water soluble

vitamin and is used as an antioxidant to preserve food. Humans cannot make

vitamin C so it must be included as part of their daily diet. It is found in most

fruits and vegetables, particularly citrus fruits, green peppers, strawberries and

tomatoes.

AIMTo determine the percentage of ascorbic acid (C6H8O6) in a tablet of vitamin C


measuring cylinder

electronic balance

3 × conical flasks

funnel

beaker

white tile

standardised 0.1 mol L–1 NaOH

vitamin C tablets

phenolphthalein

METHOD

1. Record the brand and vitamin C content of the tablet.

2. Record the accurate concentration of the NaOH solution.

3. Accurately weigh a tablet of vitamin C. Record the mass.

4. Place the tablet in a conical flask with about 50 mL of warm water.

5. Use a stirring rod to crush the tablet.

6. Add three drops of phenolphthalein.

7. Prepare the burette by rinsing and then filling with the NaOH solution.

8. Titrate the vitamin C solution with 0.1 mol L–1 NaOH solution until the faint pink colour persists for 30 seconds.

9. Repeat steps 3–8 until three concordant titres are obtained.

QUESTIONS

1. Write an ionic equation for the reaction of ascorbic acid with sodium hydroxide to produce sodium hydrogen ascorbate (NaC6H7O6).

2. Calculate the number of moles of sodium hydroxide that reacted with the ascorbic acid.

3. Calculate the number of moles of ascorbic acid present in one tablet.

4. Calculate the mass (in milligrams) of ascorbic acid present in one tablet.

5. How does this compare with the manufacturer’s value?

6. Determine the percentage by mass of ascorbic acid (vitamin C) in one tablet.

7. How would the calculated percentage of ascorbic acid in the vitamin C tablet be affected if the tip of the burette were not filled with sodium hydroxide solution before you started titrating?

8. How would the calculated percentage of ascorbic acid be affected if the burette were rinsed with water instead of NaOH?

9. Why is it important that the sodium hydroxide solution be standardised

before use in this titration?



83


Gravimetric determination of

salt in salt tablets

Sodium chloride, commonly referred to as salt, is an essential mineral that

helps carry nutrients to the cells, and regulates body functions such as blood

pressure and fluid volume. Sodium is important in nerve conduction. Chlorine is

also needed for good health; it maintains the acid–base balance, aids potassium

absorption and is present in stomach acid. In this experiment the chloride

ions from salt are precipitated as silver chloride. By finding the mass of the

precipitate, the amount of salt in a salt tablet can be determined.

AIMTo find the percentage of salt in salt tablets

APPARATUS1 salt tablet



25 mL 0.1 mol L–1 silver nitrate

20.00 mL pipette

beaker

bench mat

deionised water

electronic balance

filter funnel

filter paper

Gooch crucible

mortar and pestle

stirring rod

vacuum flask and vacuum pump

oven

METHOD

1. Record the stated mass and brand of the salt tablet.

2. Weigh one tablet.

3. Crush the tablet and dissolve it in 50 mL of deionised water in a 100 mL

beaker.

4. Pour into a 100 mL volumetric flask and make up to 100 mL with deionised

water. Mix thoroughly.

5. Take a 20.00 mL aliquot and place it in a 100 mL beaker

6. Add 25 mL of 0.100 mol L–1 AgNO3 drop by drop from a measuring

cylinder while stirring continually.

7. Weigh a Gooch crucible and filter paper.

8. Pour the sample into the crucible lined with filter paper, and wash remaining

solid into crucible with water.

9. Filter using vacuum filtration.

10. Put in oven to dry overnight.

11. Weigh dried precipitate.




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QUESTIONS

1. Write the overall and ionic equations for this reaction.

2. What is the mass of the precipitate?

3. How many moles of silver chloride are present in the precipitate?

4. How many moles of NaCl are present in the aliquot?

5. What is the mass of NaCl in one tablet?

6. Calculate the percentage mass of salt in one tablet.

7. Give two sources of error that may have affected your results.

8. Suggest a modification that would improve the design of this experiment.

Give the reasons for your suggestion.



85


Investigating iron content in iron tablets using atomic absorption spectroscopy

The mineral iron is important for delivering oxygen around the body. Iron

is a component of haemoglobin, which is found in red blood cells. Oxygen

combines with haemoglobin to form oxyhaemoglobin, which is transported to

body cells.

Anaemia is caused by an iron deficiency and sometimes an iron supplement is

advised if a change in diet does not help. The following analysis of iron tablets

required the preparation of a number of standard Fe2+ solutions. The absorbance

of these solutions was measured using an atomic absorption spectrometer. The

absorbance can be plotted against the values of the Fe2+ concentrations.

METHOD

1. A stock solution containing 50 mg L–1 Fe2+ was provided and used to

prepare the solutions listed in the table below.

2. An atomic absorption spectrometer was used to determine the absorbances

of these solutions. The following absorbances were obtained:

Concentration

Fe2+ (mg L–1) Absorbance

0.0 0.0

0.50 2.3

1.00 4.5

1.50 6.9

2.00 9.2

3. Two iron tablets were then weighed, dissolved in hot HCl, filtered, cooled

and made up to 100 mL in a volumetric flask. This solution was diluted

twice as shown below using deionised water.

5 mL 10 mL

Iron tablets

100 mL

Dissolving tablets

A

100 mL

First dilution

B

100 mL

Second dilution

C




86


4. A sample from the second dilution was analysed in an atomic absorption

spectrometer and the absorbance was found to be 2.0.

QUESTIONS

1. You have four 100 mL standard flasks and a graduated 10 mL pipette.

Explain how you would prepare the following solutions: 0.5 mg L–1,

1.0 mg L–1, 1.5 mg L–1, 2.0 mg L–1 using the stock solution.

2. Plot a graph of iron concentration in ppm (mg L–1) versus absorbance using

the data given above.

3. Use the absorbance value and your graph to find the concentration of iron

in flask C.

4. Determine the concentration of Fe2+ in the original flask (flask A).

5. Calculate the mass of iron in millgrams in the original flask.

6. Calculate the mass of iron in one iron tablet.

7. If the mass of one tablet was 0.600 g (600 mg), calculate the percentage

iron by mass in the tablet.

8. Determine the mass of FeSO4.7H2O present in 1 litre of the initial stock

solution containing 50 mg L–1 Fe2+.


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