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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
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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
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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
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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
potassium permanganate
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.
Take care when handling the
potassium permanganate
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.
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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
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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?
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Chemistry 2: Experiments EXPERIMENT 3.1
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.
Take care when handling the
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.
Take care when handling the
hydrochloric acid solution.
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10
Chemistry 2: Experiments
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
3 × 300 mL conical fl asks
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
hydrochloric acid solution.
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
hydrochloric acid solution.
EXPERIMENT 3.2
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Chemistry 2: Experiments
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
APPARATUSburette with stand
20.00 mL pipette and pipette fi ller
4 × 300 mL conical fl asks
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
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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
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Chemistry 2: Experiments
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
Chemistry 2: ExperimentsCONTINUED
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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
Chemistry 2: ExperimentsCONTINUED
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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.
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16
Chemistry 2: Experiments
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
Chemistry 2: ExperimentsCONTINUED
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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.
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Chemistry 2: Experiments
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
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19
Chemistry 2: Experiments EXPERIMENT 4.3
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?
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Chemistry 2: Experiments
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.
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21
Chemistry 2: Experiments
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
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Chemistry 2: Experiments
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
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23
Chemistry 2: Experiments
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.
Chemistry 2: ExperimentsCONTINUED
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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
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Chemistry 2: Experiments EXPERIMENT 6.1
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?
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Chemistry 2: Experiments
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
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27
Chemistry 2: Experiments
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
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Chemistry 2: Experiments EXPERIMENT 7.1
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.
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29
Chemistry 2: Experiments EXPERIMENT 7.2
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?
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Chemistry 2: Experiments
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
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31
Chemistry 2: Experiments
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
APPARATUS3 test tubes
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
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Chemistry 2: Experiments EXPERIMENT 7.5
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.
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33
Chemistry 2: Experiments
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
Chemistry 2: ExperimentsCONTINUED
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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
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35
Chemistry 2: Experiments
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
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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
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37
Chemistry 2: Experiments
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
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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
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39
Chemistry 2: Experiments EXPERIMENT 9.2
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.
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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
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41
Chemistry 2: Experiments
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
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42
Chemistry 2: Experiments
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.
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43
Chemistry 2: Experiments
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’.
TASK 2 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
[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
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TASK 3 Concentration data (mol L–1) at 30°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 0.85 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
TASK 4 Concentration data (mol L–1) at 35°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 0.80 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.
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45
Chemistry 2: Experiments
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
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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
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47
Chemistry 2: Experiments
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
APPARATUS7 test tubes
test-tube rack
universal indicator solution
10 mL measuring cylinder
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
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Chemistry 2: Experiments EXPERIMENT 11.1
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.
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49
Chemistry 2: Experiments
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?
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EXPERIMENT 11.2
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51
Chemistry 2: Experiments EXPERIMENT 11.3
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.
Chemistry 2: ExperimentsCONTINUED
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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
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53
Chemistry 2: Experiments
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.
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Chemistry 2: Experiments
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
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55
Chemistry 2: Experiments
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.
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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
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57
Chemistry 2: Experiments
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
1. Place 10 mL of water into a test tube, then measure and record the initial
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.
Chemistry 2: ExperimentsCONTINUED
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PART C: DISSOLVING AMMONIUM NITRATE IN WATERMETHOD
1. Place 10 mL of water into a test tube, then measure and record the initial
temperature.
2. Carefully add a spatula of solid ammonium nitrate 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 NH4NO3(s) NH4NO3(aq) endothermic or exothermic?
Justify your decision.
2. Draw an energy diagram for this reaction.
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.
3. Draw an energy diagram for this reaction.
PART E: DISSOLVING ANHYDROUS COPPER(II) SULFATE IN WATERMETHOD
1. Place 10 mL of water into a test tube, then measure and record the initial
temperature.
2. Add a spatula of solid anhydrous copper(II) sulfate 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 endothermic or exothermic? Justify your decision.
2. Predict the temperature change for the reverse reaction.
3. Draw an energy diagram for this 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
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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
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Chemistry 2: Experiments
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
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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
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Chemistry 2: Experiments
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
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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
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64
Chemistry 2: Experiments
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
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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
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Chemistry 2: Experiments
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
APPARATUS4 test tubes
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
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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
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Chemistry 2: Experiments
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
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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
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Chemistry 2: Experiments
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
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71
Chemistry 2: Experiments
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
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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
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Chemistry 2: Experiments
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.
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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
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Chemistry 2: Experiments
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
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Chemistry 2: Experiments
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.
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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
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Chemistry 2: Experiments
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.
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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
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Chemistry 2: Experiments
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.
EXTENDED INVESTIGATIONS
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Chemistry 2: Experiments
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
250 mL volumetric flask
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.
EXTENDED INVESTIGATIONS
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Chemistry 2: Experiments
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
APPARATUSburette with stand
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?
EXTENDED INVESTIGATIONS
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83
Chemistry 2: Experiments
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
100 mL measuring cylinder
100 mL volumetric flask
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.
EXTENDED INVESTIGATIONS
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Chemistry 2: Experiments
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.
EXTENDED INVESTIGATIONS
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Chemistry 2: Experiments
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
EXTENDED INVESTIGATIONS
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Chemistry 2: Experiments
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+.
EXTENDED INVESTIGATIONS
top related