EXPERIMENT
St. Clare's Girls' School
F.7 Chemistry
DATE :
A Study of the Oxidation States of Vanadium
Procedure
First complete a copy of Table A to assist you in identifying the different oxidation states of vanadium.
Results Table A
Ion (hydrated)
Colour
*VO2
+
Yellow
VO2+
V3+
V2+
Blue
Green
Violet
Oxidation state
Name
+
* The VO3- ion in the ammonium salt is converted to VO2 by acid:
-
VO3 (aq) + 2H+ (aq) ↔ VO2+ (aq) + H2O (l)
1.
Place about 0.25 g (one spatula measure) of ammonium trioxovanadate(V) in a conical flask and add about 25cm³ of dilute
sulphuric acid. Carefully add about 5 cm³ of concentrated sulphuric acid and swirl the flask until you a clear yellow
solution.
2.
Pour about 2 cm³ of this vanadium(V) solution into each of two test tube ready for later tests.
3.
To the conical flask add 1-2 g (one spatula measure) of zinc dust, a little at a time. Swirl the flask at intervals and record
any observed colour changes in a copy of Results Tables B.
4.
When the solution has become violet (you may need to heat the flask for this final change), filter about 2 cm³ into each of
three test tubes.
5.
To one of the three tubes add, a little at a time, an excess of acidified potassium manganate(VII) solution, shaking after
each addition, until no further change is observed.
6.
Keep the other test tubes of solutions for later tests. Answer the first two questions (next page) and complete Results Table
B as far as you can before doing further tests.
7.
To one of the tubes containing vanadium(V), add a little sodium sulphite and shake. Filter if cloudy. Now boil carefully
(at a fume cupboard) to remove excess sulphur dioxide and add about the same volume of the vanadium(II) solution.
Record your observations.
8.
To the second of the tubes containing vanadium(V), add about 2 cm³ potassium iodide solution and mix. Then add about 2
cm³ of sodium thiosulphate solution. Record your observations.
9.
Keep the last tube of vanadium(II) solution for a final experiment you may wish to do after answering question 7.
Results Table B
Test
Observation
Summary of reaction
Ammonium vanadate + acid
White solid dissolved to a yellow
solution.
+5
VO3
+5
VO2+
Vanadium(V) + zinc
Vanadium(II) + manganate(VII)
Vanadium(V) + sulphite
Add vanadium(II)
Vanadium(V) + iodide
+ thiosulphate
Vanadium(II) + ……
Questions
1.
How do you explain the first appearance of a green colour in the solution?
2.
What are the subsequent changes in colour and why do these changes occur?
You will need the following electrode potentials in order to answer some of the remaining questions:
Zn2+ (aq) + 2e- ↔ Zn(s)
3+
-
2+
V (aq) + e ↔ V (aq)
VO2+ (aq) + 2H+ (aq) + e- ↔ V3+ (aq) + H2O (l)
+
+
-
2+
E = - 0.76V
E = - 0.26V
E = + 0.34V
VO2 (aq) + 2H (aq) + e ↔ VO (aq) + H2O (l)
E = + 1.00V
SO42- (aq) + 4H+ (aq) + 2e- ↔ H2SO3 (aq) + H2O (l)
E = + 0.17V
I2 (aq) + 2e- ↔ 2 I (aq)
E = + 0.54V
3.
What did you observe when you added iodide ions to vanadium(V)? What caused this colour?
4.
Why did you add sodium thiosulphate?
5.
Why does reduction with iodide not give the same result as reduction with zinc?
6.
What did you observe when you added sulphite ions to acidified vanadium(V) solution? Does this result correspond with a
prediction made using the Eθ values? (Hint: sulphite ions and acid react to give …)
7.
How would you set about finding a suitable oxidizing agent for the oxidation of vanadium(II) to vanadium(III) and no
further? Does one appear in the table above?
EXPERIMENT
St. Clare's Girls' School
F.7 Chemistry
DATE :
A Colorimetric Method to Determine the
Formula of Complex Ions
Principle of Method
Many d-block metal ions form coloured complexes with various ligands. Their concentration can be estimated by monitoring
the absorbance using a colorimeter equipped with a specifically chosen filter. By mixing different proportions of metal ions
and ligands together and measuring their absorbance, the relative ratio of the mixture with maximum absorbance corresponds
to the optimum stoichiometry of the complex ion. Such approach to carry out investigations based on preparing a series of
mixtures with varying composition is known as continuous variation.
Procedure
1.
Prepare a different mixture in each of the 8 test-tubes by running in from burettes or graduated pipettes the volumes of the
three solutions shown in the Results table below. Cork each tube and shake thoroughly.
Tube number
1
2
3
4
5
6
7
8
Volume of ammonium sulphate / cm3
1.5
0.5
0.5
0.5
0.5
0.5
0.5
0.5
Volume of copper(II) sulphate / cm3
0.0
0.1
0.15
0.2
0.25
0.3
0.4
0.5
Volume of ammonia / cm3
0.0
0.9
0.85
0.8
0.75
0.7
0.6
0.5
Relative ratio of Cu2+ / NH3
Absorbance
2.
Calibrate the colorimeter by setting it to 0% absorbance (or 100% transmittance) when the light is shining through a
solution of ammonium sulphate filled in a cuvette.
3.
Measure the corresponding absorbance for each sample in the series of solutions that you prepared earlier.
a. Rinse the cuvette with a small amount of your first sample, then fill the cuvette ¾ of the way with the sample. Wipe
with the lens paper and place cuvette in the measuring chamber so the mark on the cuvette lines up with the mark on
the measuring chamber.
b. Select a filter of a certain colour.
c. Allow the reading to stabilize and record it.
d. Remove and rinse the cuvette and repeat steps a through b for all of the remaining samples.
e. Repeat the entire recording procedure using alternative filter(s).
Data Treatment
By plotting a suitable labeled graph, deduce the formula of the complex formed between copper(II) ion and ammonia.
Discussion Questions
1.
What is meant by transmittance and absorbance ? Explain which term is more preferable in doing experiment.
2.
What is the purpose of introducing ammonium sulphate in the above mixtures ?
3.
Sketch the schematic diagram of a colorimeter, labeling all essential components.
4.
Comment on the choice of filter employed in this experiment.
5.
In preparing the above series of mixtures, the composition where ammonia possessing the greater proportion (>50%) are
omitted. Suggest a possible reason for this.
6.
Ni2+(aq) reacts with H2NCH2CH2NH2(aq) to form a purple complex cation with the formula [Ni(en)2]2+(aq), where en
represents H2NCH2CH2NH2. Outline the experimental procedure and the data treatment to verify the formula of the
complex cation.
EXPERIMENT
St. Clare's Girls' School
F.7 Chemistry
DATE :
Hardness of water by EDTA Titration
Purposes
1.
To determine the concentrations of Ca2+ and Mg2+ in a commercial sample of bottled mineral water and
2.
To compare experimental results with the concentrations of the metal ions claimed by the manufacturer.
Introduction
The ions involved in water hardness, i.e. Ca2+(aq) and Mg2+(aq), can be determined by titration with a chelating agent,
ethylenediaminetetraacetic acid (EDTA), usually in the form of disodium salt. The titration reaction is:
Ca2+(aq) + H2Y2-(aq) → CaY2-(aq) + 2H+(aq)
(where H2Y2- represents the disodium salt of EDTA, see below)
Eriochrome Black T is commonly used as indicator for the above titration. At pH 10, Ca2+(aq) ion first complexes with the
indicator as CaIn+(aq) which is wine red. As the stronger ligand EDTA is added, the CaIn+(aq) complex is replaced by the
CaY2-(aq) complex which is blue. The end point of titration is indicated by a sharp colour change from wine red to blue.
Titration using Eriochrome Black T as indicator determines total hardness due to Ca2+(aq) and Mg2+(aq) ions. Hardness due to
Ca2+(aq) ion is determined by a separate titration at a higher pH, by adding NaOH solution to precipitate Mg(OH)2(s), using
hydroxynaphthol blue as indicator.
O 2N
HOOC
_
N
CH2
CH2
N
OOC
COOH
_
COO
N
_ OH
SO 3
EDTA (anionic form)
N
HO
Eriochrome Black T
Safety
Avoid skin contact with chemicals. Clothing contaminated with NaOH solution should
be carefully removed. Spillage adhering to skin should be immediately washed with
plenty of water.
EYE PROTECTION
MUST BE WORN
Materials and Apparatus
50% w/v NaOH solution (50 g in 100 cm3 solution)
CORROSIVE
Eriochrome Black T indicator
IRRITANT
pH 10 buffer, hydroxynaphthol blue indicator, 0.01 M
EDTA (disodium salt)
Commercial sample of bottled mineral water, apparatus required for titration, 5 cm3 measuring cylinder.
HARMFUL
Experimental Procedures
Part A: Determination of total hardness
1.
2.
3.
4.
Pipette 50 cm3 (25 cm3 x 2) of mineral water into a conical flask.
Add 2 cm3 NH3/NH4Cl pH 10 buffer followed by 3 drops of Eriochrome Black T indicator solution.
Titrate with 0.01 M EDTA until the solution turns from wine red to sky blue with no hint of red (save
the solution for colour comparison).
Repeat the experiment to obtain two concordant results.
Part B: Determination of Ca2+
1.
2.
3.
4.
Pipette 50 cm3 of mineral water into a conical flask.
Add 30 drops of 50% w/v NaOH solution, swirl the solution and wait for a couple of minutes to
completely precipitate the magnesium ions as Mg(OH)2(s).
Add a pinch of hydroxynaphthol blue (exact amount to be decided by the intensity of the resulting
coloured solution) and titrate with 0.01 M EDTA until it changes to sky blue (save the solution for
colour comparison).
Repeat the experiment to obtain two concordant results.
Results
Part A: Determination of total hardness
Final burette reading/cm3
Trial
1
2
1
2
Initial burette reading/cm3
Volume used/cm3
Average volume of 0.01 M EDTA used/cm3
Part B: Determination of Ca2+
Final burette reading/cm3
Trial
Initial burette reading/cm3
Volume used/cm3
Average volume of 0.01 M EDTA used/cm3
Calculation
1.
2.
3.
From Part A of the experiment, determine the total concentration of Ca2+ and Mg2+ ions in the mineral
water sample in mol dm-3.
From Part B of the experiment, determine the concentration of Ca2+(aq) in the mineral water sample in
mg dm-3, or ppm.
Hence, calculate the concentration of Mg2+(aq) in the mineral water sample in mg dm-3 or ppm. Compare
with the corresponding values displayed on the label of the bottle.
Discussion Questions
1.
Why are two indicators used in the experiment? Can the first indicator be used for the second titration?
2.
What are the limitations of the EDTA Titration in determining metal ion concentrations?
3.
Outline an alternative method for determining the hardness of water sample due to Ca2+(aq) ions.
EXPERIMENT
St. Clare's Girls' School
F.7 Chemistry
DATE :
To Study the Reactions of
Cyclohexane, Cyclohexene and Methylbenzene
Procedure
Carry out the following procedures using each of the hydrocarbons : cyclohexane (C6H12), cyclohexene (C6H10) and
methylbenzene (C7H8) and report your observations together with the corresponding inferences (ie. explanations and
appropriate equations) in a table form.
(A) Combustion
Put a few drops of the substance in an evaporating dish and ignite them by a lighting match. Note the degree of sootiness of the
flame as well as the residue left behind.
(Precaution : Perform the combustion in the fume cupboard)
(B) Reaction with Potassium Manganate(VII)
Put a few drops of potassium manganate(VII) into 2 cm3 of sodium carbonate solution in a test-tube, and add it to a few drops
of the substance in another tube.
(C) Reaction with Aqueous Bromine
Add bromine water slowly to a few drops of the substance in a test-tube and note any colour change. Leave the tube in direct
sunlight if there is no apparent observable change.
(D) Reaction with Concentrated Sulphuric Acid
Cool 5 cm3 of the substance under the cold water tap and then add an equal volume of cold, conc. sulphuric acid. Cool and
shake the mixture for several minutes. Then pour the mixture into an equal volume of cold water.
(Precaution : Wear safety goggles for the reaction with cyclohexene)
Questions
1.
Which of the above hydrocarbons is the most reactive ? Explain your choice with the aid of an appropriate mechanism
for the reaction (C).
2.
Account for the unreactive nature of the remaining two hydrocarbons other than the one you have chosen in Q.1.
EXPERIMENT
St. Clare's Girls' School
F.7 Chemistry
DATE :
To Study the Chemistry of the Halogens
For each of the following reactions, record your observation in a table form together with the corresponding inference, which
may include a brief explanation for the observed result and an appropriate equation for the reaction, if any.
(A) Displacement reactions of the halogens
1.
Discover the colour of solutions of chlorine, bromine and iodine in CCl4.
2.
Add chlorine water to each of the following aqueous solutions : (i) potassium iodide, (ii) potassium bromide, with CCl4
present in each case.
3.
Add bromine water to each of the following aqueous solutions : (i) potassium chloride, (ii) potassium iodide, with CCl4
present in each case.
(B) Action of acids on the halides
1.
Add 1cm3 phosphoric acid to a small sample of each of the three SOLID potassium halides (ie. KCl, KBr & KI) and
warm. Record your observations.
2.
Repeat step (1) using concentrated sulphuric acid instead of phosphoric acid WITHOUT WARMING. Record your
observations.
Caution : You should wear goggles and perform the test in a fume cupboard
(C) Action of silver(I) ions on the halides
1.
Place about 1 cm3 of silver(I) nitrate solution in each of the aqueous solution of the three potassium halides (ie. KCl, KBr
& KI). Record your observation.
2.
Expose the resulting solutions in step (1) to sunlight for about 3 minutes. Note any change.
Worksheet for Qualitative Analysis
of Inorganic Compounds
Name : ________________________________
Class : ______
Date : ____/____/____
1. Describe the solid.
2. Report on the solubility in water.
3. Estimate the pH.
4. Heat a small amount of the solid in an
ignition tube at first gently and then
strongly. Report the observation.
5. Report the observation on a flame test.
6. Report on the action of 2M HCl on the
solid.
7. Report on the action of 1M H2SO4 on
the solid.
8. Report on the action of conc. H2SO4
on the solid.
9. Report on the action of 2M NaOH on
the aqueous solution.
10. Report on the action of 2M NH3 on the
aqueous solution.
11. Report on the action of acidified
KMnO4 solution made from equal
volumes of 0.2M KMnO4 and
1M H2SO4 on the aqueous solution.
12. Report on the action of
Write your interpretation of the various observations reported above and comment on the chemical nature of the substance.
_______________________________________________________________________________________________________________
_______________________________________________________________________________________________________________
_______________________________________________________________________________________________________________
_______________________________________________________________________________________________________________
_______________________________________________________________________________________________________________
_______________________________________________________________________________________________________________
____________________________________________________________________________________________________
EXPERIMENT
St. Clare's Girls' School
F.7 Chemistry
DATE :
Preparation of Benzoic Acid from Ethyl Benzoate
Introduction
Most esters can be effectively hydrolyzed by refluxing with a slight excess of aqueous sodium hydroxide. The alcohol is
isolated directly, and the acid is removed from its sodium salt. If the alcohol is soluble in water, it is distilled out; if not, it is
extracted with ether.
In this experiment, the ester used for hydrolysis is ethyl benzoate.
O
O
O
CH2CH3
+
NaOH
O Na
+
+ CH3CH2OH
The ethanol produced is distilled out and the benzoic acid is precipitated by a strong acid, HCl. The acid solid is then further
purified by recrystallization.
Procedure
A. Hydrolysis of ester
1. Set up the apparatus for reflux.
2. Place about 5 cm3 of ethyl benzoate in the flask and add 30 cm3 of 2M NaOH.
3. Heat the mixture to boiling and reflux for about fifteen minutes. At the end of hydrolysis, the oily drops of ester
should have disappeared.
B. Separation and purification of hydrolysis products
1. When hydrolysis is complete, set up the apparatus for distillation.
2. Distil the mixture slowly. Collect the distillate at a temperature range of 76 - 80°C. Stop the distillation when the
temperature rises above 80°C.
3. Test the ethanol distilled by warming a few drops with a small amount of acidified potassium dichromate solution.
Note any observation.
4. Add 20 cm3 of concentrated hydrochloric acid to the residue in the distillation flask. Stir continuously. Allow to cool
to room temperature.
5. Filter off the benzoic acid solid by suction filtration using the Buchner funnel.
6. Purify the crude solid sample of benzoic acid by recrystallization.
7. Dry the crystals by leaving overnight in a desiccator and weigh the dried sample.
8. Determine the melting point of benzoic acid.
Questions
1.
Calculate the amount of ester used and hence determine the percentage yield of the reaction.
(Relative atomic masses : H, 1.0 ; C, 12.0 ; O, 16.0 ; density of ethyl benzoate =
g cm-3 )
2.
With the aid of a labelled diagram, show the laboratory apparatus required to carry out each of the following operations :
(a) reflux
(b) distillation
3.
Describe the experimental procedure by which you would recrystallize a sample of impure benzoic acid.
4.
How would you test the purity of the product obtained ? Comment on the purity of your sample.
EXPERIMENT
St. Clare's Girls' School
F.7 Chemistry
Investigating Some of the Properties of
A Pair of Geometrical Isomers
DATE :
EXPERIMENT
St. Clare's Girls' School
F.7 Chemistry
DATE :
The Analysis Of Aspirin Tablets
Introduction
In all pharmaceutical preparations, the manufacturer is required by law to state on the packaging the maximum amount of each
active ingredient present. In many preparations the active ingredient will only form a small percentage of the pill or tablet as a
whole. For example, in the case of many tablets, the tablet would simply disintegrate into a powder unless additives were
introduced to help the ingredients cohere into a tablet form.
In this experiment you will have the opportunity to carry out a consumer survey on the aspirin content of a number of a
commercial preparations, to see whether the manufacturer’s claims are justified.
Acetylsalicylic acid (aspirin), being dibasic, can be readily hydrolyzed by sodium hydroxide into the sodium salts of two weak
acids, ethanoic acid and salicylic acid. In this experiment, the hydrolysis is achieved using an excess of sodium hydroxide, the
excess being later found by titration with standard sulphuric acid. The equation for he hydrolysis reaction is:
COOH
O
C
CH3
COONa
OH
+ 2 NaOH
+ CH3COONa + H2O
O
Choice Of Indicators
Acid-base indicators are substances that change colour with change of pH. The observable colour change of an indicator
actually takes place over a range of pH  this is known as the indicator range. The table below shows the colours of some
common indicators in acid and base media and their pH ranges.
Colour
Indicator
Acidic
Alkaline
pH range
red
yellow
3.2 – 4.4
methyl red
yellow
red
4.8 – 6.0
bromothymol blue
yellow
blue
6.0 – 7.6
phenol red
yellow
red
6.8 – 8.4
colourless
pink
8.2 – 10.0
methyl orange
phenolphthalein
Choice of the appropriate indicators to detect the end-points in acid-base titrations depends on
(a) the pH of the solution at the equivalence point of titration (as seen from the corresponding pH titration curve)
(b) the pH range of the indicators
Generally speaking, if a visual indicator is to be used to determine the end-point of an acid-base titration, its pH range should
fall on the “vertical portion” of the titration curve. This gives a sharp colour shape at the end-point.
Procedure
Weigh accurately a definite number of aspirin tablets (two or three tablets weighing no more than 1.5 g) into a 250 cm3 conical
flask. Initiate the hydrolysis of the aspirin by adding 25 cm3 of 1.0 M sodium hydroxide by pipette, diluting with
approximately the same volume of distilled water. Warm the flask over a tripod and gauze for ten minutes to complete the
hydrolysis.
Cool the reaction mixture and transfer with washings to a 250 cm3 volumetric flask. Dilute to the mark with distilled water and
then ensure that the contents of the flask are well mixed by repeated shakings.
Titrate 25 cm3 portions of the diluted reaction mixture with the standard 0.05 M sulphuric acid provided, using phenol red
indicator. This indicator (with pH range 6.8-8.4) is most suitable for this titration because of the presence of the salts of the
two acids, though phenolphthalein is also satisfactory for this purpose.
Calculate the weight of acetylsalicylic acid in each tablet and compare your results with the manufacturer’s specification.
Compare your results with those given in different brands of aspirin.