EXPERIMENT St. Clare's Girls' School DATE : F.7 Chemistry

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) *VO2

+ VO2+ V3+ V2+

Colour Yellow 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 +5 VO3

� � 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) E = - 0.76V

V3+ (aq) + e- ↔ V2+ (aq) E = - 0.26V

VO2+ (aq) + 2H+ (aq) + e- ↔ V3+ (aq) + H2O (l) E = + 0.34V

VO2+ (aq) + 2H+ (aq) + e- ↔ VO2+

(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 DATE : F.7 Chemistry

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 DATE : F.7 Chemistry

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.

HOOC COOH

OOCN CH2 CH2 N

COO_ _

O2N

SO3

_ OH

N N

HO

EDTA (anionic form) 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.

Materials and Apparatus

50% w/v NaOH solution (50 g in 100 cm3 solution)

Eriochrome Black T indicator 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.

EYE PROTECTIONMUST BE WORN

CORROSIVE

IRRITANT HARMFUL

Experimental Procedures

Part A: Determination of total hardness

1. Pipette 50 cm3 (25 cm3 x 2) of mineral water into a conical flask. 2. Add 2 cm3 NH3/NH4Cl pH 10 buffer followed by 3 drops of Eriochrome Black T indicator solution. 3. 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). 4. Repeat the experiment to obtain two concordant results.

Part B: Determination of Ca2+

1. Pipette 50 cm3 of mineral water into a conical flask. 2. 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). 3. 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).

4. Repeat the experiment to obtain two concordant results. Results

Part A: Determination of total hardness

Trial 1 2 Final burette reading/cm3

Initial burette reading/cm3

Volume used/cm3

Average volume of 0.01 M EDTA used/cm3

Part B: Determination of Ca2+

Trial 1 2 Final burette reading/cm3

Initial burette reading/cm3

Volume used/cm3

Average volume of 0.01 M EDTA used/cm3

Calculation

1. From Part A of the experiment, determine the total concentration of Ca2+ and Mg2+ ions in the mineral water sample in mol dm-3.

2. From Part B of the experiment, determine the concentration of Ca2+(aq) in the mineral water sample in mg dm-3, or ppm.

3. 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 DATE : F.7 Chemistry

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 DATE : F.7 Chemistry

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 DATE : F.7 Chemistry

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

CH2CH3

NaOHO

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 DATE :

F.7 Chemistry

Investigating Some of the Properties of A Pair of Geometrical Isomers

EXPERIMENT St. Clare's Girls' School DATE : F.7 Chemistry

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:

OC

CH3

O

COOH

NaOH CH3COONaOHCOONa

+ + + H2O2

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

methyl orange 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

phenolphthalein colourless pink 8.2 – 10.0

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.