Experiment :

1

Title :

Preparation of standard solution and standardization of hydrochloric acid

Objective :

To prepare a standard solution of sodium carbonate and use it to standardize a given solution of dilute hydrochloric acid.

Introduction : Anhydrous sodium carbonate is a suitable chemical for preparing a standard solution (as a primary standard). The molarity of the given hydrochloric acid can be found by titrating it against the standard sodium carbonate solution prepared. The equation for the complete neutralization of sodium carbonate with dilute hydrochloric acid is

Na2CO3(aq) + 2HCl(aq) → 2NaCl(aq) + CO2(g) + H2O(l) The end-point is marked by using methyl orange as indicator.

Chemicals :

solid sodium carbonate, 0.1 M hydrochloric acid

Apparatus :

Procedure :

1. Weight out about 1.3 g of anhydrous sodium carbonate accurately using the method of “weighing by difference”.

2. Transfer the weighed carbonate to a beaker and add about 100 cm3 of distilled water to dissolve it completely. 3. After dissolving, transfer the solution to a 250.00 cm3 volumetric flask. Rinse the beaker thoroughly and

transfer all the washes into the volumetric flask. Remember not to overshoot the graduation mark of the flask. 4. Make up the solution to the mark on the neck by adding water. 5. Pipette 25.00 cm3 of sodium carbonate solution to a clean conical flask. 6. Add 2 drops of methyl orange indicator to the carbonate solution. 7. Titrate the carbonate solution with the given dilute hydrochloric acid until the colour of solution just changes

from yellow to orange. 8. Repeat the titration two times.

Calculation :

Calculate the molarity of the sodium carbonate solution prepared and the molarity of the hydrochloric acid.

Results :

Questions :

1 What is the meaning of “weighing by difference”? 2. Suggest one method other than using acid-base indicator to detect the end point of an acid-alkali titration.

Experiment :

2

Title :

Redox titration – ethane-1,2-dioic acid vs potassium permanganate

Objective :

To prepare a standard solution of ethane-1,2-dioic acid (oxalic acid) and use it to standardize a solution of potassium permanganate. Optional Activity Standardization of Hydrogen peroxide solution used the standardized potassium permanaganate solution.

Introduction :

Permanganate ion, MnO4-(aq), is a strong oxidant. Since permanganate ion is intensely coloured and its

reduction product, Mn2+(aq), is almost colourless, a self-indicating titration is possible. The addition of the first drop

of permanganate solution in excess imparts a pink colour to the solution. In strongly acidic medium, permanganate undergoes a 5 electrons reduction to manganese(II) ion: MnO4

-(aq) + 8H+

(aq) + 5e- → Mn2+(aq) + 4H2O(aq)

Ethane-1,2-dioate (oxalate) ions are oxidized according to

C C

O O

O--O(aq) → 2CO2(g) + 2e-

Potassium permanganate does not oxides oxalates in cold solution. A temperature of about 70ºC is necessary to cause the reaction to occur rapidly. Formula mass of ethane-1,2-dioic acid (H2C2O4·2H2O) is 126. A 0.05 M solution of ethane-1,2-dioic acid is to be prepared by weighing out about 1.5 g of the acid and made up to 250.0 cm3 of solution on a standard volumetric flask. Optional Activity Potassium permanganate KMnO4(aq) also reacts with hydrogen peroxide H2O2(aq) to form colourless products. Therefore, the concentration of a H2O2(aq) solution can also be determined by titrating it against the potassium permanaganate solution just standardized.

Chemicals :

3M H2SO4(aq), 0.02M KMnO4(aq), H2C2O4·2H2O(s), 0.01 M H2O2(aq)

Apparatus :

Procedure :

[Hazard Warning: oxalate is poisonous.] 1. Weigh out accurately about 1.5 g ethane-1,2-dioic acid, dissolve in water and make up to 250.00 cm3. 2. Pipette 25.00 cm3, using a pipette filler, of the prepared solution into a conical flask. Add about 25 cm3 of 3 M

H2SO4 and heat the mixture on a tripod to about 60ºC (The temperature can be estimated by using your fingers. Swirl the flask for about 10 seconds, if the side of the flask is still just too hot to be touched, the temperature of the liquid is approximately correct.)

3. Titrate with the potassium permanganate solution till a permanent pink coloration is observed. Heat the solution occasionally to maintain the temperature at 60ºC, after the addition of every 5 cm3 of KMnO4(aq). Swirl the flask continuously throughout titration.

4. Repeat the titration for two times. Optional Activity Design a procedure to standardize a solution of hydrogen peroxide using the potassium permanganate solution just standardized. Note : Heating is not required in this reaction.

Calculation : Calculate the concentration of the potassium permanganate solution in molarity (M) and in gdm-3.

Optional Activity Calculate the concentration of the hydrogen peroxide solution provided.

Results :

Questions :

1. Ethane-1,2-dioic acid is used as the primary standard in the titration. What is primary standard ? What criteria a chemical need to fulfill to be a good primary standard. Other than the ethane-1,2dioic acid and anhydrous calcium carbonate, give an example of primary standard for each of the following :

i. acid ii. base iii. oxidizing agent iv. reducing agent 2. Why is it that potassium permanganate is used as a titrant in ACIDIC MEDIUM? 3. In the above titration, what conditions should be maintained? 4. Close observation of the solution during the titration may reveal the evolution of small gas bubbles. What gas

or gases are formed? 5. Dichromate ion is also a strong oxidant. Do you think that dichromate titrations could also be self-indicating?

Half-reaction of dichromate is: Cr2O7

2-(aq) + 14H+

(aq) + 6e- → 2Cr3+(aq) + 7H2O(l)

orange green 6. For the titration between potassium permanaganate and hydrogen peroxide, write a balanced equation for the

reaction. Which chemical acts as a reducing agent and which acts as an oxidizing agent in the reaction, explain your answer briefly.

Experiment :

3

Title :

Analysis of Two commercial Brands of Bleaching Solution

Introduction : Sodium chlorate(I) (sodium hypochlorite) forms the basis of most commercial bleaches. The amount of this active ingredient available can be estimated by the following method. In this analysis, the sodium chlorate(I) is allowed to react with an excess of potassium iodide solution in the presence of acid, liberating iodine, which is then titrated against standard sodium thiosulphate solution. The reactions involved are:

ClO-(aq) + 2I-

(aq) + 2H+(aq) → I2(aq) + H2O(l) + Cl-

(aq) I2(aq) + 2S2O3

2-(aq) → 2I-

(aq) + S4O62-

(aq)

Chemicals : Commercial bleach (2 brands), 1 M KI, dilute H2SO4, 0.050 M Na2S2O3, starch indicator (freshly prepared)

Apparatus : Titration apparatus, measuring cylinder

Procedure :

1. Pipette 10.00 cm3 of the bleach into a clean 250.00 cm3 volumetric flask. Make up to the mark using water. 2. Pipette 25.00 cm3 of this solution into a conical flask, add to it 10 cm3 of 1 M potassium iodide solution and 10

cm3 of dilute sulphuric(VI) acid. [Hazard Warning: Bench dilute sulphuric(VI) acid is corrosive.] 3. Titrate this against the standard sodium thiosulphate solution. Add 1cm3 of freshly prepared starch indicator

when the reaction mixture turns light brown and continue to titrate to the end-point. Repeat your titration once.4. Repeat steps (1) to (3) with another brand of bleach.

Calculation : For each brand, work out (a) the amount of the active ingredient available in gdm-3. (b) the cost per gram of this compound.

Results :

Concentration of standard sodium thiosulphate solution provided : __________ M Brand A Trade Name: _______________ Manufacturer: _______________ Price: __________ Volume : __________

Brand B Trade Name: _______________ Manufacturer: _______________ Price: __________ Volume : __________

Questions :

1. According to your results, which of the two brands of bleach is a better buy? 2. Why should potassium iodide be present in excess? 3. What is the function of the dilute sulphuric(VI) acid ? 4. Bleaching solution may deteriorate for two main reasons. One is the attack by carbon dioxide in air according

to the equation: 2ClO-

(aq) + CO2(aq) → CO32-

(aq) + Cl2(aq) What is the other possible reason? 5. The starch indicator should not be added too early. Why?

Experiment :

4

Title :

Analysis of Aspirin Tablets

Objective :

To determine the percentage by mass of aspirin in aspirin tablets

Introduction : Aspirin is an analgesic and antipyretic drug. Analgesics are drugs which relieve pain. Antipyretics are drugs which lower body temperature. The main constituent of aspirin tablets is 2-ethanoylhydroxybenzoic acid (acetyl-salicylic acid, CH3COOC6H4COOH, m.m. 180.0). Aspirin passes unchanged through the acidic conditions in the stomach but is hydrolysed to ethanoate (acetate) ions and 2-hydroxybenzoate (salicylate) ions by the alkaline juices in the intestines. Salicylates lower body temperature rapidly and effectively in feverish patients (antipyretic action), but have little effect if the temperature is normal. They are also mild analgesics, relieving certain types of pain such as headaches and rheumatism.

CC

C

CC

CC

OH

OOC

OCH3

H

H

H

H

Although the toxic dose from salicylates is relatively large, their uncontrolled use could be dangerous. Single doses of 5 to 10 grams of salicylate have caused death in adults; and 12 grams taken over a period of twenty-four hours produces symptoms of poisoning. Pharmaceutical manufacturers are required by law to state on the packaging the amount of each active ingredient in their products. In this experiment a consumer survey on the amount of the active ingredient (2-ethanoyloxybenzoic acid, or o-acetylsalicylic acid) in different commercial brands of aspirin tablets is carried out, to see whether the manufacturers' claims are justified. 2-ethanoyloxybenzoic acid can be readily hydrolysed, using a known excess of sodium hydroxide, into the sodium salts of two weak acids, ethanoic acid and 2-hydroxybenzoic acid. The excess amount of sodium hydroxide is then estimated by a back titration with standard sulphuric(VI) acid. The equation for the hydrolysis reaction is:

CH3CO2C6H4CO2H + 2NaOH(aq) → CH3CO2Na(aq) + HOC6H4CO2Na(aq) + H2O(l) Phenol red (pH range 6.8 – 8.4) is most suitable for this titration due to the presence of the salts of two weak acids, though phenolphthalein is also satisfactory for the present purpose.

Chemicals : standard 0.5M NaOH, H2SO4 (about 0.05 M), a standard sodium carbonate solution, 2 brands of aspirin tablets, phenol red indicator, methyl orange indicator,

Apparatus : Titration apparatus, balance

Procedure :

[Hazard Warning: 0.5 M sodium hydroxide and 0.05 M sulphuric(VI) acid are irritant.] Do the Step 1 and 2 of Part II first. Part I – Standardization of sulphuric acid 1. Design a procedure to standardize the sulphuric acid provided using a standard sodium carbonate solution. Use

methyl orange as the indicator. Part II – Analysis of aspirin tablets of 2 brands (Pair up with another student and each member choose a brand). 1. Weigh accurately two aspirin tablets into a 250 cm3 conical flask. 2. Initiate the hydrolysis of the aspirin by adding 50.00 cm3 of 0.5 M sodium hydroxide from a pipette, diluting

with 25 cm3 of water. Warm the flask gently for ten minutes to complete the hydrolysis. Cool down the flask thoroughly.

3. After cooling, transfer the reaction mixture with washings to a 250 cm3 volumetric flask. Dilute to the mark with distilled water. Shake the flask to mix well.

4. Titrate 25.00 cm3 aliquots of the diluted reaction mixture with the standard sulphuric(VI) acid provided, using phenol red as indicator until the colour changes from red to orange.

5. Repeat the titration 6. Calculate the mass of 2-ethanoyloxybenzoic acid in each tablet and compare the results with the

manufacturer's specification.

Questions :

1. What is back titration ? When should it be used instead of direct titration ? 2. Determine which brand of aspirin gives the best value for money.

Experiment :

5

Title :

Enthalpy of formation of calcium carbonate

Objective :

To determine the enthalpy of formation of calcium carbonate

Introduction :

The enthalpies of reaction between calcium and hydrochloric acid calcium carbonate and hydrochloric acid are determined experimentally. By mean of an energy cycle, the enthalpy of formation of calcium carbonate can then be calculated. The following data will be needed : H2(g) + ½O2(g) → H2O(l) ∆Hf[H2O(l)] = -286 kJmol-1 C(s) + O2(g) → CO2(g) ∆Hf[CO2(g)] = -393 kJmol-1

Chemicals :

powdered calcium carbonate, 2M hydrochloric acid, calcium metal

Apparatus :

Thermometer, polystyrene cup, plastic cover

Procedure :

Work in groups of 2. A. Reaction of calcium carbonate with dilute hydrochloric acid 1. Weigh out accurately about 2 g of dry powdered calcium carbonate directly into a clean dry plastic cup. 2. Place about 50 cm3 of the hydrochloric acid in a measuring cylinder. Record its temperature every 15 seconds.

At exactly the second minute, pour the acid on the carbonate in the plastic cup, cover the cup with the cover provided.

3. Stir gently with the thermometer and record the temperature every 15 seconds until the sixth minute. 4. Tabulate your results. 5. Plot the graph of temperature of the liquid against time. 6. Calculate the no. of mole of limiting reagent used and the amount of heat evolved, hence the molar heat of

reaction between calcium carbonate and dilute hydrochloric acid. B. Reaction of calcium with dilute hydrochloric acid 1. Weigh out accurately about 0.5 g of calcium metal. 2. Using a measuring cylinder, place 50 cm3 of the hydrochloric acid in a plastic cup. Stirring carefully with the

thermometer provided, measure the temperature of the acid every 15 seconds. Tabulate your results. 3. At exactly the second minute add the metal to the acid and continue taking the temperature of the liquid every

15 seconds until the sixth minute, stir the liquid throughout the measurement. 4. Plot the graph of temperature of the liquid against time. 5. Calculate the no. of mole of limiting reagent used and the amount of heat evolved, hence the molar heat of

reaction between calcium and dilute hydrochloric acid. C. Determination of enthalpy of formation of calcium carbonate With the data provided and the results from part A and B, use Hess's law to calculate the enthalpy of formation

of calcium carbonate. Hints : 1. What is the equation for the formation of calcium carbonate from its elements under standard

conditions? [equation 1] 2. Write an ionic equation for the reaction that has taken place in A [ equation 2 ] 3. Write an ionic equation for the reaction which has taken place in B. [equation 3 ] 4. Draw an energy cycle linking equation 1, 2 and 3 together.

Results :

Specific heat capacity of water = 4.17 Jg-1K-1 Density of water = 1.00 gcm-3

Questions :

1. What does the term standard enthalpy of formation mean? 2. What are the “standard conditions” of thermochemistry? 3. Why is the exact concentration of the hydrochloric acid unimportant? 4. In this experiment, did we assume that there was no heat lost to the surrounding ? 5. What other assumption(s) has you made in this experiment ?

Experiment :

6

Title :

Enthalpy Change of Hydration of Copper(II) Sulphate(VI)

Introduction : This experiment enables an approximate determination of the enthalpy change of hydration of copper(II) sulphate(VI) to be calculated using Hess's law. The enthalpy changes when known masses of anhydrous copper(II) sulphate(VI) and copper(II) sulphate(VI)-5-water crystals are respectively dissolved in the same quantity of water are measured. From this, the molar enthalpy changes are calculated in each case.

If ∆H1 = molar enthalpy change of hydration of CuSO4 ∆H2 = molar enthalpy change of solution of anhydrous CuSO4 ∆H3 = molar enthalpy change of solution of CuSO4·5H2O then ∆H1 = ∆H2 - ∆H3

Chemicals : Anhydrous CuSO4, CuSO4·5H2O

Apparatus : Polystyrene foam cup, -10 – 110 ºC thermometer, 100 cm3 measuring cylinder, 400 cm3 beaker, cotton wool

Procedure :

A. To determine the enthalpy change of solution of copper(II) sulphate(VI)-5-water crystals

1. Put 12.5 g of anhydrous copper(II) sulphate(VI) in a weighing bottle. 2. Loosen the cotton wool and pack a 400 cm3 beaker with the loosened cotton wool with a polystyrene cup

fitted in the beaker. 3. Using a measuring cylinder, measure 100 cm3 water into the polystyrene foam cup fitted in a beaker.

Record its temperature every 15 seconds. At exactly the second minute, pour the copper(II) sulphate(VI) in the plastic cup, cover the cup with the cover provided.

4. Stir gently with the thermometer and swirl the cup occasionally. Record the temperature every 15 seconds until the sixth minute.

5. Weigh the emptied weighing bottle again to determine the actual mass of copper(II) sulphate(VI) added. B. To determine the enthalpy change of solution of anhydrous copper(II) sulphate(VI)

Repeat part A using 8.0 g of anhydrous copper(II) sulphate(VI) instead of the hydrated salt. [Hazard Warning: Anhydrous copper(II) sulphate(VI) is harmful.]

Questions :

1. State Hess's law. 2. Calculate the molar enthalpy change of solution of anhydrous copper(II) sulphate(VI), ∆H2 Assume that the

specific heat capacity of the solution is 4.2 kJ kg-1K-1. 3. Calculate the molar enthalpy change of solution of hydrated copper(II) sulphate(VI), ∆H3. 4. Work out the molar enthalpy change of hydration of copper(II) sulphate(VI). 5. What assumptions have you made in calculating ∆H2 and ∆H3? 6. What are the sources of errors in this experiment and how could they be minimized?

Experiment :

7

Title :

Estimation of strength of hydrogen bond

Introduction : Breaking or formation of intermolecular hydrogen bonds between molecules in liquids would cause an enthalpy change when the liquids are mixed. This experiment is to investigate such enthalpy changes and to measure approximate strengths of hydrogen bonds formed between molecules of ethanol and those between molecules of trichloromethane and ethyl ethanoate using simple calorimetric methods.

Chemicals : Ethanol, cyclohexane, ethyl ethanoate, trichloromethane

Apparatus : 10 cm3 and 25 cm3 measuring cylinders, 50 cm3 beaker, 250 cm3 beaker, -10 – 110 °C thermometer, cotton wool

Procedure :

[Hazard Warning: Ethanol, cyclohexane and ethyl ethanoate are flammable, trichloromethane is harmful, and tetrachloromethane is toxic.] A. To measure the strength of hydrogen bond formed between ethanol molecules

1. Using a measuring cylinder, add 5 cm3 of ethanol into an insulated 50 cm3 beaker. Measure the temperature of the liquid.

2. Then add 20 cm3 of cyclohexane to the ethanol in the beaker, mix well and record the lowest temperature attained.

3. From the temperature drop estimate the hydrogen bond strength (in kJmol-1) in ethanol. B. To measure the strength of hydrogen bonds formed between molecules of ethyl ethanoate and

trichloromethane

[Hazard Warning: Mixtures of trichloromethane and propanone have been known to explode on standing. The solvent residues from this experiment should not be disposed of into a container in which propanone is present.]

1. Measure 5 cm3 of ethyl ethanoate into an insulated 50cm3 beaker. Record its temperature. 2. Add to this 20 cm3 of trichloromethane and mix well. Record the highest temperature attained. 3. From the temperature change estimate the strength of the hydrogen bond formed between molecules of

ethyl ethanoate and trichloromethane. N.B. Actually either ethyl ethanoate or trichloromethane can be in excess. The following physical data may be useful: Specific heat capacity of glass = 0.78 kJ kg-1K-1

Liquid Formula Relative

molecular mass Density / kgdm-3

Specific heat capacity / kJkg-1K

Ethanol CH3CH2OH 46 0.81 2.44 Cyclohexane C6H12 84 0.78 1.83 Trichloromethane CHCl3 119.5 1.48 0.98 Ethyl ethanoate CH3CO2CH2CH3 88 0.90 1.92

Questions :

A. To measure the strength of hydrogen bond formed between ethanol molecules 1. Why should the beaker be insulated? 2. Is the mixing process endothermic or exothermic? 3. Account for the temperature change. 4. Explain why cyclohexane has to be used in excess in this experiment. B. To measure the strength of hydrogen bonds formed between molecules of ethyl ethanoate and

trichloromethane 1. Is the mixing process exothermic or endothermic? 2. Account for the temperature change. 3. Explain why it does not matter which liquid is used in excess.

Experiment :

8

Title :

The Methanal Clock Reaction – A Study of the Factors Affecting the Rate of a Chemical Reaction

Introduction : The reactions between methanal and hydrogensulphate(IV) ion can be represented as: H2O(l) + HCHO(aq) + SO3

2-(aq) → CH2OHSO3

-(aq) + OH-

(aq) OH-

(aq) + HSO3-(aq) → H2O(l) + SO3

2-(aq)

overall reaction: HCHO(aq) + HSO3-(aq) → CH2OHSO3

-(aq)

In the overall reaction, hydrogensulphate(IV) ion is consumed. When its concentration drops to a certain level, the pH rises sharply, and a pink colour will be observed if phenolphthalein indicator is present. By measuring the time taken for the colour of the solution to turn pink under different conditions, factors which determine the rate of the reaction can be studied.

Chemicals : 200 cm3 Solution A (0.3 M methanal), 200 cm3 solution B (0.2 M with respect to sodium hydrogensulphate(IV) and 0.05 M with respect to sodium sulphate(IV) with a little phenolphthalein added)

Apparatus : burette, boiling tube, stop watch, -10 – 110 °C thermometer, 250 cm3 beaker

Procedure :

[Note : Use 3 different burettes to measure the volume of solution A, solution B and water. Deliver the amount of the solution required to test-tube or boiling tube, then mix the solutions in one batch.] A. Effect of dilution

1. Mix the following amounts of water with 5 cm3 of solution A in a boiling tube, then add 5 cm3 of solution B to each and stir vigorously. Record the time for the pink colour to form.

Volume of water / cm3 0 2.5 5 7.5 10 12.5 15 Time / s

2. Plot a graph of volume of water against time.

B. Effect of increasing methanal concentration

1. Mix the given amounts of solution A and water in a boiling tube as listed in the table below and then add

5 cm3 of solution B. Record the time taken for the solution to turn pink.

Volume of solution A / cm3 5 7.5 10 12.5 15 Volume of water / cm3 15 12.5 10 7.5 5 Time /s

2. Plot a graph of volume of methanal against time.

3. Plot an appropriate graph to determine the order of the reaction with respect to methanal.

C. Effect of temperature

1. Prepare a solution of 20 cm3 of water and 10 cm3 of solution A in a boiling tube. Put 10 cm3 of solution B into another test-tube.

2. Immerse the boiling tube and the test tube in the hot water bath with the following temperature for 2 minutes.

3. Mix the two solutions and record the maximum temperature of the solution attained and the time taken for it to turn pink.

Approximate temperature / ºC

25 30 35 40 45

Maximum temperature reached / ºC

Time /s 5. Plot a graph of temperature against time.

N.B. In plotting a graph, usually, the independent variable is plotted as x and the dependent variable is plotted

as y.

Questions :

A. Effect of dilution 1. What effect does the amount of water have on the rate of the reaction as indicated in your graph? B. Effect of increasing methanal concentration 1. What effect does the amount of methanal have on the rate of the reaction as indicated in your graph? C. Effect of temperature 1. What effect does temperature have on the rate of the reaction as indicated in your graph?

Experiment :

9

Title :

Determination of order of rate of disproportionation of sodium thiosulphate in acidic medium by initial rate method.

Introduction : In acidic medium, thiosulphate ion disproportionate (an element undergoes oxidation and reduction at the same time) to sulphur and sulphur dioxide according to the following equation.

S2O32--

(aq) + 2H+(aq) → S(s) + SO2(g) + H2O(l)

The formation of sulphur precipitate will turn the mixture cloudy. If a paper with a cross is placed under the flask of reaction mixture, the cross will be masked by the sulphur precipitate gradually. By assuming that the concentration of the reactants, rate of disproportionation and the amount required to mask the cross are all constant in the due course of the reaction. A rate equation can be formulated as Rate0 = k[H+

(aq)]0x[S2O3

2-(aq)]0

y

1t = k’ [H+

(aq)]0s[S2O3

2-(aq)]0

y

By keeping the concentration of the acid in large excess the equation can be further simplified as

1t = k’’[S2O3

2-(aq)]0

y

ln 1t = ln k’’ + y ln [S2O3

2-(aq)]0

By plotting a graph of ln 1t vs ln [S2O3

2-(aq)]0 , the order of reaction with respect to thiosulphate ion would be

determined.

Chemicals :

Apparatus :

Experiment :

10

Title :

To Determine the Activation Energy of the Reaction between Bromide Ion and Bromate(V) Ion in Acid Solution

Introduction : During a reaction, bonds are first broken, and others are then formed. Energy is required to break certain bonds and start this process, no matter the overall reaction is exothermic or endothermic. Particles will not always react when they collide because they may not possess sufficient energy for the appropriate bonds to break. A reaction will only occur if the colliding particles possess more than a certain minimum amount of energy known as the activation energy, EA. Using the kinetic theory and probability theory, Maxwell and Boltzmann showed that the fraction of molecules with energy greater than EA Jmol-1 was given by e-EA/RT where R is the gas constant, T is the absolute temperature and e is the exponential function. This suggests that at a given temperature: Reaction rate α e-EA/RT. But as k, the rate constant for a reaction, is a measure of the reaction rate, we can write k α e-EA/RT k = A e-EA/RT

This last expression is called the Arrhenius equation. In acidic solution, the reaction between bromide ion and bromate(V) ion can be represented by

5Br-(aq) + BrO3

-(aq) + 6H+

(aq) → 3Br2(aq) + 3H2O(l) The progress of the reaction may be followed by adding a fixed amount of phenol together with some methyl red indicator. The bromine produced during the reaction reacts very rapidly with phenol (forming tribromophenol). Once all the phenol is consumed, any further bromine bleaches the indicator immediately. So, the time for the reaction to proceed to a given point may be determined.

Now, as k α reaction rate = concentration change time

time

As for the concentration change in this experiment is constant,

k α 1

time for methyl red to be bleached

k = constant

t

ln (constant

t ) = ln A + ln e-EA/RT

ln (constant) + ln 1t = ln A -

EART

ln t = [ln A - ln (constant)] - EART

so a graph of ln 1t against

1T has a gradient of -

EAR .

Chemicals : 50 cm3 solution A (0.083 M with respect to KBr and 0.017 M with respect to KBrO3), 25 cm3 0.5 M H2SO4, 50 cm3 0.01 M phenol solution, methyl red indicator

Apparatus : Beaker, boiling tube, stop watch, two -10 – 110 °C thermometers, burette

Procedure :

[Hazard Warning: 0.5 M sulphuric(VI) acid is irritant, and phenol is toxic.] 1. Place 10.0 cm3 of 0.01 M phenol solution, 10.0 cm3 solution A and 2 drops of methyl red indicator into a

boiling tube. 2. Place 5.0 cm3 of 0.5 M sulphuric(VI) acid into a second boiling tube. 3. Place both boiling tubes into a beaker of water which is maintained at about 30°C. Allow the contents of the

tubes to reach the temperature of the water bath. 4. Pour the content of the first boiling tube into the sulphuric(IV) acid, and at the same time start the stop watch.

Swirl well for 5 seconds. 5. Keep the second boiling tube in the water bath throughout the experiment. Record the time (t) taken when the

red colour fade. 6. Record also the temperature (T), to the nearest degree, of the reaction mixture at the end of the experiment. 7. Repeat steps (1) to (6), maintaining the reaction temperature at about 35 °C, 40 °C, 45 °C and 50 °C.

Calculation : Determine Ea by plotting a suitable graph. (Given: R = 8.314 J K-1mol-1) Result :

Temperature Experiment Time t / s

ln 1t

T / ºC T/K 1/T / K-1

1 2 3 4 5

Questions :

1. Give an equation for the reaction between phenol and bromine. 2. What is the use of methyl red in this experiment? 3. Should the time be taken when the methyl red starts to fade or when the methyl red becomes colourless ?

Why? 3. Based on your results, is it advisable to perform the experiment at high temperatures such as 80 °C? 4. Why is it not necessary to know how far the main reaction has proceeded at the point where the methyl red is

decolorized ? 5. The Arrhenius equation can be represented as: k = A e-EA/RT (a) Can 1/t substitute k in this equation? Why? (b) Derive an equation relating ln k and 1/T. 6. Explain why the reaction rate can be affected by temperature.

Experiment :

11

Title :

The effect of concentration changes on equilibria

Objective :

The purpose of this experiment is to find out how a system in equilibrium responds to a change in concentration of components in the mixture.

Introduction : Equilibrium I Iron(III) ions and thiocyanate ions react in solution to produce thiocyanatoiron(III), a complex ion, according to the equation : Fe3+

(aq) + SCN-(aq) d Fe(SCN)2+

(aq) pale colourless blood-red yellow Equilibrium II Bromine molecules and water molecules react in solution to produce hydrogen ions, bromide ions and hydrobromous acid (hydrobromic (I) acid) molecules, according to the equation : Br2(aq) + H2O(l) d H+

(aq) + Br-(aq) + HOBr(aq)

yellow colourless colourless colourless colourless brown Equilibrium III Dichromate(VI) ions and water molecules react in solution to produce chromate(VI) ions and hydrogen ions, according to the equation : Cr2O7

2-(aq) + H2O(l) d 2CrO4

2-(aq) + 2H+

(aq) orange yellow Equilibrium IV Bismuth trichloride molecules and water molecules react in solution to produce bismuth oxychloride precipitate and hydrochloric acid, according to the equation : BiCl3(aq) + H2O(l) d BiOCl(s) + 2HCl(aq) colourless white precipitate The appearance of the mixtures can indicate the positions of equilibria.

Chemicals : water, 0.01 M KSCN, 0.01 M FeCl3, 1M FeCl3, 1M NH4Cl, Br2(aq), H2SO4(aq), NaOH(aq), K2Cr2O7(aq), BiCl3(aq)

Apparatus :

safety spectacles, 4 test-tubes and test-tube rack, 2 teat-pipettes, spatula, glass stirring rod

Procedure :

Equilibria Reagent A Reagent B Reagent C Reagent D I 0.01M FeCl3(aq) 0.01M KSCN(aq) 1M NH4Cl(aq) 1M FeCl3(aq) II Br2(aq) H2O(l) H2SO4(aq) NaOH(aq) III K2Cr2O7(aq) H2O(l) H2SO4(aq) NaOH(aq) IV BiCl3(aq) H2O(l) H2SO4(aq) NaOH(aq) 1. Note the appearance of all the reagents. 2. Mix 1 cm3 of Reagent A to 1 cm3 of Reagent B in a test tube, shake well and note the appearance of the

mixture. 3. Add Reagent C dropwise to the mixture, shake well and note any observable change. 4. Add Reagent D dropwise to the mixture, shake well and note any observable change. 5. Repeat the procedure 2 to 4 for another equilibrium system.

Questions :

Explain any observable change observed.

Experiment :

12

Title :

Determination of the equilibrium constant by titration

Objective :

To find the equilibrium constant by titration for the reaction : Ag+(aq) + Fe2+

(aq) d Fe3+(aq) + Ag(s)

Introduction :

Chemicals :

Apparatus :

Procedure :

1. Prepare a sample of reaction mixture by adding together 70.00 cm3 each of 0.100M iron(II) sulphate and 0.100M silver nitrate solutions in a 250 cm3 conical flask.

2. Stopper the flask tightly and leave them to stand overnight. 3. Carefully pipette out into a conical flask 25.00 cm3 of the supernatant solution above the silver precipitate. 4. Titrate the 25.00 cm3 sample of solution with 0.0500 M standard potassium thiocyanate solution. When all

Ag+(aq) ions have been removed, thiocyanate ions react with iron(III) ions (also present in the equilibrium

mixture) to give a blood red color which acts as an indicator. 5. Repeat the titration and calculate Kc for this reaction.

Results :

Questions :

1. Write an equation for the reaction between silver(I) ion and thiocyanate ion. 2. Write an equation to account for the formation of the blood red color at the end point. 3. Why was it necessary to stopper the conical flask containing 25 cm3 of iron(II) sulphate and silver nitrate

solution tightly? 4. Why is it necessary to leave the solutions to stand overnight? 5. Suggest sources of error for this experiment

Experiment :

13

Title :

Determination of Ka of a weak acid using buffer solutions

Introduction : For a weak acid, HA(aq) d H+(aq) + A-

(aq)

the dissociation constant is given by Ka = [H+

(aq)][X-(aq)]

[HA(aq)]

If the acid is a weak acid and its salts are strong electrolytes, then for a mixture of the acid with one of its salts, it is possible to assume that [HA(aq)] = total acid concentration [A-

(aq)] = total salt concentration

Chemicals :

Apparatus :

Procedure :

1. Put approximately 10 cm3 portions of the different buffer solutions, in numerical order, in 8 test tubes in a rack.

2. Add 3 drops of universal indicator to each tube and shake to mix. 3. Add 5.0 cm3 of 0.2 M ethanoic acid and 5 cm3 of 0.1 M NaOH solution using burettes, into a clean, dry test

tube to produce a half-neutralized acid solution. Add 3 drops of universal indicator. Determine the approximate pH of this solution by comparing with the colour set.

4. Determine the dissociation constant for the acid.

Questions :

The pH for a half-neutralized acid equals pKa. Account for this.

Experiment :

14

Title :

Acid-base titration by double indicator method

Objective :

To determine the proportions of sodium carbonate and sodium hydroxide in a mixture solution using Double indicator method.

Introduction :

Indicator to exhibit completion of reaction I. NaOH(aq) + HCl(aq) → NaCl(aq) + H2O(l) Any indicator II. Na2CO3(aq) + HCl(aq) → NaHCO3(aq) + NaCl(aq) Phenolphthalein III. NaHCO3(aq) + HCl(aq) → NaCl(aq) + H2O(l) + CO2(g) Methyl orange The neutralization of a strong alkali by means of a strong acid can be followed by the used of indicator. Sodium hydrogen carbonate solution is definitely alkaline to it. If acid is added to a mixture of sodium hydroxide and sodium carbonate in solution, using phenolphthalein as indicator, then the pink colour of the indicator is discharged when reactions I and II are completed. If methyl orange is now added and a further quantity of acid is also added, the amount required will be that necessary to complete reaction III. But 1 mole of sodium hydrogen carbonate formed from one mole of sodium carbonate and hence the amounts of acid required for reactions II and III will be the same. Suppose the volume of acid needed to reach end-point as indicated by the phenolphthalein is x cm3 and the additional volume of acid to reach the end point as indicated by the methyl orange is y cm3 . Then the volume of acid reacting with sodium hydroxide is x - y cm3.

Chemicals :

Apparatus :

Procedure :

Pipette 25.00 cm3 of the mixture solution into a conical flask. Add two drops of phenolphthalein indicator. Titrate with the given standard hydrochloric acid (0.100M) until the pink colour is just discharged. Note the burette reading and add a few drops of methyl orange and a further quantity of acid until the yellow colour of the methyl orange changes to orange. Repeat the above procedure 3 times more.

Results :

Questions :

1. What is the mean volume of the titrant with phenolphthalein indicator. (x cm3) 2. What is the mean volume of the titrant with methyl orange indicator. (y cm3) 3. Calculate the concentration of the solution with respect to sodium carbonate and sodium hydroxide. Give your

answer in molarity and gdm-3.

Experiment :

15

Title :

To Study Acid-Base Titration by pH Measurements

Introduction : The voltage across an electrochemical cell is affected by the pH of electrolyte in which electrodes are immersed in. By measuring the voltage across two reference electrodes immersed in buffer solutions with different pH, a calibration graph can be constructed with the pH vs the voltage (V) across the 2 electrodes. Using the calibration curve obtained, the voltmeter calibrated functions as a pH meter to trace the change in pH in the course of an acid-base titration. Consequently, titration curve can be constructed to determine the end point of the titration.

Chemicals :

0.2M standard hydrochloric acid, 0.2 M sodium hydroxide, 0.2 M ethanoic acid, 0.2 M ammonia solution, buffer solutions with pH = 2.2, 4.2, 5.9, 8.0, 10.0, 12.0, copper(II) sulphate solution

Apparatus :

Burette, Pipette, Retort stand & clamp, Magnetic stirrer, magnetic bar & remover, DC 3 V and connecting wires, Antimony electrode, copper electrode, digital multimeter, crocodile clips, Wash bottle, Salt bridge–use filter paper strips

Procedure :

Part I : Construction of calibration curve 1. Calibrate the voltmeter by using the buffer solutions provided, with the pH = 2.2, 4.2, 5.9, 8.0, 10.0, 12.0.

Construct a calibration curve using the voltage measured as the vertical axis against the pH of the buffer solutions as the horizontal axis.

Part II: Titration of Sodium Hydroxide Solution with Hydrochloric Acid 1. Pipette 25.00 cm3 of 0.2 M (approximate) sodium hydroxide solution into a beaker. 2. Set up the apparatus as demonstrated. Set the magnetic stirrer in motion. 3. Run in from a burette, 5.00 cm3 of 0.2 M (accurate) hydrochloric acid. Determine the voltage of the cell. 4. Continual add 5.00 cm3 portions of acid, taking the voltage reading after each addition, until a total of

20 cm3 has been added. Then add the acid in 1.00 cm3 portions until a total of 30.00 cm3 has been added. Then revert to 5.00 cm3 portions until a total of 40.00 cm3 has been added, determining the voltage throughout.

5. Obtain the pH values of the solutions from the calibration graph. Plot the titration curve of the pH of the solution (vertical axis) against the volume of hydrochloric acid added.

Part III: Titration of Ammonia Solution with Hydrochloric Acid 1. With the original calibration curve, repeat the titration using 0.2 M (approximate) ammonia solution

instead of sodium hydroxide solution and construction of titration curve to determine the concentration of ammonia solution provided.

Wash the burette and refill the burette with the 0.2M (approximate) ethanoic acid provided for Part IV and V. Part IV: Titration of Ammonia Solution with Ethanoic Acid 1. With the original calibration curve, repeat the titration using the 0.2 M (approximate) ammonia provided

and construction of titration curve to determine the concentration of ethanoic acid provided by assuming that the concentration of ammonia determined in Part III is accurate.

Part V: Titration of Sodium Hydroxide Solution with Ethanoic Acid 1. With the original calibration curve, repeat the titration using 0.2 M (approximate) sodium hydroxide

solutions and construction of titration curve to determine the concentration of ethanoic acid provided again by assuming that the concentration of sodium hydroxide determined in Part II is accurate.

Calculation : Part II : Use the titration curve to determine the accurate concentration of the sodium hydroxide solution provided.

Part III: Use the titration curve to determine the accurate concentration of the ammonia solution provided. Part IV: Use the titration curve to determine the accurate concentration of the ethanoic acid provided. Part V: Use the titration curve to determine the accurate concentration of the ethanoic acid provided again.

Questions :

1. What differences can be observed among the three titration curve obtained ? Explain the differences briefly. 2. Compare and comment on the concentrations of ethanoic acid determined in Part IV and Part V.

Experiment :

16

Title :

Investigation of the Oxidation States of Vanadium

Introduction : Vanadium has a range of oxidation states from 0 to +5. The following experiments are designed to identify the colours of vanadium solutions for the various oxidation states, the ease with which the oxidation states can be formed and the stability of the vanadium compounds made.

Chemicals : Ammonium metavanadate, 0.1 M FeSO4, 0.1 M Fe2(SO4)3, 0.1 M KI, 0.1 M KBr, 0.1 M I2 in KI, aqueous SO2, copper powder, zinc powder, dilute H2SO4, 0.05M Na2S2O3(aq)

Apparatus : Bunsen burner, 8 test tubes and rack, filter funnel, filter paper, stoppers

Procedure :

[Hazard Warning: Ammonium metavanadate is toxic, iron(III) sulphate(VI) is irritant, and bench dilute sulphuric(VI) acid is corrosive.] A. Various oxidation states obtained from ammonium metavanadate (NH4VO3)

Colours of different Vanadium compound at different oxidation states. Oxidation State Colour in Solution Vanadium (V) Yellow Vanadium (IV) Blue Vanadium (III) Green Vanadium (II) Violet

1. Dissolve a small spatula of ammonium metavanadate in 60 cm3 of dilute sulphuric(VI) acid to make a V(V)

solution in a conical flask. 2. Transfer 10 cm3 of the solution to a test tube and stopper it for later use. 3. Add a little zinc powder to the remaining solution in the conical flask with constant swirling and warming to

reduced V(V) to Vanadium compounds with lower oxidation states. When the solution turns blue, filter 10 cm3 of the solution to a test tube and stopper it for later use.

4. Repeat the heating and filtering procedure for the green and violet solutions.

N.B. Extra zinc powder should only be added when most of the zinc powder in the solution have been dissolved.

The violet solution can be air oxidized, keep it away from air.

B. Identification of colours of vanadium compounds and a study of these compounds Make a large copy of Table 1 on your raw data sheet. Using the chart of Eo values and the colour of vanadium

containing ions at different oxidation states, predict the likely experimental result when substances are mixed as in the table. This should cover the expected oxidation state of vanadium in the final solution and the likely observations.

(Be careful about the mixing of coloured solutions even without chemical reactions. A blue and yellow solution can give a green one on mixing. If iodine is present, the colour of other products can only be seen if the iodine is converted to colourless iodide ions by sodium thiosulphate or removed in a solution of 1,1,1-trichloroethane.)

3. Test them out by performing the actual experiments. Record the results also in the table.

Table 1 Substances to be mixed Prediction Experimental Result V(V) + FeSO4(aq) V(V) + KI V(V) + SO2(aq) V(V) + Cu V(V) + Zn V(IV) + V(II) V(V) + V(III) V(IV) + V(III) V(IV) + Fe2(SO4)3 V(III) + Fe2(SO4)3 V(IV) + KI V(V) + KBr V(III) + I2 V(IV) + I2

Questions :

1. What is the oxidation state of vanadium in ammonium metavanadate? 2. What are the colour changes observation when V(V) is being reduced to V(II), explain the colour change

briefly ? 3. Does all experiment result match your prediction ? If not, explain why does the predication fail.

Experiment :

17

Title :

The variation of boiling point with composition for a binary liquid mixture

Objective :

The purpose of this experiment is to construct boiling-point / composition curves for mixtures of two different liquids.

Introduction :

Works in groups of 3, each member choose one of the following binary mixture 1. trichloromethane and methyl ethanoate; 2. ethanok and cyclohexane; 3. propan-1-ol and propan-2-ol. You measure the boiling-points of these mixtures at various compositions and construct boiling-point composition curves from your results. For simplicity, you plot volume composition rather than mole fraction - the curves have the same general shape. Share your data with your members in the group and plot a boiling point composition curve showing the liquid composition curve only for each of the binary mixture.

Chemicals :

trichloromethane CHCl3, methyl ethanoate CH3CO2CH3, ethanol C2H5OH, cyclohexane, C6H12 propan-1-ol, CH3CH2CH2OH, propan-2-ol, CH3CH(OH)CH3

Apparatus :

safety spectacles, quick-fit apparatus, thermometer, 0-100 °C, ± 0.1 °C, anti-bumping granules, Bunsen burner, gauze and tripod

Procedure :

1. Set up a suitable assembly for reflux, with the flask positioned over a tripod and gauze. It is important that the thermometer is positioned so that it will dip into the liquid mixture but it must not touch the walls of the flask.

2. Choose one of the three systems shown and ensure that the other students in your group investigate the other

systems. In the remaining procedure steps we refer to the components of each system as A or B, as shown.

System Component A Component B 1 trichloromethane, CHCl3 methyl ethanoate, CH3CO2CH3 2 ethanol, C2H5OH cyclohexane, C6H12 3 propan-1-ol, CH3(CH2)2OH propan-2-ol, CH3CH(OH)CH3 3. Transfer 10.0 cm3 of A from the burette to the pear-shaped flask containing a small spatula measure of anti-

bumping granules. Heat the flask gently until the liquid just begins to boil. 4. Record the boiling-point of A. 5. Turn off the Bunsen burner and allow the apparatus to cool for about two minutes. 6. Measure 2.0 cm3 of liquid B from the burette into a test-tube and pour this liquid down the condenser into the

pear-shaped flask. 7. Reheat the flask gently until the liquid mixture just boils and record its boiling-point. 8. Repeat stages 5, 6 and 7 with further additions of 2 cm3 portions of component B until a total of 10 cm3 of B

has been added. 9. Allow the apparatus to cool and ask your teacher how you can safely dispose of the mixture. 10. Repeat the experiment from step 3, this time starting with 10 cm3 of liquid B in the flask and adding 2 cm3 of

component A after each boiling point determination until a total of 8 cm3 of A has been added. 11. Plot a graph of boiling-point (y-axis) against percentage composition by volume (x-axis). 12. Collect the results for the other two systems from your members and plot the results in exactly the same way.

Results :

Volume of Component A / cm3

10 10 10 10 10 10 0 2 4 6 8

Volume of Component B / cm3

0 2 4 6 8 10 10 10 10 10 10

Volume % of Component A

Boiling point / ºC Questions :

1. For each boiling-point/composition curve, sketch an approximate vapour pressure composition curve showing both liquid composition curve and vapour composition curve.

2. Study the vapour pressure curves and classify each mixture as ideal or as showing positive or negative

deviations from ideality.

Experiment :

18

Title :

Distribution equilibrium

Objective :

The purpose of this experiment is to determine the value of the distribution coefficient for the equilibrium that exists when ammonia is distributed between water and 1,1,1-trichloroethane.

Introduction :

In this experiment you shake some ammonia solution with 1,1,1-trichloroethane (relative density 1.34) to establish equilibrium, and then determine the concentration of ammonia in each solvent by titration. This enables you to calculate the distribution coefficient, KD. NH3(tce) d NH3(aq) (tce = 1,1,1-trichloroethane)

KD = [NH3(aq)][NH3(tce)]

Chemicals :

50 cm3 ammonia solution 1M NH3, 1,1,1-trichloroethane CH3CCl3, methyl orange indicator solution, hydrochloric acid 0.010M HCl (standardized), hydrochloric acid, 0.50 M HCl (standardized)

Apparatus :

safety spectacles, measuring cylinder, separating funnel 150 cm3, 2 beakers 100 cm3, pipette 10 cm3 and safety filler, 2 conical flasks 150 cm3, wash-bottle of distilled water, white tile, burette 50 cm3 and stand

Procedure : Part 1 (Group experiment) 1. Pour about 100 cm3 of ammonia solution into a separating funnel. 2. Pour about 100 cm3 of 1,1,1-trichloroethane into the same separating funnel. 3. Holding the tap firmly in position with one hand and the stopper with the other, shake the separating funnel

vigorously for about ten seconds. Release the pressure inside by loosening the stopper for a moment. 4. Continue shaking for about half a minute, releasing the pressure every ten seconds. Set aside until two layers

separate. Part 2 (Individual experiment) 5. Transfer the lower organic layer to a beaker. Using a dry pipette, transfer 10.0 cm3 to a conical flask. 6. Add about 20 cm3 of distilled water, a few drops of indicator and titrate the mixture with 0.010 M HCl until

the yellow solution just changes to red and remains red after shaking. (It may take a few moments for all the ammonia to transfer from the organic layer and react with the acid.)

7. Titrate two more 10 cm3 samples. 8. Transfer the aqueous layer to a beaker. Rinse the pipette thoroughly, transfer 10.0 cm3 to a flask, add about 20

cm3 of water and a few drops of indicator solution, and titrate to the endpoint with 0.5 M HCl. 9. Titrate two more 10 cm3 samples.

Results :

Questions :

1. Calculate the average concentration of ammonia in the organic layer. 2. Calculate the average concentration of ammonia in the aqueous layer. 3. Calculate the distribution coefficient.

Experiment :

19

Title :

Separation of Amino Acids by Paper Chromatography

Introduction : A mixture of unknown amino acids can be separated and identified by means of paper chromatography. The positions of the amino acids in the chromatogram can be detected by spraying with ninhydrin, which reacts with α-amino acids to yield highly coloured products.

Chemicals : 2% ammonia, propan-2-ol, aluminium foil, ninhydrin spray (2% solution of ninhydrin in ethanol), separate solutions of 0.05 M glycine, tyrosine, leucine, and aspartic acid in 1.5 % HCl in 4 test tubes, an unknown containing one to four of the above amino acids at a concentration of about 0.05 M each in 1.5 % HCl

Apparatus : Capillary tube, Whatman chromatography paper, beaker, oven

Procedure :

1. Mix 10 cm3 of 2 % ammonia solution with 20 cm3 of propan-2-ol in a clean, 500 cm3 beaker, and cover tightly with a piece of aluminium foil. This would be used as the solvent for the experiment. [Hazard Warning: Propan-2-ol is flammable.]

2. On a clean sheet of chromatography paper with size about 12 cm by 22 cm, mark a light pencil line parallel to the bottom and about 1.5 cm away (Figure 1). Along this line mark ten light crosses ("×") at intervals of about 2 cm. Label each cross as shown in Figure 1. ("U" represents the unknown amino acid mixture.)

3. Using capillary tubes, place a small amount of each appropriate solution on its two positions along the line on

the chromatography paper. Avoid getting the spot on the paper larger than about 2 mm in diameter. Let the paper dry for a few minutes in air. Add a second portion of the unknown to one of its two positions, to make certain that sufficient quantities of each component of the unknown will be present for good visual observation when the paper is developed.

4. Roll the paper into a cylindrical form. Staple the ends together in such a fashion that they do not touch each other (Figure 2). Otherwise the solvent will flow more rapidly at that point and form an uneven solvent front.

5. When the spots on the cylindrical paper are dry (it may be necessary to place the paper in an oven at about 100°C for a short time), place it carefully in the beaker of solvent, and cover carefully and tightly with the aluminium foil. Make sure that the paper does not touch the wall of the beaker.

Let the solvent rise up the paper for at least 1.5 hours. If the time is shorter, the components may not be sufficiently separated for easy identification. Remove the paper and place it upside down on the desk top to dry. When most of the solvent has evaporated, open the cylinder by tearing it apart where it was stapled and hang it in a fume cupboard. Spray the paper lightly but completely with a solution of ninhydrin, and leave the paper in the fume cupboard until the spray solution is dry.

[Hazard Warning: The ninhydrin solution should be kept off the body because it reacts with proteins in the skin to form a rather long-lasting purple discoloration. The teacher should ensure that students wear laboratory gowns, gloves and safety spectacles in carrying the experiment.]

Place the paper in an oven at 100°–110°C for about 10 minutes, or until all the spots have developed. 8. Circle each spot with a pencil, and measure the distance each spot traveled (use the centre of the spot for

measurement). Measure the distance the solvent traveled at each position, and calculate the Rf values for each amino acid. Determine the composition of the unknown by visual comparison of spot colours and by comparing the Rf values.

Rf value = distance travelled by the colour spot

distance travelled by the solvent front

Questions :