G.C.E. (Advanced Level)

Chemistry

Practical Instructional Manual

(for the syllabus implemented from 2009)

Department of Science, Health and Physical Education

Faculty of Science and Technology

National Institute of Education

G.C.E. (Advanced Level)

Chemistry

Practical Instructional Manual

(for the syllabus implemented from 2009)

©National Institute of Education

1st Print - 2011

Department of Science, Health and Physical Education

Faculty of Science and Technology

National Institute of Education

-

ii

Guidance Prof. W.M. Abeyrathne Bandara, Director General,National Institute of Education (NIE)Mr. Lal Wijesinhge Assistant Director General(Curriculum Development), NIE

Direction Mr. C.M.R. Anthony, Director, Department of Science, Health andPhysical Education, NIE

Coordination Mr. A. D. A. de Silva, Chief Project Officer, NIE

Resource Committee: Internal : Mr. C. M.R. Anthony - Director, Department of Science, Health &

Physical Education.Mr. A.D.A de Silva - Chief Project Officer, NIEMr. L. K. Waduge - Chief Project Officer, NIEMrs. Malini Ragavachari - Project Officer, NIE

External : Prof. H.D. Gunawardana - University of ColomboProf. W.D.W Jayathilaka - University of Sri JayawardanapuraMr. M.A.P. Munasinghe - Chief Project Officer, NIE (Retired)Miss. S. Velupillai - Hindu College, Colombo 04Mr. S. Thillainathan - Hindu Ladies College, Colombo 06Mr. R.N.T. Bandara - Dharmarajah College, Kandy.Mr. P.G. S Peramuna - Pinnawela Central College, RabukkanaMr. Bandula Ranasinghe - Wesly College, BorallaMiss. C.A.N. Perera - Devi Balika Vidyala, BorallaMrs. N. Thirunavukarasu - Hindu Collge, Colombo 4. (Retired)Mrs. M. S. Athukoralla - Taxila Central Collge, HoranaMrs. Sriyani Mallika - Dharamapala Vidyala - Pannipitiya

Type Setting : Ms.R.R.K. Pathirana - NIE

Cover : Miss Vignesan Seshamani - Gr. 10 -Hindu Collge, Colombo 04

Web Site : www.nie.lk

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vi

Contents

Page1. Identification of the properties of cathode rays .. .. 012 Observing components of the visible range of the electromagnetic spectrum.. 043 Observation of the atomic spectrum of hydrogen .. .. 044 Experimental determination of the molar volume of a gas .. .. 055 Experimental determination of the relative atomic mass of magnesium .. 076 (I) Experimental comparison of enthalpy of dissolution .. .. 08

(II) Experimental determination of the enthalpy of neutralization ofan acid / base

(III) Experimental determination of enthalpy of a displacementreaction

7 Experimental verification of Hess’s law .. .. 108 Comparison of the reactions of s and p block metals with water and acids .. 129 Examination of the solubility of salts formed by s and p block elements .. 1310 Examination of the thermal stability of nitrates, carbonates and bicarbonates

of s - block elements .. .. 1411 Preparation of allotropic forms of sulphur .. .. 1512 Preparation of sulphur dioxide and examination of its properties .. .. 1613 Preparation of chlorine and examination of the properties of halogens .. 1814 Identification of halides .. .. 1915 Observing the reactions of Copper(II) and Cobalt(II) salts with

hydrochloric acid .. .. 2016 Observation of colours of manganese in oxidation states of +2, +4, +6

and +7 .. .. 2117 Observing reactions of alkenes and alkynes .. .. 2218 Examination of the properties of alcohols .. .. 2419 Examination of the properties of phenol .. .. 2620 Tests for aldehydes and ketones .. .. 2721 Examination of the properties of carboxylic acids .. .. 2822 Experimental determination of the effect of concentration on the

reaction between Mg and acids .. .. 2923 Experimental determination of the effect of concentration of reactants

on the reaction between sodium thiosulphate and nitric acid .. .. 3124 Experimental determination of the effect of concentration of reactants on the

reaction between iron (III) ions and potassium iodide .. .. 3325 Experimental study of the characteristics of a dynamicequilibrium

using Fe3+ / SCN- system .. .. 34

v

26 Experimental study of the effect of temperature on the equilibrium systemof NO2 and N2O4 .. .. 35

27 Experimental determination of the solubility product of calcium hydroxide .. 3628 Experimental determination of the distribution coefficient for the

distribution of ethanoic acid in water and butanol .. .. 3729 Preparation of an indicator using flowers provided and experimental

determination of its pH range .. .. 3930 Experimental determination of the acidic/basic/neutral nature of aqueous

solutions of salts by testing pH. .. .. 4031 Determination of approximate pH of a given solution using pH indicators .. 4132 Measuring the conductivity of a given solution using the Wheatstone

bridge principle .. .. 4333 Experimental determination of the variation of conductivity with concentration,

temperature and the nature of the electrolyte .. .. 4434 Making different types of electrodes and measuring their electrode

potentials using the potentiometer .. .. 4535 Measurement of the electromotive force of various cells using

the potentiometer .. .. 4636 Comparison of efficiency of cells .. .. 4637 Experimental determination of the decomposition potential of a

simple electrolytes .. .. 4838 Preparation of hydrogen and oxygen gases using electrolysis of water. .. 4939 Investigation into the corrosion of iron as an electrochemical process .. 5040 Electroplating with nickel and copper .. .. 5141 Tests for selected cations .. .. 5242 Tests for selected anions .. .. 5443 Recrystallization using pure solvents/solvent mixtures .. .. 5744 Experimental determination of the moisture content of a sample of food .. 5845 Determination of the stoichiometry of the reaction between barium

chloride and sulphuric acid (Precipitation method) .. 5946 Determination of the stoichiometry between salicylic acid and iron(III) ions

(Colorimetry method) .. .. 6447 Experimental determination of the concentration of a thiosulphate solution using

potassium iodate .. .. 6648 Separation of components of chlorophyll using paper chromatography .. 6749 Experimental determination of the oxygen percentage in air by volume .. 6850 Experimental determination of the alkalinity of a sample of water .. 7251 Experimental determination of the dissolved oxygen in a sample of water .. 7552 Experimental determination of the total dissolved solid (TDS) in a

sample of water .. .. 78

1

Practical No. 01

Identification of the properties of cathode rays(Teacher demonstration)

Expected learning outcomes :1. Identifies the elements of the electric circuit used to operate a cathode ray tube.2. Observes the properties of cathode rays.

Materials and Equipments : Cathode ray tubes Induction coil 6 V direct current electric source

(a lead accumulator or a nickel/ iron cell or a power pack or if electricity is notavailable use a powerpack)

Connecting wires A switch

Instructions : You are provided with special Crooke’s tubes in order to test the properties of cathode

rays. Several such tubes are represented in Figure 1 (a). There is a gas under low pressure in these tubes. A very large potential difference has

to be applied across the two terminals. Figure 1 (b) indicates some of the circuits usedfor this purpose.

(i) (ii)

(iii)Figure 1 (a) Cathode ray tubes

Flourescent screen

2

(iv) (v)

(vi)Figure 1 (b) Few circuits that can be used to operate cathode ray tubes

Induction coil

Vibrator

Power supply unit

Copper Sheet

N.B. : Do not allow any part of the body to come in contact with high voltage instruments

or connecting wires, when a current flows through a circuit. Keep the key of the circuit open to cut off the flow of current, when the cathode

ray properties are not being tested.

Experiment I: Set up the equipment as indicated in the circuit diagram (iv) or the circuit diagram

(v). Here the cathode ray tube shown in (i) is used. Close the key in the circuit andallow a current to flow through. Turn the screw A left or right, so that the vibratorvibrates to and for producing electrical sparks between A and B.

Then on the wall D, opposite the cathode, a light green glow appears (if this doesnot happen change the direction of the flow of current by changing the terminals ofthe battery). This occurs as a result of the cathode rays. Under these circumstances,the rays that emerge from the cathode are known as cathode rays.

Now open the key and stop the flow of current through the circuit. Bring the metalliccross in the tube to the vertical position and close the circuit again allowing the flowof current. What do you observe? Open the key and stop flow of the current.

3

Experiment II : Connect the cathode ray tube given in figure (ii) (with paddle wheel) to the circuit. See

that the tube is placed horizontally. Close the key and allow the current to flow. What canyou observe in the cathode ray tube?

Open the key and stop the flow of current. In order to change the direction of flow ofcurrent, change the connections to the battery terminals. Close the key and observewhat happens inside the tube.

Experiment III : Open the key to stop the flow of the current and connect the cathode ray tube given in

figure (iii) to the circuit. In order to subject the rays to an electrical field, arrange theequipment according to figure (vi). Place a copper strip above parallel to the cathode raytube and arrange the circuit as shown so that it becomes the negative plate.

Now, allow the current to flow through the circuit and observe what happens inside thetube Open the circuit and stop the flow of current. Then place the copper strip below thetube and arrange the circuit again. Close the circuit, allow a current to flow and observewhat happens inside the tube.

Experiment IV: Use the same cathode ray as tube used for Experiment III. Here the cathode ray tube

is subjected to a magnetic field. Arrange the circuit as in figure (iv) or (v) and connect thecathode ray tube (iii) to the circuit. Observe what happens in the cathode ray tube.

Move the south pole of a magnet closer and perpendicular to the zinc sulphide screenand observe what happens in the same way. Then, bring the north pole of the magnetcloser and observe what happens.

Note : Based on the observations of the above experiments, describe the properties of cathode

rays.

4

Practical No. 02

Observing components of the visible range of the electromagneticspectrum

Expected learning outcomes :1. Observes and records the nature of the visible range of the electromagnetic spectrum

using a suitable device.

Materials and Equipments : A glass prism A source of white light A screen

Instructions : Disperse a beam of white light using a prism and obtain the spectrum on as screen.

Let the students observe how the colours are resolved.

Practical No. 03

Observation of the atomic spectrum of hydrogen

Expected learning outcomes:1. Gets to know the basic facts that should be considered in constructing an electrical

circuit.2. Presents the hydrogen spectrum as a piece of evidence in support of the existence

of electrons in different energy levels.3. Observes hydrogen spectrum and states its main characteristics.

Materials and Equipments : Equipment used to identify the properties of cathode rays Hydrogen discharge tube Diffraction grating

Instructions : Connect the hydrogen gas discharge tube to the terminals of the circuit used to study

the properties of cathode rays. Observe the emergent radiation with the equipmentprovided to observe the spectrum.

N.B.The hydrogen gas discharge tube essential to study the atomic spectrum of hydrogenhas not been provided to schools as at present. Therefore, for the time being it issufficient to give a description using a coloured picture of the line spectrum ofhydrogen.

5

Practical No. 04

Experimental determination of the molar volume of a gas

Expected learning outcomes:1. Acquires the skill of collecting gases and knowledge about the methods of

measuring their volumes.2. Determines the volume occupied by one mole of oxygen gas under laboratory

conditions.

Materials and Equipments : Solid potassium permanganate Cotton wool Apparatus indicated in the diagram Measuring cylinder (500 cm3) Thermometer

Instructions : Take into a boiling tube a small amount of lightly heated potassium permanganate (Why

heated?) and push a cotton wool plug through the mouth of the tube. (Why do you needcotton wool?). Then, weigh the boiling tube with the contents. Set up the apparatus asshown in either the Figure 4 (a) or Figure 4 (b) and gently heat the boiling tube.

It is sufficient to collect about 350 - 400 cm3 of oxygen gas. If equipments are arrangedas in Figure 4 (a), an equal volume of water is displaced to the vessel. Stop heating whenthis volume of water is 350 - 400 cm3 and allow the apparatus to cool to the roomtemperature (Why?). After equalising water levels, run all the water in the rubber tube tothe measuring cylinder. Measure the volume of displaced water.

If the apparatus are set as in Figure 4 (b), equalise the water levels in the measuringcylinder and the trough and measure the volume of the gas collected.

Note room temperature and atmospheric pressure. Weigh the boiling tube with its content. Using relevant readings calculate molar volume of oxygen under room temperature and

pressure and under S.T. P.

N.B.1. The cotton plug should not be pushed far down as it could get burned.2. If Figure 4 (b) set up is used, the end of the delivery tube should reach almost the top

of the measuring cylinder.

6

Note :Discuss about the errors of the experiments, how they can be minimized and theassumptions made in the calculation.

Cotton woolOxygen

KMnO4

HeatWater

Figure 4 (b)

Cotton wool

KMnO4Heat Water Water

Figure 4 (a)

7

Practical No. 05

Experimental determination of the relative atomic mass of magnesium

Expected learning outcomes :1. Determines under the room temperature and pressure, the volume of hydrogen gas

liberated by a known mass of magnesium.2. Calculates the relative atomic mass of magnesium using the amount of moles of

hydrogen liberated.

Materials and Equipments : Magnesium ribbon, 0.02.-0.04 g (about 3.5 cm) Burette Beaker Dilute hydrochloric acid, 25 cm3

Water

Instructions :

Take a piece of magnesium ribbon about 3.5 cm long and clean with a sand paper.Weigh it accurately. Add 25 cm3 of dilute hydrochloric acid to a burette and then fill itcompletely with water by adding water slowly along the side of the burette. Cover thepiece of magnesium loosely with cotton wool, make it into a lump and trap at the upperend of the burette.

Nearly half fill a 250 cm3 beaker with water. Close the upper end of the burette with thethumb, quickly turn it upside down and dip the mouth in the water in the beaker. Clampthe burette vertically. Before gas bubbles rise up, open the stopcock of the burettequickly and bring the liquid level to the final mark of the burette. Close the stopcock andtake this as the initial liquid level. When all the magnesium has reacted, take the finalliquid level in the burette. The volume of hydrogen liberated is given by the differencebetween these two levels.

Magnesium ribbonBurette

Burette

BeakerMagnesium ribbonWater

Dilute hydrochloric acid

Cotton wool

Figure 5 (a)

8

Note : Under the room temperature and pressure under which the experiment was carried out,

calculate the amount of moles of hydrogen gas liberated by the mass of magnesium used. Through relevant computation, find the relative atomic mass of magnesium. Discuss the errors of the experiment, the ways of minimizing them and assumptions made

in the calculation.

Practical No. 06

(I) Experimental comparison of enthalpy of dissolution(II) Experimental determination of the enthalpy of neutralization of

an acid / base(III) Experimental determination of enthalpy of a displacement

reaction

Expected learning outcomes :1. Develops the skill of measuring temperature change in a chemical reaction.2. Compares the enthalpies of dissolution.3. Determines the enthalpy change of an acid - base neutralization reaction.4. Determines the enthalpy of a displacement reaction experimentally.

Materials and Equipments : 50 cm3 of 2.0 mol dm-3 sodium hydroxide solution 50 cm3 of 2.0 mol dm-3 potassium hydroxide solution 50 cm3 of 2.0 mol dm-3 nitric acid 50 cm3 of 2.0 mol dm-3 hydrochloric acid 50 cm3 of copper sulphate solution (prepared by dissolving 5 g in 50 cm3 of

water) Beakers, 100 cm3

A 250 cm3 beaker Measuring cylinders, 100 cm3

A thermometer Insulating material A glass rod

9

Instructions :

Experiment (I)Take a few water soluble salts available in laboratory and prepare 250 cm3 of aqueoussolutions of each. Note the temperature variation. Using Q mc T calculate thedissolution enthalpies and compare.

Experiment (II)Take 50 cm3 of 2 mol dm-3 sodium hydroxide solution into a beaker and measure itstemperature. Into another beaker take 50 cm3 of 2 mol dm-3 hydrochloric acid solutionand take its temperature (Wash the thermometer well before putting it into the acid solutionafter taking the temperature of the base). Add at once, the solution in one beaker to theother and take the maximum steady temperature. Calculate the enthalpy of neutralisationof the reaction using Q mc T . Repeat the experiment using potassium hydroxide andnitric acid solutions.

Experiment (III)Take 50 cm3 of the copper sulphate solution to a beaker and measure its temperature.Add 1g of zinc powder to it at once. Stir the solution with a glass rod and measure themaximum temperature. Using Q mc T , calculate the enthalpy change associatedwith the reaction of one mole of zinc.

N.B.Conducting above experiment in insulated vessels, more accurate results can be obtained.

Note :

Discuss errors in each experiment and assumptions made in calculations. Discuss measures that can be taken to minimize the errors.

10

Practical No. 07

Experimental verification of Hess’s law

Expected learning outcomes :1. Shows that one system could be converted to another through two different

pathways.2. Indicates an enthalpy change by an enthalpy diagram.3. Shows by means of an enthalpy diagram that the enthalpy changes in the

two pathways are similar and verifies Hess’s law.

Materials and Equipments : 2.0 mol dm-3 hydrochloric acid, 250 cm3

Four 400 cm3 beakers 250 cm3 measuring cylinder Thermometer (0 - 200 0C) Solid sodium hydroxide, 20 g

Instructions :In this experiment a known mass of sodium hydroxide is converted to aqueous sodiumchloride through two pathways.

Method I(a) Measure 125 cm3 of 2.0 mol dm-3 hydrochloric acid into a beaker and note

its temperature. Measure 125 cm3 of water into another beaker and note itstemperature also. Now add the 125 cm3 acid to the water, mix thoroughly andnote the temperature again (you will see that the temperature change in this dilutionis negligible).

(b) Measure 250 cm3 of 1.0 mol dm-3 hydrochloric acid into a beaker and note itstemperature (t1). Weigh 10 g of solid sodium hydroxide and put it into a beakeras quickly as possible. Then add the acid into this beaker and dissolve the sodiumhydroxide with stirring. Note the maximum temperature of the solution (t2).

Method II(a) Measure 125 cm3 of water into a beaker and note its temperature (t3). Weigh

10 g of solid sodium hydroxide into another beaker; add water into it, stirthoroughly until sodium hydroxide dissolves and note the maximum temperature(t4) of the solution.

(b) Measure 125 cm3 of 2.0 mol dm-3 hydrochloric acid into a beaker and note itstemperature (t5). Note the temperature (t6) of the sodium hydroxide solutionprepared in II (a) and add the hydrochloric acid solution quickly, stir thoroughlyand note the maximum temperature (t7).

11

Reactions that occur in both methods are represented below.

Figure 7 (a)

Determine the enthalpy change associated with the reaction by using appropriatecalculations.

Represent the relevant enthalpy changes by an enthalpy diagram and verify Hess’slaw.

Note : Discuss the experimental errors and assumptions made in calculations.

2 mol dm-3 NaOH 125 cm3

1 mol dm-3 HCl 250 cm3

2 mol dm-3 HCl 125 cm3

water 125 cm3

NaOH 10 g

1 mol dm-3 NaCl250 cm3

2 mol dm-3HCl125 cm3

solid NaOH10 g

Final substancesInitial substances

H3H2

H1

The enthalpy changehere is neglegibly small'

Method I

Method II

12

Practical No. 08

Comparison of the reactions of s and p block metals with water andacids

Expected learning outcomes:1. Observes the reactions of s - block elements with

(a) Air ( b) Water (c) Dilute acids2. Identifies the horizontal and vertical variation of the reactivity of s - block and p -

block metal elements.

Materials and Equipments : s - block elements available in the laboratory and aluminium Dilute hydrochloric acid or dilute sulphuric acid Phenolphthalein indicator A trough of water Test tubes

Instructions :• Examine how the s - block elements and aluminium are stored up in the

laboratory. Cut a piece of sodium into two and expose the cut surfaces to the air for about five

minutes and observe. Take a 5 cm long strip of magnesium, apply sand paper on a part of it . Expose it to

air for about 15 minutes and observe what happens. Heat a piece of magnesium ribbon in air. What do you observe? Fill a trough half - full with water, add a little of phenolphthalein indicator. Clean

a tiny piece of sodium with a filter paper and put it to water (as sodium reactsvery vigorously with water care should be taken). What are your observations?

Put a piece of magnesium ribbon cleaned with sand paper into a test tube containingwater. Add a little phenolphthalein also. What are your observations?

Put a clean piece of mangesium ribbon to another test tube containing water and alittle phenolphthalein. Heat and observe.

• Take a clean piece of magnesium ribbon into a test tube and add dilutehydrochloric acid. What are your observations? (As the reaction is extremelyvigorous with sodium, do not attempt it.)

13

Practical No. 09

Examination of the solubility of salts formed by s and p blockelements

Expected learning outcomes :1. Examines and compares the solubility of s and p block salts.

Materials and Equipments :• Water soluble salts of Be, Mg, Ca, Sr, and Al• Sodium salts (or potassium salts) given in the table

• Test tubes

Instructions :• Prepare solutions of relevant salts.• Add dropwise the sodium salt solution to solutions of the group II salts and make

observations.

0.1

mol dm

-3

solu

tio

n $

1 c

m3

1 m

ol dm

-3

NaC

l

/1 c

m3

1 m

ol dm

-3

NaB

r

/1 c

m3

1 m

ol dm

-3

NaI

/1 c

m3

1 m

ol dm

-3

NaO

H

/1 c

m3

1 m

ol dm

-3

Na

2C

O3

/1 c

m3

1 m

ol dm

-3

NaH

CO

3

/1 c

m3

1 m

ol dm

-3

NaN

O2

/1 c

m3

1 m

ol dm

-3

NaN

O3

/1 c

m3

1 m

ol dm

-3

Na 2

SO

3

/1 c

m3

1 m

ol dm

-3

Na 2

SO

4

/1 c

m3

Be2

+

Mg

2+

Ca2

+

Sr2

+

Ba2

+

Al3+

Table 9 (a) 9 (a

14

Practical No. 10

Examination of the thermal stability of nitrates, carbonates andbicarbonates of s - block elements

Expected learning outcomes :1. Finds a pattern in the thermal stability of the nitrates, carbonates and

bicarbonates of the s - block elements.

Materials and Equipments :• Nitrates and carbonates of the following metals

and bicarbonates of group I metals • LimewaterSodium • Pyrex test tubesPotassium • Test tubesMagnesium • Delivery tubeCalcium • Test tube holderBarium • Bunsen burnerStrontium • Litmus

Instructions :• Carbonates and bicarbonates : Take about 1g of solid carbonate/bicarbonate

into a pyrex tube and fix a stopper carrying a delivery tube. Insert the free end ofthe delivery tube into about ½ cm of limewater in a test tube. Heat the pyrex tube.

Is there a change in the limewater inside the test tube? Do this experiment with allthe carbonates and bicarbonates separately. Compare the time taken to change thecolour of limewater.

• Nitrates: Take each nitrate into a pyrex tube and heat. Hold a glowing splinter atthe mouth. If a coloured gas is liberated, compare the time taken for the liberationof the coloured gas.

15

Practical No. 11

Preparation of allotropic forms of sulphur

Expected learning outcomes :1. Acquires the skills of preparing rhombic, monoclinic and plastic forms of sulphur.2. Shows that there is a variation in the allotropic forms of the same element.

Materials and Equipments :• Powdered sulphur• Three evaporating dishes (medium size)• Carbon disulphide

Instructions :Preparation of rhombic sulphur

• Add a little of sulphur powder to an evaporating dish. Also add carbon disulphideadequate to dissolve all the sulphur added. Allow to stand till carbon disulphide slowlyevaporates. Observe the separated crystals of rhombic sulphur.

Preparation of monoclinic sulphur• Add about 5 g of powdered sulphur to an evaporating dish and heat till the suphur

melts. Allow to cool slowly. Pierce the crust formed on the surface in two places with aglass rod and pour out the liquid underneath. Observe the needle - shaped crystals ofmonoclinic sulphur formed beneath the crust and on the sides of the dish.

Preparation of plastic sulphur• Heat sulphur nearly to boiling and add to cold water. Observe the rubbery plastic

sulphur formed.

Note : Even boiling tubes can be used in place of evaporating dishes in this experiment.

Note :Rhombic and monoclinic sulphur are the crystalline allotropes of sulphur. Plastic sulphur andsulphur obtained as a milky precipitate in many instances are amorphous forms of sulphur.

16

Practical No. 12

Preparation of sulphur dioxide and examination of its properties

Expected Learning outcomes :1 Prepares sulphur dioxide in the laboratory.2 Understands oxidising and reducing properties of sulphur dioxide.3 Develops the skill of testing for sulphite ions.

Materials and Equipments :• Copper turnings• Concentrated sulphuric acid• Litmus papers• Potassium permanganate• Dilute nitric acid• Dilute sulphuric acid• Coloured flower petals• Potassium dichromate• Sodium sulphite

Instructions :

• Tip copper turnings to a boiling tube to a height of about 0.5 cm and add concentratedsulphuric acid to height of nearly 1 cm. Fix a delivery tube with a bung and heatshowly. Sulphur dioxide gas begins to evolve.

• Sulphur dioxide is also produced by the reaction between the sulphites and dilute acids.Therefore. this method can also be used to prepare sulphur dioxide in the labortory.

N.B.Do not touch sulphuric acid as it may cause burns. If excess copper and sulphuric acidwere used the contents may spill over on heating. Therefore, care should be taken not totake more copper and sulphuric acid than the amounts given.

Heat

Boiling tube

Figure 12 (a)

conc.

17

As soon as the experiment is over, disconnect the equipment, empty the contents of theboiling tube into the sink and open the tap.

Do the experiment quickly to avoid excessive fuming.

Do the following tests for sulphur dioxide gas

1. Observe the odour of the gas2. (a) Bring the gas into contact with flower petals. what happens to the colour of the

flower petals ?(b) Pass the gas into moist blue and red litmus papers. What happens to the colour

of the litmus papers ?3. (a) Pass sulphur dioxide into potassium dichromate/ potassium chromate solution

acidified with dilute sulphuric acid ?What do you observe ?

(b) Pass sulphur dioxide gas througth a solution of potassium permanganateacidified with dilute sulphuric acid.What are your observations ?Why dilute hydrochloric acid is not used in tests 3(a), 3(b) and 3(c) instead ofsulphuric acid for acidification ?

4. Pass hydrogen sulphide gas through a solution of sulphur dioxide. What is the colour ofthe precipitate formed?

Note :Present explanations for the above observations using oxidising and reducing proper-ties of sulphur dioxide.

18

Practical No. 13

Preparation of chlorine and examination of the properties ofhalogens

Expected learning outcomes :1. Prepares and collects chlorine using simple equipment.2. Observes the reactions shown by halogens.

Materials and Equipments :• Solid potassium permanganate • Litmus• Conc. hydrochloric acid • Boiling tubes• Sodium hydroxide solution • Delivery tube• Ammonia solution • Gas jars (with lids)• Copper turnings or gauze • Bunsen burner• Iron wool or powder

Instructions :Take a little (about 2 g) potassium permanganate to a boiling tube; add conc. hydrochloricacid; fix a delivery tube; collect several cylinders of the liberating gas by upwarddisplacement of air (Boiling tubes may be used instead of gas cylinders). Close the cylinderswith their lids.

What is the smell and the colour of the gas liberated?

N.B. : • Do not heat potassium permanganate with concentrated hydrochloric acid. It is not healthy

to inhale chlorine gas. If there is a fume cupboard, prepare and collect the gas inside it.

Reactions of halogens:• Hold wet litmus papers or wet, coloured petals to

the gas.• Invert a cylinder ( or boiling tube) of chlorine in a

trough of water. Observe the level of water andtest the solution with coloured materials.

• Pass chlorine to a test tube containing a little sodiumhydroxide solution and test the solution with acoloured petal.

• Put heated copper gauze and iron wool into chlorinegas cylinders separately and observe.

• Add little water ( 2 or 3 drops), shake and observe.• Add 2 or 3 drops of ammonia solution to a chlorine

gas cylinder and observe what happens.

Concentratedammonia

Chlorine gascylinder

Lid

Figure 13 (a)

19

Practical No. 14Identification of halides

Expected learning outcomes:1. Develops the skill of identifying chloride, bromide and iodide ions when they exist

separately in solution.

Materials and Equipments :• Solid sodium chloride • Ammonia solution• Solid sodium bromide • Dilute nitric acid• Solid sodium iodide • Starch solution

(could use any salt containing these halide ions) • Litmus (red and blue)• Silver nitrate solution • Boiling tubes• Conc. sulphuric acid • Test tubes• Lead acetate or lead nitrate solution • Glass rod• Chlorine water • Bunsen burner• Chloroform or carbon tetrachloride

Instructions :Do the following tests separately for chlorides, bromides and iodies. Compare the observations.

Test 1Take a little solid halide into a test tube; add a little conc. sulphuric acid and heat the testtube.i. Note the colour of the gas liberated.ii. Test the gas with wet blue litmus and wet red litmus papers.iii. Test all the gases with a filter paper soaked in starch solution.iv Test all the gases with a glass rod dipped in ammonia solution.

Do the following tests with aqueous solutions of the halides.

Test 2Take the three halide solutions separately into test tubes. Acidify each with dil. nitric acid.Add silver nitrate solution, a little at a time and observe whether precipitation occurs. If sowhat are their colours?Add ammonia solution to these precipitates and note down the observations.

Test 3Add lead nitrate solution to the three halide solutions in separate test tubes. Observe thecolours of the precipitates formed. Dilute the precipitates with water and heat, then allowto cool again. Note the observations.

Test 4Add a little chlorine water to the three halide solutions separately and then a few drops ofchloroform. Shake the tubes well and observe. What colour do you see in the chloroformlayer?Explain the observations for each test.

20

Practical No. 15

Observation of reactions of cobalt(II) and copper(II) ions withhydrochloric acid

Expected learning outcomes :1. States that the ‘d’ block elements form complex ions.2. Acquires the skill of preparing a few complex compounds.3. States that the hydrated ions of those complex compounds are coloured.

Requirements :• A dilute copper sulphate solution• A solution of cobalt(II) ions• Conc. hydrochloric acid• Hydrogen peroxide solution• Dilute and concentrated ammonia solutions

Instructions :• Add reagents to the respective solutions of ions as indicted in the following tables and

observe the changes.

For Cu2+ ions:

For Co2+ ions:

• Give the formulae of complex ions formed by Co2+ and Cu2+ ions giving respectivecolours in the above instances.

Cu2+ (aq)

Colour of thesolution

Add a diluteammonia solutondropwrise

Add a concentratedammonia solutiondropwise

Add concentratedhydrochloric aciddropwise

Co2+ (aq)

Colourof thesolution

Add a diluteammoniasolutiondropwise

Add a concen-trated ammoniasolutiondropwise

Add conc. HClsolutiondropwise

Add hydrogenperoxidedropwise

21

Practical No. 16

Observation of colours of manganese in oxidation states of +2, +4,+6 and +7

Expected learning outcomes :1. Identifies the colours of ions of manganese in different oxidation states.2. Acquires the skill of convertting ions of manganese from one oxidation state to another

Materials and Equipments :• Dilute aqueous potassium permanganate solution• Dilute sulphuric acid• Concentrated potassium hydroxide or sodium hydroxide solution• Hydrogen peroxide solution• Concentrated hydrochloric acid solution

Instructions :• Put about 2 cm3 of a dilute potassium permanganate solution to a boiling tube, acidify

slightly with dilute sulphuric acid and add concentrated potassium hydroxide solution tillthere is a colour change. Observe the colour change and add about two drops ofhydrogen peroxide to the solution. Observe the colour change. Also aaa concentratedhydrochloric acid to the solution and observe the changes in colour.

Note :• The +7 oxidation state of Mn, which is present in an aqueous solution of KMnO4 is

purple in colour. Concentrated potassium hydroxide converts +7 oxidation state ofmanganese to +6 state. It is green in colour.

2KMnO4(aq) + 2KOH(aq) 2K2MnO4(aq) + H2O(l)

• Hydrogen peroxide reduces K2MnO4 to MnO2 of + 4 state whcih is brown . Whenconc. HCl is added to MnO2, it gets converted to Mn2+ ions. Mn2+ ions are almostcolourless.

MnO2(s) + 4HCl(aq) MnCl2(aq) + 2H2O(l) + Cl2(g)

So, different oxidation states of manganese exhibit different colours.

22

Practical No. 17

Observing reactions of alkenes and alkynes

Expected learning outcomes :

1. Developes the skill of preparing ethene and ethyne in the laboratory.2. Observes the reactions of ethene and ethyne.3. Compares the properties of ethene and ethyne.

Materials and Equipments :• Ethanol• Alumina• Calcium carbide• Alkaline potassium permanganate• Acidified potassium permanganate• Ammoniacal cuprous chloride• Cotton wool

Instructions :The apparatus for the preparation of ethene is indicated in the Figure 17 (a).

Figure 17 (a)

(A rubber tube fixed to a glass tube and stopper could be used instead of the deliverytube).

Take ethanol to a height of about 2cm to a boiling tube. With the help of a glassrod insert sufficient cotton wool to soak ethanol. Admit about 1g of alumina and pushit half - way down the tube.

Fix the tube to a support and arrange the equipment as shown. Then heat alumina.Collect the gas liberated into several boiling tubes by the downward displacement ofwater. They contain ethene.

Alumina

Ethene

Water

Deliverytube

Heat

Cotton woolsoaked withethanol

23

1. Light the gas by bringing a lighted splinter close to the test tube. What is the colourof the flame? As soon as the flame is extinguished, add a few drops of lime water,close the test tube and shake well.

2. (a) Add a few drops of bromine water.(b) Add a few drops of alkaline potassium permanganate.(c) Add a few drops of acidified potassium permanganate.

3. (a) Add ammoniacal silver nitrate solution.(b) Add ammoniacal cuprous chloride solution.

Note :Another way of doing the test 2 and 3 is to take about 1 cm3 of the reagent into a test tubeand passing the gas into it.

Take some calcium carbide into a boiling tube and arrange the apparatus as shown inFigure 17 (b).Add few drops of water at a time .

1. Take about 2.5 cm3 of each of the following regents to a test tube and pass ethyne gasas shown in Figure 17 (b). Alkaline potassium permanganate Acidified potassium permanganate Ammoniacal cuprous chloride Ammoniacal silver nitrate

Note the observations. After the tests are carried out, wash the tubes thoroughly.

2. Hold a flame to the end of the delivery tube and allow the liberated ethyne to burn.

Calciumcarbide

Water

Figure 17 (b)

24

Practical No. 18

Examination of the properties of alcohols

Expected learning outcomes :

1. Understands the reactivity of the - OH group in alcohols by reacting them with various reagents.2. Explains the observations associated with the reactions of alcohols with given reagents.

Materials and Equipments :• Ethanol• Methanol• Benzyl alcohol and other available alcohols• Acetic acid• Sodium salicylate or salicylic acid• Acidified potassium dichromate• Acidified potassium permanganate• Alkaline potassium permanganate• Sodium carbonate• Litmus paper

Instructions :• Take about 1 cm3 of each of the three alcohols into separate test tubes and do the

following activities. Compare the results obtained.• Test the alcohol with blue and red litmus paper.• Add solid sodium carbonate and observe.• Add a freshly cut piece of sodium and observe. Hold a glowing splinter to the

mouth of the tube. What is the gas liberated?For these tests dry alcohol should be used. Why?

When the piece of sodium is completely dissolved,vapourise the product ( as indicted in thediagram). What is the remainder? Add a fewdrops of water to the remainder and shake. Putblue and red litmus separately and observe. Add adrop of phenolphthalein to it. What is the colourobtained? What is the reason for it?

WaterAlcohol

Heat

Figure 18 (a)

25

• (a) Take about 1 cm3 of the alcohol and add about 1cm3 of glacial acetic acid. Addabout 5 drops of conc. sulphuric acid and warm. Pour the product to cold waterin a test tube and observe the smell.

(b) Take about 1 cm3 of the alcohol and add about 0.5 g of sodium salicylate orsalicylic acid . Add about 5 drops of conc. sulphuric acid and warm. Pour theproduct into a test tube or a beaker containing cold water and observe the smell.

• . Take 1cm3 each of acidified potassium dichromate, acidified potassiumpermanganate and alkaline potassium permanganate to three test tubesseparately. Add three drops of the alcohol to each. Repeat the same test withthe other alcohols also. What colour changes do you observe in the solutions?How can you explain the colour changes? Examine the smell of the products.

N. B. : As the alcohols are inflammable, keep the bottles away from the burners.

26

Practical No. 19

Examination of the properties of phenol

Expected learning outcomes:

1. Understands the acidic property of phenol.2. Reports the observations of special tests for phenol.

Materials and Equipments :• Phenol• Sodium hydroxide solution• Dilute hydrochloric acid• Sodium carbonate solution• Liquid bromine• Dilute ammonia solution• Dilute ferric chloride solution• Litmus

N.B. :Phenol when comes into contact with the skin causes burns. Therefore it should be handledwith care.

Instructions :Take a little phenol separately into test tubes and do the following tests.• Add 1cm3 of water and shake. Test the solution with red and blue litmus papers.• Add 1cm3 of sodium hydroxide solution and shake thoroughly. Also add 1cm3 of

dilute hydrochloric acid to the same solution.• Add 1cm3 of sodium carbonate solution.

Note the observations.

Special tests for phenol :1. Take 1cm3 liquid bromine into a test tube add a crystal (or two drops) of phenol.

2. Take 1cm3 of dilute ferric chloride solution into a test tube and add one drop ofdilute ammonia. Add one crystal (or one drop) of phenol to this neutral ferric chloridesolution.

N.B.• If phenol comes into contact with the skin it should be washed with a very dilute

solution of bromine. Bromine reacts immediately with phenol and removes it.

• Some carbon compounds form coloured complexes with ferric ions in neutral medium.Phenol gives a purple colour, p- nitrophenol gives a red colour, o-nitrophenol doesnot give a special colour.

27

Practical No. 20

Tests for aldehydes and ketones

Expected learning outcomes :1. Makes aldehydes and ketones to react with common oxidising agents and makes

observations.2. Understands that aldehydes can be easily oxidised to carboxylic acids by common

oxidising agents but ketones cannot be so easily oxidised.

Materials and Equipments :• Methanal• Ethanal• Benzaldehyde• 2- Propanone• Acetophenone• Acidified potassium permanganate• Acidified potassium dichromate• Ammoniacal silver nitrate solution• 2,4 - Dinitrophenylhydrazene• Fehling’s solution (A and B)

Instructions :• Do the following tests with the aldehydes methanal, ethanal and benzaldehyde and

the ketones, propanone and acetophenone. Take about 1cm3 of each, note theobservations and tabulate them.

1. Mix with a little water. Examine whether there is a separation of layers.2. (a) Add about two drops of acidified potassium permanganate solution.

(b) Add about two drops of acidified potassium dichromate solution.Observe the colour changes in each.

3. Add about 3 cm3 of Fehling solution ( made by mixing equal volumes of FehlingA and B) and heat the mixture. What are the compounds that give a brick - redprecipitate?

4. Add about 5 cm3 of ammoniacal silver nitrate ( Tollen’s reagent) and heat thesolution in a water bath. Note the observations.

5. Add a few drops of 2, 4 - dinitrophenyl hydrazine (2, 4 - DNP). Note the observations.

N.B.Use 2, 4 - DNP reagent in aqueous solution for water - soluble crboxyl compounds. Forcarboxyl compounds that are insoluble in water, 2. 4 - DNP in methanol should be used.

28

Practical No. 21

Examination of the properties of carboxylic acids

Expected learning outcomes :1. Shows by simple tests that compounds containing the carboxylic group show

acidic properties.2. Prepares esters by reacting carboxylic acids with alcohols.3. Demonstrates that methanoic acid has reducing properties.

Materials and Equipments :• Methanoic acid • Dilute hydrochloric acid• Ethanoic acid • Mercuric chloride solution• Benzoic acid • Concentrated sulphuric acid• Sodium metal • Ammonium hydroxide solution• Sodium carbonate • Neutral ferric chloride solution• Sodium hydroxide solution • Ammoniacal silvernitrate solution (Tollens’

reagent)

Instructions :• Do the following experiments using methanoic acid, ethanoic acid and benzoic

acid. If the acid used is a liquid, take about 1cm3, but if the acid is a solid useabout 0.5 g. Tabulate the observations.1. Mix with 1cm3of water and see whether it dissolves.2. Mix well with 2 cm3 of dilute sodium hydroxide solution. See whether the

acids dissolve.3. Put a small piece of sodium and test the gas liberated with a burning splinter

(It is essential that the acid is anhydrous).4. Add to a solution of sodium carbonate and see whether gas bubbles are

evolved.5. Take about 1cm3of ethanol, add carboxylic acid (or its sodium salt) and

then a few drops of conc.sulphuric acid and heat gently for about oneminute. Pour this mixture into cold water in a beaker and observe whethera pleasant smell is given.

6. Add about 3 cm3 of water to 0.5 g of mercuric chloride and prepare asolution. Heat the acid with a few drops of this mixture. Which of the acidsgive a white precipitate? Add a few drops of hydrochloric acid and seewhether the precipitate dissolves.

7. Add acidified potassium permanganate solution. Mix well and observewhether a colour change occurs.

8. Add about 2 cm3 of ammoniacal silver nitrate (Tollens’ reagent) and heat ina water bath. Observe whether a silver mirror is formed.

9. Add ammonia solution until it is just alkaline. Test the basic nature with ared litmus paper. Heat the solution until the smell of ammonia disappears.Add several drops of neutral ferric chloride to the solutions so neutralized.Observe the colour change.

29

Practical No. 22

Experimental determination of the effect of acid concentration on thereaction between magnesium and acids

Expected learning outcomes :1. Collects experimental data to identify the variation pattern of the reaction rate with

respect to the concentration of a given reactant.2. Determines the order with respect to hydrogen ion concentration in the reaction

between mangesium and acids.

Materials and Equipments : Pieces of clean magnesium ribbon about 3 cm long A boiling tube A rubber bung with a hole fitting the boiling tube About 400 cm3 of 1.0 mol dm-3 (approx.) hydrochloric acid A glass tube about 5 cm long fitting the hole of the bung A stop watch A beaker

Instructions : Using a piece of string or a rubber band mark a level on the boiling tube about 2 cm

from the bottom.

Boilingtube

Magnesium ribbon

Rubber bung

Glass tube

Magnesium ribbonAcid solution.

Mark (rubber band)HydrogengasMark

Figure 22 (a)

30

Fix the glass tube to the bung so that one end of the glass tube and the bottom surfaceof the bung are at the same level (The glass tube should be tightly fixed to the bung).Insert 0.5 cm of one of the magnesium ribbons into the space between the tube and thebung (In all experiments the length of the magnesium ribbon jutting out of the bungshould be equal).

Using the 1.0 mol dm-3 hydrochloric acid solution, prepare 100 cm3 of 0.8 mol dm-30.6 mol dm-3 , 0.4 mol dm-3 , 0.2 mol dm-3 hydrochloric acid solutions each (0.8 moldm-3 hydrochloric acid solution can be prepared by adding water to 80 cm3 of 1.0 moldm-3 hydrochloric acid solution until the total volume is 100 cm3). Fill the boiling tubecompletely with 1.0 mol dm-3hydrochloric acid solution and fix the rubber bung (withthe piece of magnesium ribbon) quickly. At the same time start the stop watch andinvert the boiling tube. Measure the time taken to reach the level of the solution to themark ( Mark on the boiling tube can be adjusted, so that the time is about 8-10seconds for 1.0 mol dm-3 solution).

Repeat the experiment using new magnesium ribbons and acid solutions of differentconcentrations (In every case use the same boiling tube and the same mark).

Using the time taken to produce a constant volume of gas, with acid solutions ofdifferent concentrations, determine the order of the reaction with respect to the acid.

31

Practical No. 23

Experimental determination of the effect of concentration of reactantson the reaction between sodium thiosulphate and nitric acid

Expected learning outcomes :1. Collects experimental data to identify the variation pattern of the reaction rate with

respect to each reactant.2. Determines the order of the reaction with respect to each reactant.

Materials and Equipments : A sodium thiosulphate solution (40 g dm-3 approx.) 3.0 mol dm-3 (approx.) nitric acid A 50 cm3 beaker A stop watch Boiling tubes Measuring cylinders

Instructions : Draw a cross on a white paper and place the beaker containing the measured

volume of the thiosulphate solution on it. Mix water and acid in a boiling tube. Addthis mixture in the boiling tube to the solution in the beaker while starting the stopwatch keeping the eye at a constant height above the beaker. Measure the time takenfor the disappearance of the cross.

(a) Determining the relationship between the rate of the reaction and thethiosulphate ion concentration Mix the following solutions as described above and measure the time taken for the

cross to disappear.

Volume of thiosulphate Volume of Volume of Time (t)/ssolution / cm3 acid / cm3 water / cm3

25.0 5.0 -20.0 5.0 5.015.0 5.0 10.010.0 5.0 15.0 5.0 5.0 20.0

Pay attention to the following points when you take readings.

Table 23 (a)

32

Do not shake the beaker; allow sulphur to settle down freely. Observation should be made at the same height by the same student using the same

beaker and the cross. After every experiment wash the beaker well to remove all the sulphur deposited on

the bottom. Wipe well the under surface of the beaker every time you place it on thecross.

(b) Determining the relationship between the rate of the reaction and the hydrogenion concentration Mix the following solutions as described below.

Table 23 (b)

Volume of thiosulphate Volume of Volume of Time (t) / ssolution/cm3 acid /cm3 water/cm3

25.0 5.0 -25.0 4.0 1.025.0 3.0 2.025.0 2.0 3.025.0 1.0 4.0

Obtain readings as described above.

33

Practical No. 24

Experimental determination of the effect of concentration of reactantson the reaction between iron(III) ions and potassium iodide

Expected learning outcomes :1. States that the rate of the above reaction changes with the change in concentration of

iron(III) ions.2. Determines the order of the reaction with respect to iron(III) ions.

Materials and Equipments : 10 cm3 and 25 cm3 measuring cylinders 0.1 mol dm-3( approx.) acidified ammoniumiron(III) sulphate solution 0.1 mol dm-3( approx.) potassium iodide solution 0.0006 mol dm-3( approx.) sodium thiosulphate solution (containing starch) Small beakers A stop watch

Preparation of solutions : Ammoniumiron(III) sulphate solution

Add 50 cm3 of 1.5 mol dm-3 sulphuric acid and some water to 12.0 g ofammoniumiron(III) sulphate and heat till the salt dissolves. Dilute the solution to250 cm3 with water.

Potassium iodide solutionDissolve about 4.00 g of potassium iodide in water and dilute the solution to250 cm3.

Sodium thiosulphate solution (with starch)Dissolve about 0.25 g of sodium thiosulphate in some water. Add 2.00 g of starch toabout 50 cm3 of water and heat till it dissolves. Mix the two solutions and dilute to250 cm3 with water.

Instructions : As the following table indicates, measure iron(III) ion solution to one beaker

(beaker A) and relevant volumes of other solutions to another beaker (beaker B).

Experiment Volume of Volume of Volume of Volume ofwater/cm3 acidified Fe3+ KI Na2S2O3 solution

No. solution/cm3 solution/cm3 with starch /cm3

1 - 25.0 10.0 15.02 5.0 20.0 10.0 15.03 10.0 15.0 10.0 15.04 15.0 10.0 10.0 15.05 20.0 5.0 10.0 15.0

Table 24 (a)

34

Figure 25 (a)

Add the solution in A to that in B while starting the stop watch. Observe the colourof the solution well. As soon as the blue colour appears stop the stop watch andmeasure the time. Repeat same for the other combinations of solutions and measurethe time taken for the appearance of the blue colour.

Using the data you obtained, determine the order of the reaction with respect toiron(III) ions through relevant calculations.

Practical No. 25

Experimental study of the characteristics of a dynamicequilibriumusing Fe3+ / SCN- system

Expected learning outcomes :1. Understands that all the reactants and products are present in a chemical system in

equilibrium.

Materials and Equipments : 10 cm3 of 0.050 mol dm-3 (approx.) ferric chloride or ferric nitrate solution 100 cm3 of 0.20 mol dm-3 (approx.) ammonium thiocyanate or potassium thiocyanate

solution. Sodium hydroxide or disodiumhydrogen phosphate solution Test tubes

Instructions : Mix 5 cm3 of the ferric ion solution with 5 cm3 of thiocyanate ion solution. What colour

is the mixture? Dilute the above solution about five times with water ( until the solution becomes ‘ light

tea’ colour). Transfer about 5 cm3 portions of this diluted solution to four test tubes. Keep one as the

control. To one test tube add a few drops of conc. ferric chloride solution or a crystalof ferric alum. To another test tube add a few drops of conc. ammonium thiocyanatesolution or a crystal of potassium thiocyanate. To the other test tube add a few drops ofdilute sodium hydroxide solution or disodiumhydrogen phosphate solution. Compare thecolour of each solution with that of the control.

Iron (III) ionsolution

KI + Na2S2O3+ Starch + Water

BeakerA Beaker B

35

With the above observations explain that all the reactants and products are present in achemical system in equibilirium.

Practical No. 26

Experimental study of the effect of temperature on the equilibriumsystem of NO2 and N2O4

Expected learning outcomes :1. Understands that the equilibrium point of an equilibrium system can be changed by

changing the temperature.

Materials and Equipments : Copper turnings Concentrated nitric acid (Avoid contact with the skin) Four boiling tubes Suitable rubber bungs Ice chips Hot water 3 beakers A delivery tube with a stopper

Instructions : Take about 5 g of copper turnings to a boiling tube, add about 3 cm3 of conc. nitric

acid and fix the delivery tube. Fill three dry identical boiling tubes with the gas evolvedand stopper them (The intensity of the colour of the gas in the boiling tubes should bemore or less equal. This gas should not be inhaled).

Keep one boiling tube in water at room temperature as the control, the second tubein ice and the other in a hot water bath at about 70-80 0C. After some time comparethe intensities of colour with that of the control tube. Then interchange the tubes keptin ice and hot water. After some time, compare the colour intensities with previouscolour intensities.

Observe the changes in the intensities of colour in each tube. What is the reason forthe colour changes observed?

36

Practical No. 27

Experimental determination of the solubility product of calciumhydroxide

Expected learning outcomes :1. Acquires the ability to determine the hydroxide ion concentration of a saturated aqueous

solution of calcium hydroxide.2. Demonstrates experimentally that the ionic product of various aqueous solutions saturated

with calcium hydroxide is a constant.

Materials and Equipments : Solid calcium hydroxide 0.10 mol dm-3 (approx.) standard sodium hydroxide solution 0.10 mol dm-3 (approx.) standard hydrochloric acid solution A burette A pipette Phenolphthalein Conical flasks

Instructions : Number five conical flasks as 1, 2, 3.... and prepare five solutions by adding different

volumes of sodium hydroxide solution and water into them according to the tablegiven below.

Flask Volume of 0.10 mol dm-3 Water / cm3

sodium hydroxide solution /cm3

1 - 100.0 2 25.0 75.0 3 50.0 50.0 4 75.0 25.0 5 100.0 -

To prepare systems 2, 3 and 4 use a burette or a pipette. But to prepare systems 1and 5 a measuring cylinder can be used (Discuss the reason for this).

Add excess of solid calcium hydroxide (about 1 g) to each of the above flasks, stirwell and allow to stand for about 15 minutes. Filter the solution mixture in flask 1 toa dry beaker.

Using a pipette, measure two portions of 25.0 cm3 of the filtrate and add to a conicalflask. Add 1-2 drops of phenolphthalein and titrate with the standard hydrochloricacid solution .Filter other four systems as well. Titrate the filtrate with the acid asbefore and record the results.

Using the above data find the ionic product [Ca2+(aq)][OH-(aq)]2 in the systems 1-5.Then, determine the solubility product of calcium hydroxide.

Table 27 (a)

37

Practical No. 28

Experimental determination of the distribution coefficient for thedistribution of ethanoic acid in water and butanol

Expected learning outcomes :1. Acquires the skill of determining the concentration of ethanoic acid in butanol and

water layers when ethanoic acid is in equilibrium between butanol and water.2. Using experimental data, determines the ratio between the concentrations of ethanoic

acid in water and butanol at equilibrium.

Materials and Equipments : About 120 cm3 of butanol 250 cm3 of approximately 1 mol dm-3 ethanoic acid solution 250 cm3 of approximately 0.5 mol dm-3 sodium hydroxide solution Five reagent bottles Burettes Pipettes Titration flasks Funnels Phenolphthalein

Instructions : 1 mol dm-3 ethanoic acid solution can be prepared by diluting 15 cm3of glacial

acetic acid ( 99% w/w) to 300 cm3 with water. As indicated in the following table,mix the respective volumes of liquids / solutions separately in the five reagent bottleslabeled 1-5. Use burettes to measure the volumes of liquids/ solutions.

System Butanol/cm3 1 mol dm-3 ethanoic Water/cm3

acid/cm3

1 20.0 40.0 - 2 20.0 35.0 5.0 3 20.0 30.0 10.0 4 20.0 25.0 15.0 5 20.0 20.0 20.0

Allow the systems to stand for 10-15 minutes to reach the equilibrium. Put thesystem in the first bottle to the burette (The upper layer is butanol). After theseparation of the layers, run 10.00 cm3 of the aqueous layer to a titration flask andadd 1 – 2 drops of phenolphthalein. Titrate this solution with the prepared sodiumhydroxide solution and take readings.

Table 28 (a)

38

Remove the remaining portion of the aqueous layer left in the burette carefully and take10.00 cm3 of the butanol layer to a titration flask. Add about 10 cm3 of water and 1-2drops of phenolphthalein to the flask and titrate with the sodium hydroxide solution . Takethe readings. Using the readings complete the following table.

Using experimental data, suggest a value for the distribution coefficient for the distributionof ethonoic acid between water and butanol.

System

12345

Volume of NaOHsolution requiredfor 10.00 cm3ofthe aqueouslayer /cm3

Volume ofNaOHsolutionrequired for10.00 cm3ofbutanol layer/cm3

[CH3COOH]aq/mol dm-3

[CH3COOH]butanol/mol dm-3

3

3

butanol[CH COOH]

[CH COOH]aq

Table 28 (b)

39

Practical No. 29

Preparation of an indicator using flowers provided and experimentaldetermination of its pH range

Expected learning outcomes :1. Uses that extracts of plant materials as indicators.2. Determines the pH range of a prepared indicators.

Materials and Equipments : 1.00 mol dm-3hydrochloric acid solution, 1.00 mol dm-3 sodium hydroxide solution . pH indicators available in the laboratory (phenolphthalein, methyl orange,

litmus, universal indicator). Plant materials (e.g. Argyreia flowers promegnate flowers, Clitoria flowers). Measuring cylinder (10 cm3) or a burette Test tubes Droppers Mortar and the pestle Funnels Filter papers Two rubber bands

Instructions :(a) Preparation of plant extracts Using the mortar and the pestle, crush each type of the plant material with some

water. Filter them and prepare the extracts. Blue flowers are more suitable for this.

(b) Preparation of solutions of pH 1 – 13 pH of 1.0 mol dm-3 hydrochloric acid solution is approximately 0. Take a test tube and add 1 cm3 of water to it from a burette or a measuring cylinder.

Mark the level. Remove water, add 10 cm3 of water and mark the level. This canbe used to measure 1 cm3 and 10 cm3 approximately. Thus, this can be used todilute a volume ten times.

Add 1.00 mol dm-3 HCl solution to the 1 cm3 mark and then water to the 10 cm3

mark. Shake well . The pH of the resulting solution is approximately 1. Repeat this method to prepare solutions of pH 2, 3, 4, 5 and 6 by successive

dilution. Distilled water can be used as a solution of pH 7. pH of 1.0 mol dm-3 sodium hydroxide solution is approximately 14. By successive

dilution of this solution as described above, prepare solutions of pH 13, 12, 11, 10,9 and 8.

Thus you have a set of solutions of pH 0- 14.

40

(c) Determination of pH ranges of indicators Take 15 test tubes and label them from 0-14. Add about 5 cm3 of the solutions you

have prepared into them. Add two drops of a laboratory indicator (phenolphthalein) into them and observe

the colour preferably on a white background. Note the pH values of the test tubesbeyond which a clear difference in colour can be observed. Repeat same with otherlaboratory indicators (methylorange, litmus, universal indicator).

Repeat same with the plant extracts. Sometimes more drops (5 – 10) may berequired to see the colour clearly.

Determine the pH range of the indicators by your observations.

Practical No. 30

Experimental determination of the acidic/basic/neutral nature ofaqueous solutions of salts by testing pH.

Expected learning outcomes :1. Shows that solutions of all salts are not neutral.2. By experimental observation, determines the acidic, basic or neutral nature of a salt

solution.

Materials and Equipments : Sodium chloride Sodium acetate Ammonium acetate Zinc chloride Aluminum chloride Magnesium sulphate Universal indicator or pH papers Test tubes A 10 cm3 measuring cylinder

Instructions : Take about 0.5 g of each salt .Put them separately to test tubes, add about 10 cm3

of water and prepare solutions by dissolving the salt. Take about 10 cm3 of waterused to prepare the solutions into another test tube.

To every test tube, either add two drops of the universal indicator or put a piece ofpH paper. Compare the colour of each solution with that of water and decide theapproximate pH of the respective salt solutions. Using your results complete thefollowing table.

41

Salt solution NaCl CH3CO2Na CH3CO2NH4 ZnCl2 AlCl3 MgSO4 Water

Approximate pH

Note : Highlight the relationship between the nature of the salt and the acidic/basic/neutral

nature of its solution.

Practical No. 31

Determination of approximate pH of a given solution using pHindicators

Expected learning outcomes :1. Determines the approximate pH value of a given solution using pH indicators.

Materials and Equipments : Indicators having different pH ranges Several different solutions to determine approximate pH value pH papers /pH meter Droppers Test tubes

Instructions : To several test tubes take 3 cm3 portions of each of the solutions whose pH has to be

determined. Add about 2-3 drops of pH indicators into each test tube (To one tube add only one

indicator). Using the data given in the following table determine the approximate pH value. Confirm the conclusion using pH papers or a pH meter.

Indicator pH range Lower colour Upper colour0 - 1.61.1 - 2.93.1 - 4.44.2 - 6.45.0 - 8.06.6 - 8.48.3 - 10.010.1– 12.0

Methyl violetThymol blueMethyl orangeMethyl redLitmusPhenol redPhenolphthaleinAlizarin yellow

YellowRedPinkRedRedYellowColourlessYellow

VioletYellowYellowYellowBlueRedRedOrange

Table 30 (a)

Table 31 (a)

42

pH range

Violet

Yellow

Yellow

Yellow

Blue

Red

Red

Orange-red

Red

Red (pink)

Red

Red

Yellow

Colourless

Yellow

Methyl violet

Thymol blue

Methyl orange

Methyl red

Litmus

Phenol red

Phenolphthalein

Alizarin yellow

Figure 31 (a)

4

4

414

43

Practical No. 32

Measuring the conductivity of a given solution using the Wheatstonebridge principle

Expected learning outcomes :1. Uses a conductivity cell to measure the electrical conductivity of solutions.2. Uses a Wheatstone bridge circuit to measure the conductivity.3. Compares the conductivities of various solutions .4. Determines approximately, how the conductivity of a solution varies with its

concentration.

Materials and Equipments : Meter bridge 4.5 V battery ( 3 x 1.5 V cells) Oscillator Earphone or a speaker

Instructions : Make an immersion type conductivity cell as indicated in the following figures.

Figure 32 (a)

Set the meter bridge as follows using an oscillator or a signal generator (available inthe physics laboratory) as the electrical source.

1 cm

2 cm

1cm

1 cm

Lipstickcovering(insulator)

Narrowcarbon rods

Glass rod

5-10 mm

Rubberbung

Cut this layer

Rubberlayer

Punctured layer

44

Figure 32 (b)

Balance the bridge using a 0.010 mol dm-3 KCl solution. Calculate cell constant of the celltaking the conductivity of the solution at room temperature as 1500S cm-3 Using differentsolutions and a sample of water, calculate their conductivities. Adjust the resistance R2sothat the balance points obtained somewher in the middle of the slide wire of the meterbridge.

Note : Prepare five KCl solutions of concentrations 0.0010 mol dm-3 to 0.050 mol dm-3. Determine

their conductivities and plot a graph to show the variation of the conductivity withconcentration.

Discuss the fact that by the conductivities of various water samples, one can get an idea oftheir salinity.

Practical No. 33

Experimental determination of the variation of conductivity withconcentration, temperature and the nature of the electrolyte

Of the factors concentration, temperature and the electrolyte, changing one at a time examinethe effect of the factor concerned. Follow the procedure described in Experiment No. 32to measure the conductivity.

Osillator ( 1000 Hz )

R2(500 )

Covered part

45

Practical No. 34

Making different types of electrodes and measuring their electrodepotentials using the potentiometer

Expected learning outcomes:1. Prepares different electrodes.2. Measures using the potentiometer, the e.m.f of the cell made by connecting an electrode

with the standard silver- silver chloride electrode.3. Determines the electrode potentials of the electrodes prepared using silver - silver

chloride standard electrode and Ecell = E - E Rf.

Materials and Equipments : A silver wire 6 cm long Carbon electrode Dilute HCl solution or a NaCl solution 3V battery Porous partition Potentiometer Moving coil voltmeter Digital voltmeter

Instructions : Using a 6 cm long silver wire as the anode and a carbon electrode as the cathode,

electrolyse a dilute hydrochloric or a sodium chloride solution. A potential differenceof 3V is sufficient. Grayish white AgCl is deposited at the end of the silver wire. Bydipping this in a Cl- ion (KCl) solution, Ag(s)/AgCl(s)/Cl-(aq) electrode is obtained.

Connect the electrodes you have constructed and measure the e.m.f. of cells using thepotentiometer. Also measure their e.m.f. by an ordinary moving coil voltmeter and amodern digital voltmeter. Confirm your results by measuring the e.m.f. of a new 1.5 Vcell available in the market.

Calibrate the potentiometer using a standard Weston cadmium cell. If it is not available,use either a Daniel cell or a normal new 1.5 V battery.

Figure 34 (a)

2 2

1 1

E lE l

46

Connect the cell terminals correctly as shown in the diagram. Confirm the polarity ofthe cell using a common voltmeter. If not, a balance point cannot be obtained. Measurethe e.m.f. of the cells you made with the potentiometer and record.

Practical No. 35

Measurement of the electromotive force of various cells using thepotentiometer

Instructions : Measure the e.m.f of various cells available in the laboratory using the method followed

in Experiment No. 34.

Practical No. 36

Comparison of efficiency of cells

Expected learning outcomes :1. Demonstrates that the condition of a cell does not depend only on its size and

chemistry.2. Compares the efficiencies of cells of same size and chemistry available in the

market.3. Selects suitable cells for day to day usage according to needs.

Materials and Equipments : Three cells of type R 20 but of different colours (e.g. black, red, blue) Voltmeter (1-2 V or 1-5 V) Three 5 resistors Connecting wires

Instructions : Construct the circuit as indicated in the Figure 36 (a).

Figure 36 (a)

V

1.5 V

5

47

Get the voltmeters reading every 15 minutes after connecting the circuit. It is adequateto connect the voltmeter only when the measurement is taken. Complete theexperiment within one hour (to get four readings) and disconnect the circuit. Nextday repeat the experiment and obtain four readings. Likewise repeat the experimentseveral days (six days would be sufficient) within one hour each day. Repeat theexperiment for different cells. Plot the readings obtained as follows.

Voltage between the terminals /V

Date

Figure 36 (b)

Note : Different curves would be obtained for different cells of same size and voltage.

Discuss how the shapes of these curves vary with the efficiency of the cells. Compare the results obtained for cells which differ (a) in size (b) chemically. Discuss how these curves are important in selecting cells for a particular task. Explain how the causes for these differences change with the construction of the

cell.

48

Practical No. 37

Experimental determination of the decomposition potential of a simpleelectrolytes

Expected learning outcomes :1. Obtains experimental data for drawing I – E curve for simple electrolytes.2. Draws I -E curves for simple electrolytes.3. Compares the decomposition potential of different electrolytes.

Materials and Equipments : Potentiometer (or a reversible resistance) 4.5 V battery Voltmeter Ammeter Carbon electrodes Copper electrodes Dilute sulphuric acid solution Dilute copper sulphate solution Dilute Sodium hydroxide solution

Figure 37 (a)Instructions : Construct the circuit given in Figure 37 (a). Adjusting the potentiometer or the

reversible resistance, observe how the current flowing through the solution changeswith the applied potential difference. Plot the current against the potential differenceand obtain the decomposition potential. Cut off electricity and see whether there isany potential difference between the electrodes.

Also obtain I –E curves for the electrolysis of(a) Copper sulphate solution using carbon electrodes.(b) Copper sulphate solution using copper electrodes.(c) Sodium hydroxide solution using carbon electrodes.

Note : Discuss the fact that for the conduction of an electric current through a solution,

electrode reactions should occur at the electrodes. Explain why there is a potential difference in the system even after breaking the circuit

after electrolyzing the dilute sulphuric acid solution. Present the concepts of polarization and over voltage. Discuss observations during the electrolysis of copper sulphate solution.

AV

Dilute sulphuricacid

i

49

Practical No. 38

Preparation of hydrogen and oxygen gases using electrolysis of water.

Expected learning outcomes :1. Names electrolysis as an energy storing process.2. Shows that hydrogen and oxygen are the products of hydrolysis of water.3. Compares the volume of the gases liberated in hydrolysis of water.

Materials and Equipments : Small plastic basin Two test tubes Two narrow carbon electrodes Two rubber stoppers fitting the test tubes Acidulated water 4.5 V battery Paraffin wax Cork borer

Instructions :

Figure 38 (a) Set the circuit as shown in the figure. Initially the two test tubes should be filled with acidulated water. Fix the stoppers loosely to the test tubes so that water can come out. Provide an electric current and observe the collection of gases in the test tubes.

Acidulated water

50

Practical No. 39

Investigation into the corrosion of iron as an electrochemical process

Expected learning outcomes :1. Investigates the effect of magnesium, zinc, copper and lead on corrosion of iron in a

neutral gel medium containing sodium chloride.

Materials and Equipments : Petri dishes (5) Iron nails about 5 cm long Strips of magnesium, zinc, copper and lead metals about 4 cm long Sand papers Hydrochloric acid Agar agar Sodium chloride Potassium ferricyanide Phenolphthalein

Instructions : clean the iron nails, first heat them with conc. hydrochloric acid, wash with water

and dry. Then clean them further with sand paper. Clean the metal stripes also withsand paper. Around the middle of each nail wind the metal strips tightly.

Weigh 7 g of sodium chloride and 5 g of agar agar approximately. Heat about 250cm3 of water in a beaker. Then add sodium chloride and agar to hot water and stirwell with a glass rod till they dissolve. The solution should be boiled at least for 10minutes. When the solution becomes viscous, add phenolphthalein, remove the beakerfrom the flame and allow to cool while stirring. When the solution has cooled a little,add about 1 cm3 of the potassium ferricyanide solution to the mixture. Stir the mixturewell to expel the air bubbles as much as possible. Put this viscous agar mixture intofive petri dishes.

Dip horizontally, the four metal pairs in semi solid agar in four petri dishes. In theother petri dish dip the iron nail.

After some time, agar solidifies and the colour of agar in the areas around the nailsbegins to change. Keep the dishes for about one day, observe changes and record.

N.B. When boiling, the solution should not overflow. Avoid contact of the potassiumferricyanide solution with the body.

51

Practical No. 40

Electroplating with nickel and copper

Expected learning outcomes :1. Electroplates metals using simple practical electroplating systems.

Materials and Equipments : A beaker Battery Ammeter Variable resistance

Electroplating of copperCuSO4.5H2O (200-250 g), conc. sulphuric acid (15-25) cm3, very small amount ofgelatin dissolved in 1 cm3of water.

Suitable temperature: 20 - 40 0C,Current density: 20 - 50 mA cm-2

Anode: Copper metal

Electroplating of nickel250 g of NiSO4

.7H2O, 45 g of NiCl2, 30 g of boric acid (H3BO3) and a few of saccharinedissolved in 1 cm3 of water.

Suitable temperature: 40 - 70 0CCurrent density: 20 - 50 mA cm-2

Anode: Nickel metal

Instructions :

Figure 40 (a)

Construct the circuit as shown in the diagram. Take the electrolyte into the beaker and connect the object to be electroplated as the

cathode. Electrolyse using a suitable current under suitable temperature.

52

Practical No. 41

Tests for selected cations

Expected learning outcomes :1. Conducts tests to identify the cations Mg2+(aq), Ca2+(aq), Ba2+(aq), Al3+(aq),

Pb2+(aq), Fe2+(aq), Fe3+(aq), Ni2+(aq), Cu2+(aq) and Zn 2+(aq) in aqueousmedium.

2. States that the comparison of solubility of compounds can be used to identifycations.

3. States that some cations form soluble complexes.4. Observes the colours and name the cations of the d- block.

Materials and Equipments : Approximately 0.1 mol dm-3 solutions of Mg 2+(aq), Ca2+(aq), Ba2+(aq), Al3+(aq),

Pb2+(aq), Fe2+(aq), Fe3+(aq), Ni2+(aq), Cu2+(aq), Zn2+(aq) (prepared bydissolving soluble salts of corresponding metals in distilled water)

H2SO4 solution (2 mol dm-3 approx.) NH4OH solution (4 mol dm-3 approx.) NaOH solution (4 mol dm-3 approx.) Test tubes

N.B. When preparing solution of iron(II) and iron(III), they should be made acidicinitially.

Instructions :Test ITake about 1 cm 3 of solutions of Mg 2+ (aq), Ca 2+(aq), and Ba 2+(aq) separately to

test tubes. Add about 1 cm 3 of 2 mol dm -3 of Sulphuric acid solution to each ofthem and mix. Record observations.

Test IIa) Take about 1 cm 3 of Al 3+(aq) and Pb 2+(aq), solutions separately to test tubes .Add

about 1cm 3 of 2 mol dm-3 H2SO4 solution to each of them, mix and recordobservations.

b) Take about 1 cm 3 of Al 3+(aq) and Pb 2+(aq) solutions separately to test tubes andadd 4 mol dm-3 NaOH solution dropwise to each of them . Add reagent in excess ifprecipitation occurs. Record observations.

Test IIIa) Take about 1 cm 3 of Fe 2+(aq), Fe 3+(aq), Ni 2+(aq), Cu 2+(aq) and Zn2+(aq)

solutions separately to test tubes. Add 4 mol dm-3 NaOHsolution dropwise toeach of them and mix. Add reagent in excess if precipitation occurs. Recordobservations.Heat and record observations.

53

b) Take about 1 cm3 of Fe2+(aq), Fe3+(aq) ,Ni2+(aq), Cu 2+(aq) and Zn2+(aq) solutionsseparately to test tubes. Add 4 mol dm-3NH4OH solution dropwise to each of themand mix. Add the reagent in excess if precipitation occurs. Record observations.

Note : In test I, Mg 2+(aq) does not give a precipitate. Ca 2+(aq) gives a thin precipitate or

turbidity whereas Ba 2+(aq) gives a thick white precipitate. (Ksp(CaSO4) is 9.0 x 10-4

mol2 dm-6 and Ksp(BaSO4) is 1.0 x 10-10 mol2 dm-6. Using these values explain theabove observations).

In test II (a), Al 3+(aq) does not give a precipitate but Pb 2+(aq) gives a whiteprecipitate (Ksp(PbSO4) is 2.0 x 10-8 mol2 dm-6 ).

In test II (b), both Al 3+(aq) and Pb 2+(aq) give white precipitates. The precipitategiven by Al 3+(aq) dissolves in excess NaOH giving a clear solution.Write balancedequations for the above observations.

In test III (a) Fe 3+(aq), Fe 2+(aq), Cu2+(aq) and Ni 2+(aq) give precipitates, that areinsoluble in excess NaOH.Indicate their colours and write balanced equations toexplain the relevant reactions. On heating, colours of Fe(OH)2 and Cu(OH)2 change.Write equations to explain these colour changes.

The white precipitate given by Zn 2+dissolves in excess NaOH giving a clear solution.Write a balanced equation to explain this observation.

In test III (b) all solution give precipitates. On adding excess ammonia solutionfollowing observations can be made.

Fe (OH)2 (s) Does not dissolveFe (OH)3 (s) Does not dissolveCu (OH)2 (s) Dissolves giving a deep blue solutionNi (OH)2 (s) Dissolves giving a blue solutionZn (OH)2 (s) Dissolves giving a colourless solution

Write balanced equations to explain the relevant reactions.

54

Practical No. 42

Tests for selected anions

Expected learning outcomes :1. Identifies the reactions of the halides ions in aqueous solution.2. Describes the action of dilute nitric acid on silver halides.3. Examines the solubility of silver halides in dilute ammonia.4. Detects SO4

2- , SO32-

, and S2- ions in aqueous solution.5. Detects NO3

- and NO2- ions in aqueous solution.

6. Detects phosphate ions by precipitation.7. Identifies carbon dioxide evolved.8. Develops the skill of using the delivery tube.

Identification of halide ions

Materials and Equipments : 0.10 mol dm-3 of solutions the following ions.

Cl-(aq), Br-(aq), I-(aq)(Prepared by dissolving corresponding sodium, potassium or ammonium salts)

0.10 mol dm-3 silver nitrate solution Dilute nitric acid 1.0 mol dm-3 ammonia solution Test tubes

Instructions : Take about 1 cm-3 of each halide solution to test tubes separately.

Treat them as follows and record observation. Add few drops of dilute nitric acid to the solution and sbout 1cm3 of silver nitrate

solution. Add ammonia solution dropwise to the precipitates formed.

Note : Ksp(AgCl) = 1.0x10-10 mol2 dm-6

Ksp(AgBr) = 7.7x10-13 mol2 dm-6

Ksp(AgI) = 8.3x10-17 mol2 dm-6

Ag+(aq) + 2NH3(aq) [Ag(NH3)2]+ (aq)K = 1.7x107 dm6 mol-2

Using the above data explain the solubility of AgCl(s), AgBr(s) and AgI(s) inammonia solution.

Silver halides (X= Cl , Br, I) are not soluble in dilute HNO3. Why?

55

Identification of SO42-, SO3

2- and S2- ions

Materials and Equipments : Na2SO3, Na2SO4, Na2S ( or potassium salts) Dilute sodium hydroxide solution Dilute hydrochloric acid Dilute nitric acid 0.01 mol dm-3 Ba2+(aq) solution , 0.01 mol dm-3 Ni2+(aq) solution , 0.01 mol dm-3

Cu2+(aq) solution Test tubes

Instructions :Using 1.0 cm3of each solution of ions, carry out the following tests and recordobservations.

Note : BaSO3 dissolves in dil. HNO3 , whereas BaSO4 is not soluble in dil. HNO3. Why?

Ksp (BaSO4 ) = 1.1 x10-10 mol2 dm-6

H2S(aq) H+(aq) + HS-(aq);1aK = 1.0 x10-7mol dm-3

HS- (aq) H+(aq) + S2-(aq) Ka2 = 1.0 x10-14 mol dm-3

In the light of above data, explain that Ni2+ precipitates in basic media but Cu2+

precipitates in both acidic and basic media as sulphides.

Solution Test Observation

SO42-(aq) (i) Add 1cm3 of Ba2+ solution. ..........................

Then add dil. HNO3 . ..........................

SO32-(aq) (i) Add 1cm3 of Ba2+ solution. ..........................

Then add dil. HNO3 . ..........................

S2-(aq) (i) Add dil. HCl solution. .......................... Then add Cu2+ solution. ..........................

(ii) Add dil. HCl solution ..........................

Table 42 (a)

56

Identification of NO2- and NO3

- ions in aqueous solution

Materials and Equipments : 0.1 mol dm-3 NO2

-(aq) and 0.1 mol dm-3 NO3-(aq) solutions (approx).

Concentrated sulphuric acidDilute nitric acid

0.1 mol dm-3 Fe2+(aq) solution Aluminium powder Concentrated sodium hydroxide solution Red litmus paper Concentrated hydrochloric acid Aniline or p-aminosulphonic acid Ice cubes Conical flask Thermometer

Instructions :1. NO2

-

To 1 cm3 of the NO2- solution add 1 cm3 of dil. HCl and cool . to about 0-50C.

Add aniline and then add NaOH to make the solution alkaline .Then add phenol.Observe (coupling). If p-amino sulphonic acid is used, cooling is unnecessary.

2. NO3-

Add 4 cm3 of freshly prepared iron(II) ion solution to 1cm3 of NO3-(aq) solution

and then add 3-4 cm3 of conc. H2SO4 slowly down the side of the test tube(Brown ring test).

3. NO2- and NO3

-

Take 1cm3 of NO2- or NO3

-or a mixture of NO2- and NO3

- solution. Add Alpowder / Devarda’s alloy / Zn dust and then conc.NaOH (aq) .Gently warm thesolution if necessary. Test the gas liberated with moist red litmus paper or a papersoaked in Nessler’s reagent.

Note : Discuss the relevent reactions.

Identification of CO32- ions

Materials and Equipments : A solid soluble carbonate (e.g. Na 2CO3) Lime water Dilute hydrochloric acid Aqueous solution of barium chloride or calcium chloride Delivery tube Test tubes

57

Instructions : Add dil. HCl to a solid carbonate ( 0.5 g) in a test tube. Pass the liberated gas

immediately into the lime water using the delivery tube. Continue passing the excess gasinto lime water.

Add BaCl2(aq) /CaCl2(aq) to the solution of carbonate ions. Then add dil. HNO3 to the product Record observations Write balanced equations for the above reactions and explain the observations.

Identification of PO43- ions

Materials and Equipments : PO4

3- or HPO42- or H2PO4

- solutions, 0.1mol dm-3

1 mol dm-3 barium chloride solution Dilute nitric acid

Instructions : Mix 1 cm3 of BaCl2(aq) to 1cm3of the phosphate ion solution. Then add dilute HNO3 to the above solution. Record your observation in each case.

Practical No. 43

Recrystallization using pure solvents/solvent mixtures

Expected learning outcomes :1. Develops the skills of purification using charcoal followed by recrystallization with pure

solvents/ solvent mixtures.

Materials and Equipments : Animal charcoal Alcohol Salicylic acid, 2 g (mixed with a very small quantity of a dye available in the market) Filter funnel Filter paper

Instructions : Mix the impure salicylic acid sample with 1g of animal charcoal. Dissolve it in a minimum

volume of 84% alcohol - water mixture (charcoal exists in the solid phase). Keep themixture in a water bath and heat to a temperature of 65-70 0C for 15 - 20 minutes.Filter while hot through a filter paper and allow to cool. Salicylic acid recrystallizes.

Compare the recreystallized sample with the initial crude sample.

58

Practical No. 44

Experimental determination of the moisture content of a sample offood

Expected learning outcomes :1. Applies gravimetric method successfully to determine the moisture content of a given

food sample.

Materials and Equipments : Sample of a pulse (e.g. dhal, gram, green gram, cow pea etc.) or a condiment such

as clove, pepper, turmeric etc. Mortar and pestle Petri dish or a crucible Balance (triple beam, chemical or electronic) Temperature adjustable oven Sand bath with fine sand (In case where an oven is not available) Thermometer Burner

Instructions : Add about 10 g of the food sample to the mortar and grind into a powder with the pestle. Weigh a clean and dry petri dish or crucible or evaporating dish. Add about 5.00 g –

6.00 g of the crushed food material into a petri dish and weigh as much accurately aspossible within the limits of the balance (By these readings ,the weight of the wet foodsample can be obtained).

Keep the petri dish in the oven and adjust the temperature to 105 0C. Let the food sample be in the oven at least for 30 minutes. Take the petri dish out, allow to cool to the room temperature in a desiccator and weighing. Keep the petri dish again in the oven for about another 10 minutes. Cool and weigh

again. If a further loss in weight is noted, heat again and repeat weighing until you get a constant

weight. From your readings find the weight of the food sample and the weight of the moisture in

it. Give the moisture content of the food sample as a percentage by weight.

(In case where an oven is not available) Heat the petri dish with the food material in a sand bath at 105 0C for 30 minutes and

weigh. Repeat heating and weighing until a constant weight is obtained.

59

Note :Try to answer the following questions. Why temperature should not be allowed to rise beyond 105 0C ? What is the reason for taking readings repeatedly until a constant weight is

obtained? Can the same method be used to find the content of water in a hydrated salt such as

CuSO4.5 H2O ? Comment.

Practical No. 45

Determination of the stoichiometry of the reaction between bariumchloride and sulphuric acid (Precipitation method)

Expected learning outcomes :1 Determines experimentally the molar ratio between the reactants using a measurable

change during the reaction.2. Develops the ability of drawing and interpreting graphs.3. Determines the unknown concentration of a sulphuric acid solution using precipitation

method.

Materials and Equipments : 0.5 mol dm-3 barium chloride solution 0.5 mol dm-3 sulphuric acid solution About 10 tests tubes of equal cross section Two burettes or graduated pipettes

xxxx

Food sample

Sand

Petri dish

Metal tray

Heat

Thermometer

Figure 44 (a)

60

Instructions : Number test tubes 1-10 as given in the following table, and measure and mix the

volumes of the respective solutions using burettes or graduated pipettes

Table 45 (a)

Shake the test tubes well and allow to stand vertically for about one day for theprecipitates to settle.

Repeat the experiment keeping the volume of sulphuric and solution constant andvarying the volume of barium chloride solution [(Table 45 (b)].

Table 45 (b)

Plot the height of the precipitate against i). the volume of the barium chloride solutionand ii). the volume of the sulphuric acid solution.Your may get a graph such as a one in Figure 45 (a).

Test tube no.

Volume of the 0.5mol dm-3 BaCl2 solution /cm3

Volume of the 0.5mol dm-3 H2SO4

solution /cm3

1 2 3 4 5 6 7 8 9 10

5 5 5 5 5 5 5 5 5 5

1 2 3 4 5 6 7 8 9 10

Test tube no.

Volume of the 0.5mol dm-3 BaCl2 solution /cm3

Volume of the 0.5mol dm-3 H2SO4

solution /cm3

1 2 3 4 5 6 7 8 9 10

1 2 3 4 5 6 7 8 9 10

5 5 5 5 5 5 5 5 5 5

Height of the precipitate /cm

Volume of 0.5 mol dm-3

BaCl2/ H2SO4 solution/cm3

0 1 2 3 4 5 6 7 8 9

Figure 45 (a)

61

Using this graph, the volumes that give the maximum amount of the precipitate canbe found.

Note : How do you select the test tube in which the maximum amount of the precipitate has

formed ? Assume that the amount of the precipitate is proportional to the height of theprecipitate.

Observe the instance where maximum amount of the precipitate has formed. From itget the molar ratio in which the maximum amounts of the reactants have reacted. Thisis the stoichiometric ratio for the reactants.

N.B.In determining the higher stoichiometric ratios (e.g. 3:1, 4:1, etc.). using the continuousvariation method described above, practical difficulties arise. Need of a large number oftest tubes, volumes of solutions exceeding the capacity of test tubes are such barriers.

As a solution for this, an alternative method can be used where the mole fractions of thetwo reacting solution are varied keeping the total volume constant. As an exampleconsider the determination of the stoichiometry of the reaction between lead nitrate andpotassium chloride [Table 45 (c)].

If the height of the precipitate is plotted against the volumes of the solutions mixedfollowing graph is obtained.

Test tube no.

Volume of the 1.0 moldm-3 Pb(NO3)2solution/cm3

Volume of 1.0 moldm-3 KCl solution/cm3

1 2 3 4 5 6 7 8 9 10 11

1 2 3 4 5 6 7 8 9 10 11

11 10 9 8 7` 6 5 4 3 2 1

Table 45 (c)

62

Figure 45 (b)The molar ratio of the reactants in the test tube which gives the maximum height of theprecipritate is the stoichiometry of the reaction.

Experiment II

Materials and Equipments :• 100 cm3 of 5 mol dm-3barium chloride solution• 100 cm3 of 5 mol dm-3 sulphuric acid solution• 50 cm3 of sulphuric acid solution of unknown concentration

(The concentration should be within the range of 1-5 mol dm-3)• Ruler ( marked in cm and mm)

Instructions :• Prepare 100 cm3 portions of 1 mol dm-3, 2 mol dm-3, 3 mol dm-3, 4 mol dm-3

sulphuric acid solutions by diluting the 5 mol dm-3 sulphuric acid solution.• Mix the solutions as given in Table 45 (d). Allow the precipitate to settle and measure

its heights.

1 2 3 4 5 6 7 8 9 10 11 12 Volume of 1.0 mol dm-3Pb(NO3)2 solution/ cm3

Volume of 1.0 mol dm-3KCl solution/ cm3

Height of the precipiate/cm

11 10 9 8 7 6 5 4 3 2 1 0

63

Table 45 (d)

• Plot the height of the precipitate against the concentration of sulphuric acid.

Height of the precipitate/cm

Figure 45 (c)

• As indicated in Figure 45 (c), obtain the unknown concentration using the point whichcorresponds to its height.

Test tube no.

Volume of 5 moldm-3 BaCl2solution /cm3

Volume ofH2SO4 solution /cm3

1 2 3 4 5 6

5 5 5 5 5 6

5 5 5 5 5 51 mol dm-3 2 mol dm-3 3 mol dm-3 4 mol dm-3 5 mol dm-3 unknown

concentration

Height in tube 6

H2SO4 concentration/ mol dm-3 1 2 3 4 5

64

Practical No. 46

Determination of the stoichiometry between salicylic acid andiron(III) ions (Colorimetry method)

Expected learning outcomes:1. States that the intensity of the colour of the complex formed with salicylic is

proportional to the concentration of iron (III).2. Deduces the stoichiometry of the reaction using the maximum intensity of the colour

of the complex.3. Determines the formula of the iron (III) salicylate complex.4. States that concentration of any coloured complex / compound in aqueous solution

can easily be found with colorimetric apparatus.

Materials and Equipments : Ammoniumferric sulphate [(NH4)2SO4

.Fe2(SO4)3.24H2O]

(250 cm3 of 0.001 mol dm-3 Fe3+ solution) 250 cm3 of 0.001 mol dm-3 salicylic acid solution Test tubes of equal cross section (9) Funnels Burettes Measuring cylinders Volumetric flasks

Instructions : Add volumes of the respective solution into the test tubes as indicated in tables

given below.

Table 46 (a)

Test tube No. 1 2 3 4 5 6 7 8 9

Volume of the 0.001 mol dm-3

Fe3+ solution /cm3 1.0 2.0 3.0 4.0 5.0 6.0 7.0 8.0 9.0

Volume of the 0.001 mol dm-3

salicylic acid solution /cm3 5.0 5.0 5.0 5.0 5.0 5.0 5.0 5.0 5.0

Volume of water /cm3 9.0 8.0 7.0 6.0 5.0 4.0 3.0 2.0 1.0

65

Table 46 (b)

Test tube no. 1 2 3 4 5 6 7 8 9

Volume of the 0.001 mol dm-3

Fe3+ solution /cm3 5.0 5.0 5.0 5.0 5.0 5.0 5.0 5.0 5.0

Volume of the 0.001 mol dm-3

salicylic acid solution /cm3 1.0 2.0 3.0 4.0 5.0 6.0 7.0 8.0 9.0

Volume of water /cm3 9.0 8.0 7.0 6.0 5.0 4.0 3.0 2.0 1.0

After mixing the solutions, observe the intensity of colour of the solutions in the test tubeswith the naked eye and note from which tube an equal intensity in colour is observable.

Note : It can be seen that the intensity of colour is equal from test tube 5 onwards. Calculate the

molar ratio of reactants in test tube 5. It gives the stochiometry of the reaction betweeniron(III) ion and salicylic acid. The formula of the iron(III) salicylate complex can be deducedas follows.

The test tube with maximum colour intensity is selected by observing with the naked eye. Ifcolorimetric equipments are used, more accurate readings can be obtained easily.

Preparation of the solutions required :1. Weigh 12.0 g of ammoniumferric sulphate (ferric alum) accurately and dissolve as much

as possible in 10 cm3 of dilute sulphuric acid. Then add 150 cm3 of the same acid to itand boil (use a beaker). Cool the solution and transfer it to a 250 cm3volumetric flask.Add some water to the beaker stir well and add same to the volumetric flask. Dilute to250 cm3 mark with water. The resulting solution is 0.01 mol dm-3 . Pipette out 25.0 cm3

of this solution, add to a 250 cm3 volumetric flask and dilute it by adding water to the250 cm3mark. The concentration of Fe+3 ions in this solution is 0.001 mol dm-3 .

2. Weigh 3.45 g of salicylic acid accurately and dissolve it well in about 50 cm3of methanol.Add about 150 cm3 of water and dilute to 250 cm3 with methanol. The resulting solutionis 0.1 mol dm-3. A 0.001 mol dm-3 solution can be obtained by diluting 1.0 cm3 of thissolution to 100.0 cm3 with water.

N.B.1. It is better that the iron (III) solution made using ferric alum is prepared a day before the

experiment. This is because the dissolution of acidic ferric alum is slow.2. Salicylic acid solution should be prepared just before the experiment. This is to avoid

precipitation of salicylic acid with the evaporation of methanol.

CO

OH

Fe+ Fe3+

OH

COOH

+ H+

o

2+

66

Practical No. 47

Experimental determination of the concentration of a thiosulphatesolution using potassium iodate

Expected learning outcomes:1. Understands the practical uses of redox reactions.2. Develops the skills in iodometric titrations.

Materials and Equipments : Potassium iodate Balance Potassium iodide Pipette Sodium thiosulphate Measuring cylinder, 100 cm3

Soluble starch Burette Distilled water Titration flasks Sodium carbonate or chloroform Volumetric flask

Instructions : Prepare sodium thioulphate solution (0.1 mol dm-3 approx.) with a small amount of

either sodium carbonate or chloroform added to it. Prepare 0.02 mol dm-3 potassium iodate solution.

[Weigh accurately about 1.10 g of pure dry potassium iodate. Dissolve it in distilledwater and dilute to 250.00 cm3 using a 250 cm3 volumetric flask].

Take 25.00 cm3 portion of the above potassium iodate solution into a titration flask,using a pipette.

Add 1g of potassium iodide and 20 cm3 of 1.0 mol dm-3 sulphuric acid and shake todissolve the solid.

Titrate the liberated iodine with the given sodium thiosulphate solution kept in theburette.

When the colour of the solution in the titration flask has become pale yellow, dilute thesolution up to about 200 cm3 with distilled water.

Add 2 cm3 of freshly prepared starch solution and continue the titration, until thecolour changes form blue to colourless.

Repeat the procedure with another similar portion of potassium iodate.

Note : Why is sulphuric acid added to the titration flask ? Why should the starch solution be prepared fresh? What is the purpose of using starch? Explain why starch is added at the instance where the solution becomes pale yellow,

but not at the very beginning of the titration. Why do you dilute the solution in the titration flask before adding starch? Give reasons for adding a small amount of Na2CO3 or CHCl3 when preparing the

thiosulphate solution.

67

Practical No. 48

Separation of components of chlorophyll using paper chromatography

Expected learning outcomes:1. States that it is possible to separate the components in a mixture using paper

chromatography.2. States that separation of components depends on the difference of partition

coefficients of compounds in stationary and mobile phases.

Materials and Equipments : Boiling tube Rubber bung (split in the middle) Filter papers or blotting papers A few green leaves Mortar and pestle Ethanol Capillary tube

Instructions :

Figure 48 (a)

Cut a strip of filter paper about 15 cm long and 1.5 cm wide. Make one end pointed. Grind a few leaves with ethanol and water. Then filter and obtain a concentrated filtrate. Place a drop of the filtrate to be analyzed on the strip of the filter paper using a

capillary tube at a distance of about 3 cm from the pointed end. Arrange the apparatus as shown in the diagram so that the tip of the paper dips in the

mixture of water and ethanol. When the advancing front of the solvent approaches the top of the strip, remove

and dry it in air.

Rubber bung

Boiling tube

Strip of filter paper

Chlorophyll extract

Ethnol + Water

68

Note : Which is the stationary phase? Which is the mobile phase? What is the relationship between the distance travelled by the pigments, their

respective partition coefficients and their respective Rf values. Being soaked, the solvent migrates up the strip by capillary action carrying the

constituents of the sample along with it at different speeds according to their respectivepartition coefficients.

Students may use mixtures of ethanol and water of different proportions and comparethe separations of different pigments.

Practical No. 49

Experimental determination of the oxygen percentage in air by volume

Expected learning outcomes :1. Acquires the understanding and skill of determining the percentage of oxygen by

volume in air.

Method I

Materials and Equipments : Boiling tube Test tube Rubber bung Delivery tube Concentrated ammonium chloride solution Copper gauze (8 cm x 3 cm) Ammonia solution Measuring cylinder, 10 cm3

Measuring cylinder, 100 cm3

Instructions : Delivery tube

Ammonia solutionTest tube (B)

Boiling tube (A)

Copper gauzeAmmonium chloridesolution

Figure 49 (a)

69

Add 10.0 cm3of concentrated ammonium chloride solution at room temperature into theboiling tube A. Fix the copper gauze in the middle of the boiling tube. Add about 15 cm3

of aqueous ammonia solution into the test tube B. Fix the bung with the delivery tubetightly to A and dip the free end in the ammonia solution. Tilt the apparatus from time totime so that the ammonium chloride solution wets the wire gauze and observe any colourchange.

Observe that the solution in A turns blue. When shaken for about 10 minutes the colourbegins to diminish. As this happens the flow of solution into A ceases. Remove B and letany liquid left in the delivery tube flow into A. Measure the total volume in A.

Fill the boiling tube with the copper gauze with water and fix the bung with the deliverytube to it. Fill the boiling tube and the delivery tube completely with water and measurethat volume of water also. Record your readings as follows.

Total volume of the boiling tube + delivery tube = V1 cm3

Volume of the ammonium chloride solution taken initially = 10.0 cm3

The total volume of the solution in A after the experiment = V2 cm3

N.B.The apparatus should be kept air tight during the experiment.

Note : Here all the oxygen in the entrapped volume of air should be removed. How is this

effected? This is effected by making oxygen to react with a metal. Copper is the metalused in this experiment. What is the reason for it?

If a reactive metal such as magnesium is used, it easily reacts with oxygen to formmagnesium oxide. If all the oxygen in the volume of air is to react with magnesium, theoxide film formed on the metal should be removed. Magnesium oxide is a basic oxide.To remove it, A should contain an acid. In place of an acid, why an acidic solution ofammonium chloride is used here? If a reactive metal such as magnesium is used, is tractswith the acid liberating hydrogen gas which is added to the volume of air.

What type of metals do not evolve hydrogen by reacting with acids? Copper which is below hydrogen in the electrochemical series can be used for this

purpose. Why copper is preferred ?

22 ( ) ( ) 2 ( )Cu s O g CuO s

When copper oxide dissolves in the solution Cu2+ ions come into the solution turning itblue. When oxygen in entrapped air reacts with copper, ammonia solution is drawn intoA from B. Then why does the intensity of blue colour in the solution in A increase?

223 3 4( ) 4 ( ) ( ) ( )Cu aq NH aq Cu NH aq

light blue deep blue

70

When all the oxygen in the volume of air is used up in the reaction , the intensity of bluecolour decreases again due to the following reaction.

23 4 3 2( ) ( ) ( ) 2 ( ) ( )Cu NH aq Cu s Cu NH aq

colourlessDiamminecopper(I) ion is formed here.

What can be expected if the colourless solution is exposed to air after the reaction ? Theintensity of blue colour increases again. Why ?

Using the readings obtained, the percentage of oxygen by volume in air can becalculated as follows.

Volume of air in the apparatus = (V1-10) cm3

Volume of oxygen in that volume of air = (V2-10) cm3

Percentage of oxygen in the sample of air =2

1

10 10010

VV

What are the probable errors in this experiment ?

Method IIMaterials and Equipments : Boiling tube Rubber bung Solid ferrousammonium sulphate or ferrous sulphate Solid sodium hydroxide Trough of water Measuring cylinder

Instructions :

pic

Rubber bung

Water + Ferrous sulphate + Sodium sulphate

71

Fill the boiling tube fully with water and fix the cork. Remove the cork and measure thevolume of water in the tube. Add 10 cm3 of water into the tube and also about 5 pelletsof sodium hydroxide. Also add about 3g of solid ferrousammonium sulphate or ferroussulphate and fix the cork tightly. Observes the colour of the product.

Shake the tube thoroughly for about 15 minutes. What change can be observed in thereaction mixture ? Place the inverted tube in the trough of water and open the bung underwater. What happens to the liquid level inside the tube? Keep the tube as it is for abouttwo minutes. Equate the liquid levels and measure the volume of the solution in the tube.Did the volume in the tube increase or decrease? Record your readings as follows.

Total volume of the boiling tube = V1cm3

Volume of water added = 10 cm3

Volume of water in the tube after the experiment = V2cm3

Note : When the inverted tube is placed in water, the water level rises. This implies that a part

of air inside the boiling tube has been removed. What products are formed by the reactionbetween ferrous sulphate and sodium hydroxide? Accordingly, the colour of theprecipitate anticipated here is dirty green. But the precipitate formed is brown in colour.This is due to the oxidation of ferrous hydroxide into ferric hydroxide.

Taking above observation and composition of air into consideration, infer which gas hasbeen removed from air during the reaction.

Which component of air acted as the oxidant here? Write balanced ionic equation for thereaction occurred.

Water Bung

Figure 49 (b)

72

What other salts can be used in place of ferrous salts?

Since excess ferrous sulphate is used, it can be assumed that all oxygen has reacted. Areagent bottle may also be used for this experiment instead of a boiling tube.

Total volume of the boiling tube = V1cm3

Volume of water added = 10 cm3

Volume of air in the tube = (V1-10) cm3

Volume of water in the tube after the experiment = V2cm3

Volume of oxygen in air in the tube = (V2-10) cm3

Percentage of oxygen in the sample of air =2

1

( 10) 10010)

VV

.Practical No. 50

Experimental determination of the alkalinity of a sample of water

Expected learning outcomes :1. States the ions affecting the alkalinity of a sample of water.2. Determines the alkalinity of a sample of water by a titration carried out with

phenolphthalein and methyl orange.

Materials and Equipments : Burette Pipette Titration flask Funnel Burette stand Beakers 0.02 mol dm-3 hydrochloric acid solution Phenolphthalein Methyl orange Stand

Instructions : Using a 100.00 cm3pipette or a burette, accurately measure 100.00 cm3of the

sample of water that has to be examined. Add 2 – 3 drops of phenolphthalein andtitrate with 0.02 mol dm-3 HCl solution. Note the burette reading at the end point.

Add methyl orange to the resulting mixture and titrate again with the 0.02 mol dm-3

HCl solution. Note the burette reading at the end point.

73

Note : The concentration of 2

3CO ions can be calculated using the burette reading ofthe first titration.

23 3H (aq) CO (aq) HCO (aq)

The concentration of 3HCO can be calculated using the burette reading of thesecond titration.

3 2 3H (aq) HCO (aq) H CO (aq) The change in pH of the solution and the variation of the concentrations of

species 2 3H CO , 3HCO and 23CO during the titration can be shown as follows.

Figure 50 (a)

For H2CO3, p 1aK , is 6.37 and p2aK is 10.33. Henderson equation can be

used to calculate the pH at points ‘a’ and ‘b’.At the instance ‘a’:

23 3H (aq) CO (aq) HCO (aq)

pH = pKa2+ log10

SaltAcid

pH = pKa2 + log10

23

3

[Co (aq)][HCO (aq)]

Since, [ 2-3CO (aq) ] = [ -

3HCO (aq) ] at ‘a’,

log10

SaltAcid

= 0

2a pH = pK

pH

[ 2 3H CO (aq)] [ -3HCO (aq) ]

ab

Concentration/mol dm-3

74

Find the approximate pH at ‘b’. When the pH of the water sample is known, we can get an idea of the quantities of

H2CO3, HCO3- and CO3

2- present in the sample. If OH- ions are present in the water samples, its pH is about 11. But, common water

samples do not have such high pH values. The total alkalinity of the sample of water is given in terms of the composition of

calcium carbonate (CaCO3).

Burette reading ( in the presence ofphenolphthalein + methyl oragen) = X cm3

Amount of H+0.021000 X mol

+2+

-3

-3

[H (aq)][Ca (aq)] =2

0.021000 1000 mol dm

2 1000.02 1000 = × 100 g dm2 100 1000

X

X

^as mass of CaCO3&

-3

-3

0.02 1000 × 100 = × 1000 mg dm2 100 1000

= 10 mg dm

X

X

Total alkalinity = 10 X mg dm-3

75

Practical No. 51

Experimental determination of the dissolved oxygen in a sample ofwater

Expected learning outcomes :1. Determines the amount of oxygen dissolved in water quantitatively.2. Acquires the skill of standardizing a sodium thiosulphate solution of approximately

known concentration by an iodometric titration.3. Acquires the skill of determining quantitatively the amount of iodine released by an

iodide using sodium thiosuphate.

Materials and Equipments : Manganese sulphate solution (prepared by dissolving about 4 g of manganese sulphate

in 10 cm3 of water) Alkaline potassium iodide solution (prepared by dissolving about 5g of sodium

hydroxide and about 1.5g of potassium iodide in 10 cm3 of water) Starch solution (prepared by dissolving about 1g of starch in 10 cm3 of hot water) 0.01 mol dm-3 sodium thiosulphate solution (prepared by dissolving 2.48 g of

N a 2S2O3.5H2O in 1000 cm3 of water). Add a little of sodium carbonate or

chloroform to this solution to prevent precipitation of sulphur. Concentrated sulphuric acid 1.0 mol dm-3 sulphuric acid solution Two reagent bottles, 250 cm3

Pipette, 25 cm3

Burette Measuring cylinder, 10 cm3

Titration flasks Funnels Glass tubes

Instructions :(a) Standardization of the sodium thiosulphate solution prepared Weigh 1.0-1.5 g of pure dry potassium iodate accurately. Dissolve in cold distilled

water and make up to 250 cm3 in a volumetric flask. Pipette out 25.00 cm3 of thissolution into a 250 cm3 volumetric flask and dilute again by adding distilled water tothe 250 cm3 mark (The concentration of this solution is nearly 0.01 mol dm-3). Measure25.00 cm3 of this solution to a titration flask and add 1g of potassium iodide and 5 cm3

of 1 mol dm-3 sulphuric acid. Titrate this solution with the sodium thisulphate solutionkept in the burette. Swirl the titration flask well. When the solution assumes a strawcolour, add about 200 cm3 of distilled water and 2 cm3of starch solution and continuethe titration. When the blue colour disappears stop addition of the solution and takethe reading.

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(b) Determination of the amount of oxygen dissolved in water

Fill the two reagent bottles completely with the sample of water in which the oxygencontent has to be found (The bottles should not contain air). As soon as the bottles arefilled, add 1 cm3 of the manganese sulphate solution and 2 cm3 of the alkaline potassiumiodide solution separately to the two bottles with a pipette or a glass tube. The tip ofthe pipette or the glass tube needs to be kept below the water level while transferringthe liquid. Here a little water may overflow. Close the bottle with the stopper andshake thoroughly. This should be done at the site of the water source. Take the bottlesto the laboratory.

After about 10 minutes (when the precipitate formed has settled ) to each bottle addabout 2 cm3 of concentrated sulphuric acid. Close the stopper and shake well. After10 minutes, pipette out 50.00 cm3 of the solution in the bottle to a titration flask andtitrate with the 0.01 mol dm-3 sodium thiosulphate solution kept in the burette till thesolution becomes light yellow in colour. Then add about 3 cm3 of the starch solutionand continue the titration till the blue colour disappears. Repeat same for the otherbottle.

Calculation :(a) Determination of the concentration of the sodium thiosulphte solution In acid medium,

potassium iodate oxidises iodide ions to iodine. (The reactions relevant to the calculationare given under competency level 14.1.2 of the Grade 13 Teacher Instuction Manual.)

(b) Determination of the amount of oxygen dissolved in water Manganese sulphate reactswith sodium hydroxide to form manganese hydroxide.

22Mn (aq) 2OH (aq) Mn(OH )(s)

(b) This manganese hydroxide precipitate is oxidized to manganese dioxide by the oxygendissolved in water.

In acidic medium, manganese dioxide oxidises iodide ions to iodine. ——— (3)

——— (4)

(3) + (4); 22 2 2MnO (s) 4H (aq) 2I (aq) Mn (aq) 2H O(l) I (aq)

Sodium thiosulphate titrates the liberated iodine.———— (5)

-2 2

-2 2 2

2 2 2 2

O (g) + 2H O(l) + 4e 4OH (aq) ----(1)

2Mn(OH) (s) + 4OH (aq) 2MnO (s) + 4H O(l) + 4e ---(2)

(1) + (2) 2Mn(OH) (s) + O (g) 2MnO (s) + 2H O(l)

22 2MnO (s) 4H (aq) 2e(aq) Mn (aq) 2H O(l)

22I (aq) I (aq) 2e

2I (aq) 2e 2I (aq)

77

2 22 3 4 62S O (aq) S O (aq) 2e ————— (6)

(5) + (6); 2 22 2 3 4 6I (aq) 2S O (aq) 2I (aq) S O (aq)

According to the stoichiometry of the above reactions:

22 3 2 2 2

12S O I MnO O2

22 3 2

1

4S O O

i.e. One mole of 22 3S O is equivalent to

1

4moles of O2

If the burette reading is Vcm3, the amount of 22 3S O spent =

C V mol1000 4

Amount of 2O in 50 cm3 of water =C V 1000 mol

1000 4 50

Amount of 2O in 1000 cm3 of water =C V 1000 mol

1000 4 50

Composition of 2O in water(in mg dm-3) =

= 160 CV mg dm-3

-3C V 1000 32 1000 mg dm1000 4 50

78

Practical No. 52

Experimental determination of the total dissolved solid (TDS) in asample of water

Expected learning outcomes :1. Understands that there are dissolved solids in water.2. Develops the skills in carrying out gravimetric methods.3. Compares the TDS values in different water samples.

Materials and Equipments : Water samples (collected from different places e.g. fresh water, tap water marine

water, etc.) Sand bath (or oil bath) Crucible Wire gauze Bunsen burner Triple beam balance Pipette

Instructions : Take a very clean, dry crucible and weigh it (W1). Accurately measure 25.0 cm3 of a water sample using a pipette and transfer it

carefully to the weighed crucible. Heat the crucible on a sand bath (or oil bath) by maintaining the temperature at 105 –

130 0C. Evaporate the water sample (in the crucible) to dryness on the sand bath. Then allow

it to cool down to the room temperature preferably in a desiccator and weigh thecrucible with its residue.

Continue heating for another 20 minutes and weigh it at room temperature. Continue heating for another 30 minutes and weigh it at room temperature [If a

constant mass is not obtained, continue the procedure until a constant mass isobtained (W2)].

Repeat the experiment using another sample of water.

Note :

TDS value of water = 32 1( ) 1000 1000mg dm25

w w