Chemistry - Activity Resources (teacher's Edition) - 3 Years

SPN 21 CHEMISTRY CHEMISTRY

ACTIVITY ACTIVITY RESOURCES RESOURCES Teacher’s Edition

OH SING SENG GOH SIAH CHING N RAZIMI HJ MUDA ROSLENA HJ MUNEL SRI YANI HEPNIE 1

Contents

Page

INTRODUCTION TO CHEMISTRY

3

KINETIC PARTICLE THEORY

4–8

CHEMICAL FORMULA

9

TYPES OF COMMON CHEMICAL REACTIONS

10 – 25

STOICHIOMETRY AND MOLE CONCEPT

26 – 29

EXPERIMENTAL CHEMISTRY

30 – 33

ACIDS, BASES AND NEUTRALIZATION

34 – 38

SALTS

39 – 46

QUALITATIVE ANALYSIS

47 – 55

METAL AND EXTRACTION

56 – 64

THE PERIODIC TABLE

65 – 66

ENERGY FROM CHEMICALS

67 – 70

ELECTROLYSIS

71 – 76

SPEED OF REACTIONS

77 – 83

REVERSIBLE REACTIONS

84 – 86

REDOX

87 – 88

ATMOSPHERE AND ENVIRONMENT

89

ORGANIC CHEMISTRY

90 – 99

REFERENCES

100

2

Activity 1.4

INTRODUCTION TO CHEMISTRY Aim:

Use mnemonics to familiarize with name and symbol of first row of common transition metals.

Example: Use mnemonics to memorise the first row of the common transition metals. Scandium, Titanium, Vanadium, Chromium, Manganese, Iron, Cobalt, Nickel, Copper, and Zinc Scary Tim Very Crooked Man I Call Nick Corporal Zee

-

3

Scandium Titanium Vanadium Chromium Manganese Iron Cobalt Nickel Copper Zinc

Activity 2.2

KINETIC PARTICLE THEORY Aim:

Using role play to demonstrate the movement of particles in solid, liquid and gas.

Procedure: 1. A group of students (6- 9), stand at the front of the room acting as atoms. 2. Standing next to each other in 2 or 3 rows with arms linked they represent a solid- particles closed together, moving slightly (vibrating) at their fixed position. 3. As heat energy is applied the students move further away from each other and eventually the links break. 4. Allowing students to move randomly passing each other but still remain close together. 5. Further heating the students move freely away from each other at high speed.

4

Activity 2.3

KINETIC PARTICLE THEORY Aim:

To determine the melting point of naphthalene using cooling curve.

Apparatus: Thermometer Retort stand Boiling tube

Conical flask Bunsen burner Stopwatch

Materials: Solid naphthalene Procedure: 1. Clamp boiling tube on retort stand. 2. Add 3 spatulas of powdered naphthalene into a boiling tube and insert a thermometer. 3. Heat the naphthalene until all has melted (about 85 C ). 4. Then, leave it to cool in a conical flask. 5. Record the temperature for every 20 seconds until it falls to about 60 C . 6. Record the results in a table.

5

Results: Cooling of naphthalene Time (s)

Temperature ( C )

Time (s)

0

160

20

180

40

200

60

220

80

240

100

260

120

280

140

300

Temperature ( C )

Analysis of data: 1. Plot a graph of temperature against time for the cooling of naphthalene. 2. From the graph deduce the melting point of naphthalene. 3. The melting point of naphthalene is …………………… C

6

Activity 2.4

KINETIC PARTICLE THEORY Aim:

To determine the purity of ethanol by determining its boiling point.

Apparatus: Thermometer

Tripod stand Wire gauze Bunsen burner Stopwatch Porcelain chip

3

Beaker 250 cm Retort stand Boiling tube Stirrer or glass rod Materials: Water Ethanol sample (Caution: ethanol is flammable)

Procedure: 1. Quarter fill the boiling tube with your ethanol sample. 2. Add a porcelain chip to the boiling tube to ensure it does not ‘froth up’ too much on boiling. 3. Set up apparatus as shown below.

7

4. Heat the water gently and stir continuously to ensure an even temperature around the boiling tube. 5. Continue heating until the ethanol boil. This is when the bubbles start to appear from the porcelain chip.

Results: 1. Record the highest reading on the thermometer. ……………….. C 2. Allow the ethanol to boil for one minute to see if the temperature change. Record this temperature. ……………… C Questions: 1. The boiling point of ethanol is 78 C . Is your ethanol sample pure? ……………………………………………………………………………………………………………… 2. What effect does an impurity have on the boiling point of a substance? ……………………………………………………………………………………………………………… 3. Why must we heat the ethanol in a water bath and not heat it directly with a Bunsen flame? ………………………………………………………………………………………………………………

8

Activity 5.1

CHEMICAL FORMULA Aim:

To work out the formula of ionic compound using card games.

Materials: A set of cards representing common anions and cations. Each card has the symbol of the ion written on it. Procedure: 1. Group the students (2 to 3 students in a group). 2. Give each group a set of cards. 3. To get a correct formula, join the shapes to form a rectangle. 4. The formula can then be read or copied from the card. For example, to find the formula of copper(II) chloride, a complete rectangle is formed by joining one copper(II) ion card and two chloride ion cards (see below). Hence the formula of copper(II) chloride is CuCl 2 .

Questions: Now use the card to work out the chemical formulae of the following ionic compounds (a) (b) (c) (d) (e) (f) (g) (h) (i)

Potassium chloride Zinc chloride Copper(II) oxide Potassium sulphate Potassium manganate(VII) Sodium hydrogencarbonate Potassium dichromate(VI) Magnesium hydroxide Sodium hydroxide

(j) Iron(II) sulphide (k) Sodium sulphate (l) Iron(III) hydroxide (m) Ammonium nitrate (n) Iron(II) nitrate (o) Iron(II) sulphate (p) Iron(III) sulphate (q) Ammonium sulphate

9

Activity 6.2

TYPES OF COMMON CHEMICAL REACTIONS Aim:

To show relative reactivity of metals with water.

Apparatus: Test tubes with rack Pair of forceps Knife or scalpel

White tile Water trough (for sodium only)

Materials: Sodium (store in oil) – teacher’s demonstration Calcium Magnesium Iron

Copper Distilled water

Procedure: 1. Place distilled water in four test tubes to a depth of 5 cm. 2. Drop a piece of calcium into a test tube filled with water. 3. Observe the reaction carefully and answer the following questions: (a) Does the metal float or sink in water? (b) Does the metal react vigorously? 4. Test the gas given out using lighted splint 5. Dip a piece of red litmus to the solution in the test tube. Is there a colour change? 6. Repeat step 1, 2, 3, and 4 for the other metals. 7. Record all your observations in a table provided under the Results section.

10

Results: Observations Element

Float or Sink

Lighted splint test

Vigorous or Not vigorous

Effect on red litmus paper

Magnesium Copper Calcium Sodium Iron From the results arranged the five metals in the order of decreasing reactivity. Most reactive

Least reactive

............................., .............................., ............................., ..............................., .............................. Questions: 1. Which group of the Periodic Table does sodium belong? ………………………………………………………………………………………………………..... 2. What is the common name for this group of metals? …………………………………………………………………………………………………………. 3. Name the alkali formed when sodium reacts with water. …………………………………………………………………………………………………………. 4. Name the gas produced when sodium reacts with water. ………………………………………………………………………………..................................... 5. Which other metals (listed above) will produce similar reaction with cold water? …………………………………………………………………………………………………………. 6. Name this type of reaction. …………………………………………………………………………………………………………. 7. Write a balanced equation for the above reaction. …………………………………………………………………………………………………………. 11

Activity 6.3

TYPES OF COMMON CHEMICAL REACTIONS Aim:

To show reaction between metals with dilute hydrochloric acid (Reactivity Series of Metals)

Apparatus: Test tubes with rack Chemicals: Zinc foil Magnesium ribbon Copper foil

Iron wire Dilute hydrochloric acid solution

Procedure: 1. Place dilute hydrochloric acid in a test tube to a depth of 2 cm. 2. Drop a piece of zinc into the test tube. 3. Test the gas given out using a lighted splint. 4. Repeat step 1, 2 and 3 for magnesium, copper and iron. 5. Record all your observations in a table provided in the Results section.

12

Results: Metals

Observations

Lighted splint test

Copper Zinc Magnesium Iron From the results arranged the four metals in the order of decreasing reactivity. Most reactive

Least reactive

....................................., ......................................, ....................................., .......................................

Questions: 1. Write the chemical formula for hydrochloric acid. .………………………………………………………………………………………………………… 2. Name the gas given out when dilute acid reacts with a metal. ................................................................................................................................................. 3. Name the metals (listed above) other than zinc that produces similar reaction with acid. …………………………………………………………………………………………………………. 4. What would you expect the reaction to be if potassium is used instead of zinc in the above reaction? …………………………………………………………………………………………………………. 5. Although copper is a metal, it does not react with dilute hydrochloric acid. Why? …………………………………………………………………………………………………………. 6. Name the salt formed when zinc reacts with dilute hydrochloric acid. …………………………………………………………………………………………………………. 7. Name this type of reaction …………………………………………………………………………………………………………. 8. Write a balanced equation for the above reaction. …………………………………………………………………………………………………………. 13

Activity 6.4

TYPES OF COMMON CHEMICAL REACTIONS Aim:

To show reaction of carbonates with dilute hydrochloric acid.

Apparatus: Test tube with rack Boiling tube

Delivery tube (or plastic syringe) Spatula

Chemicals: Calcium carbonate powder Copper(II) carbonate powder Dilute hydrochloric acid solution Limewater Procedure: 1. The test tube is filled with dilute hydrochloric acid to a depth of 3 cm. 2. One spatula of powdered calcium carbonate is added to the dilute hydrochloric acid. 3. Pass the gas given out into limewater. 4. Repeat step 1 and 3 for powdered copper(II) carbonate. 5. Record all your observations in the table provided.

Calcium carbonate powder Boiling tube

Dilute hydrochloric acid

14

Results: Observations

Limewater test

Calcium carbonate

Copper(II) carbonate

Questions: Calcium carbonate 1. Name the gas given out in the reaction. …………………………………………………………………………………………………………. 2. Is calcium carbonate soluble in water? …………………………………………………………………………………………………………. 3. Name the salt formed when calcium carbonate reacts with dilute hydrochloric acid? …………………………………………………………………………………………………………. 4. Is this salt soluble in water? …………………………………………………………………………………………………………. 5. Name this type of reaction. …………………………………………………………………………………………………………. 6. Write a balanced equation for the above reaction. …………………………………………………………………………………………………………. Copper(II) carbonate 1. What is the colour of copper(II) carbonate? ……………………………………………………………………………………………………........ 2. Is copper(II) carbonate soluble in water? …………………………………………………………………………………………………………. 3. Name the gas given out in the above reaction. …………………………………………………………………………………………………………. 4. Write a balanced chemical reaction for the above reaction. …………………………………………………………………………………………………………. 15

Activity 6.5

TYPES OF COMMON CHEMICAL REACTIONS Aim:

To show precipitation reaction.

Apparatus: Test tubes Materials: Potassium iodide solution Lead(II) nitrate solution Silver nitrate solution

Sodium sulphate solution Barium chloride solution Dilute hydrochloric acid solution

Procedure: 1. Pour potassium iodide solution into a test tube (about 2 cm depth) and add an equal volume of lead(II) nitrate solution and observe. 2. Leave the mixture to stand for a few minutes and observe. 3. Repeat step 1 and 2 for: (a) silver nitrate and dilute hydrochloric acid solutions (b) Sodium sulphate and barium chloride solutions

Results: Reaction

Initial observations

(a) Potassium iodide + Lead(II) nitrate solution

(b) Silver nitrate + Dilute hydrochloric acid

(c) Sodium sulphate + barium chloride

16

When left to stand

Questions: 1. Name the precipitate formed in experiment (a), (b) and (c). Experiment (a): ………………………………….. Experiment (b): ………………………………….. Experiment (c): ………………………………….. 2. Write the balanced chemical equations for experiment (a), (b) and (c). Experiment (a): …………………………………………………………………………………………... Experiment (b): …………………………………………………………………………………………... Experiment (c): ……………………………………………………………………………………………

17

Activity 6.6

TYPES OF COMMON CHEMICAL REACTIONS Aim:

To show displacement reaction between metals.

Apparatus: Sand paper

Beaker 50 cm3 Measuring cylinder Materials: Magnesium ribbon Lead(II) nitrate solution Zinc foil Copper(II) sulphate solution

Silver nitrate solution Magnesium chloride solution Copper foil

Procedure: 1. Clean a strip of magnesium ribbon with a sand paper. 2. Pour about 20 cm3 of lead(II) nitrate solution into a beaker. 3. Immerse the magnesium ribbon in the lead(II) nitrate solution. 4. Leave the mixture aside for few minutes, and observe. 5. Repeat the above experiment using the following: (i)

Zinc foil and copper(II) sulphate solution,

(ii)

Magnesium ribbon and silver nitrate solution,

(iii) Copper foil and magnesium chloride solution.

18

Results: Reaction

Observations

Magnesium + lead(II) nitrate solution

Zinc + copper(II) sulphate solution

Magnesium + silver nitrate solution

Copper + Magnesium chloride solution

19

Conclusion

Activity 6.7

TYPES OF COMMON CHEMICAL REACTIONS Aim:

To show displacement reaction between halogens.

Apparatus: Test tubes Materials: Chlorine water Bromine water Sodium bromide solution

Potassium iodide solution Iodine Sodium chloride solution

Procedure: 1. Pour potassium iodide solution into a test tube (about 2 cm depth). 2. Add an equal volume of chlorine water and observe. 3. Record your observations in the table provided. 4. Repeat the experiment for the following reaction: (i)

Chlorine water and sodium bromide solution,

(ii)

Bromine water and sodium chloride solution,

(iii) Bromine water and potassium iodide solution, (iv) Iodine and sodium chloride solution, (v) Iodine and sodium bromide solution.

20

Results: Halogen

Halides NaCl (aq)

NaBr (aq)

KI (aq)

Chlorine

Bromine

Iodine

Questions: 1. From the results, arrange the three halogens in the order of decreasing reactivity. Most reactive

Least reactive

…………………………., …………………………., ………………………….. 2. Write the balanced chemical equation for the following reactions: (a) Chlorine and sodium bromide ……………………………………………………………………………………………………….. (b) Chlorine and potassium iodide ……………………………………………………………………………………………………….. (c) Bromine and potassium iodide ………………………………………………………………………………………………………..

21

Activity 6.8

TYPES OF COMMON CHEMICAL REACTIONS Aim:

To show thermal decomposition of carbonates.

Apparatus: Pyrex glass test tube Test tube with rack Test tube holder

Delivery tube (or plastic syringe) Bunsen burner

Materials: Copper(II) carbonate powder Ammonium carbonate powder

Limewater Zinc carbonate powder

Procedure: 1. Put some powdered copper(II) carbonate into a Pyrex test tube and heat it strongly. 2. Pass the gas given out through limewater and observe. 3. Record your observations in the table provided. 4. Repeat the experiment with ammonium carbonate and zinc carbonate.

22

Results: Reaction

Observations

Limewater test

(a) Copper(II) carbonate

(b) Zinc carbonate

(c) Ammonium carbonate

Questions: 1. What is the colour of the residue in experiment (a)? ……………………………………………………………………………………………………………… 2. What is the colour of the residue in experiment (b)? ……………………………………………………………………………………………………………… 3. Do you see any residue in experiment (c)? Why? ……………………………………………………………………………………………………………… 4. Write the balanced chemical equation for all the reactions. (i)

Experiment (a): ……………………………………………………………………………………..

(ii)

Experiment (b): ……………………………………………………………………………………..

(iii) Experiment (c): ……………………………………………………………………………………..

23

Activity 6.9

TYPES OF COMMON CHEMICAL REACTIONS Aim:

To show direct reactions by heating

Part 1: Heating iron with sulphur Apparatus: Crucible Pipeclay triangle Tripod stand

Tong Spatula Bunsen burner

Materials: Iron filings Sulphur powder Procedure: 1. Mix one spatula full sulphur powder with an equal amount of iron filing in a crucible. 2. Heat the mixture strongly and observe.

Results: Observation: …………………………………………………………………………………………………... Questions: 1. Name the product of the reaction. ……………………………………………………………………………………………………………… 2. Write the equation for the reaction. ……………………………………………………………………………………………………………… 24

Part 2: Heating sulphur with air (oxygen) Apparatus: Crucible Pipeclay triangle Tripod stand

Tong Spatula Bunsen burner

Materials: Sulphur powder Blue litmus paper

Strip filter paper Acidified potassium dichromate(VI)

Procedure: 1. Add one spatula full sulphur powder into a crucible. 2. Then, heat strongly and observe. 3. Test the gas given out using moist blue litmus paper and acidified potassium dichromate(VI) paper. 4. Record your observations in the table Results: Test

Observations

Using moist blue litmus paper

Using acidified potassium dichromate(VI) paper

Questions: 1. Name the product of the reaction. ……………………………………………………………………………………………………………… 2. Write the equation for the reaction. ………………………………………………………………………………………………………………

25

Activity 7.1

STOICHIOMETRY AND MOLE CONCEPT Aim:

To prepare standard solution of copper(II) sulphate.

Part 1: Preparation of 0.1 mol dm 3 copper(II) sulphate solution Apparatus: Glass rod

Volumetric flask 250 cm3 Beaker Balance Materials: Distilled water Copper(II) sulphate crystals

Procedure: 1 Use the balance to weigh _______ g of copper(II) sulphate pentahydrate crystals. 2

Dissolve the copper(II) sulphate in distilled water inside a beaker.

3

Pour the solution into a 250 cm3 volumetric flask.

4

Add distilled water until the graduation mark on the neck of the graduated flask.

Results: The solution prepared is 0.1 mol dm 3 copper(II) sulphate solution. The solution contains …………………. mole CuSO 4 .5H2 O in 1 dm3 solution.

26

Part 2: Preparation of 0.01 mol dm 3 copper(II) sulphate solution from 0.1 mol dm 3 copper(II) sulphate solution. Procedure: 1

Pipette 25 cm3 of copper(II) sulphate solution from Part 1.

2

Pour the solution into a 250 cm3 volumetric flask.

3

Add distilled water until the graduated mark on the neck of the graduated flask.

Results: The solution prepared is 0.01 mol dm 3 copper(II) sulphate solution. The solution contains ……………… mole CuSO 4 .5H2 O in 1 dm3 solution. Questions:

Dilution factor x Original concentrat ion  Final concentrat ion 1. Calculate the dilution factor from the above equation? ………………… X

27

Activity 7.2

STOICHIOMETRY AND MOLE CONCEPT Aim:

To determine the percentage purity of sodium carbonate in a mixture of sodium carbonate and ammonium carbonate.

Apparatus: Bunsen burner with matches Tripod stand Crucible

Pipeclay triangle Tongs Spatula

Materials: Sample of impure sodium carbonate (sodium carbonate mixed with ammonium carbonate) Procedure: 1. Weigh the mass of the crucible. 2. Then add in sample of impure sodium carbonate and weigh. 3. Place the crucible on the tripod stand, and heat it strongly for five minutes. 4. Leave the crucible to cool and weigh the content again.

28

Results: (a) Mass of the empty crucible

=

……………….. g

(b) Mass of the crucible + impure sample of sodium carbonate

=

……………….. g

(c) Mass of impure sample of sodium carbonate [(b) – (a)]

=

……………….. g

(d) Mass of crucible after heating

=

……………….. g

(e) Mass of pure sodium carbonate [(d) – (a)]

=

……………….. g

Formula to calculate percentage purity: percentage purity =

mass of pu re subs tan ce in samp le × 100% mass of sa mple

Questions: Calculate the percentage purity of sodium carbonate.

29

Activity 8.1

EXPERIMENTAL CHEMISTRY Aim:

To obtain copper(II) sulphate crystals from a mixture of copper(II) sulphate and sand.

Apparatus: Stirring rod Evaporating dish Bunsen burner Tripod stand

250 cm3 beaker 250 cm3 conical flask Filter paper Filter funnel Materials: A mixture of copper(II) sulphate with sand Procedure: 1. Put the mixture in a beaker 2. Add water to the mixture

3. Stir with a glass rod to make sure that copper(II) sulphate dissolved. 4. Filter the mixture; collect the filtrate in a conical flask. 5. Wash the residue on the filter paper with water. 6. Evaporate the filtrate until saturated, and leave it to crystallize.

30

Results: 1. The residue is

…………………………………

The filtrate is

………………………………...

The solvent is

…………………………………

2. Explain why the residue is washed with water. ……………………………………………………………………………………………………………… ……………………………………………………………………………………………………………… 3. Give an example of a mixture of two other substances that can be separated by this method. ………………………………………………………………………………………………………………

31

Activity 8.5

EXPERIMENTAL CHEMISTRY Aim:

To separate various dyes in food colouring and measure the R f values.

Apparatus: Boiling tube Filter paper Ruler

Pair of scissors Distilled water

Materials: Mixture of food colouring Procedure: 1. Cut a strip of filter paper such that it can fit neatly into a boiling tube. It should also be slightly longer than the boiling tube. 2. Draw a baseline with a pencil, about 2 cm away from the bottom tip. 3. Mark a tiny spot on the middle of the baseline with food colouring. 4. Put about 1cm depth of distilled water into the boiling tube. 5. Mount the filter paper strip in the boiling tube. 6. Leave the apparatus to stand for a short while. 7. Observe the solvent front as the water travels up the paper. 8. Remove the piece of filter paper when the solvent front reaches just below (1cm) the top of the paper. [DO NOT let the solvent front go beyond the top of the paper] 9. Allow the paper to dry.

spot

baseline

32

Results: Attach your chromatogram in the space below.

Questions: 1. What is the principle behind paper chromatography? ……………………………………………………………………………………………………………… ……………………………………………………………………………………………………………… 2. Calculate the R f for each colour.

33

Activity 9.1

ACIDS, BASES AND NEUTRALIZATION Aim:

To neutralise hydrochloric acid by titrating with sodium hydroxide solution.

Apparatus:

Materials: Solution P is dilute hydrochloric acid of unknown concentration Solution Q is 0.1 mol dm3 sodium hydroxide solution Methyl orange or screened methyl orange indicator. Procedure: 1. Fill the burette with solution P until 0 mark. Make sure that no air bubbles are trapped at the tip of the burette. 2. Pipette a 25.0 cm3 portion of Q into a conical flask. Add a few drops of either screened methyl orange or methyl orange indicator. 3. Titrate the solution Q with the solution P from the burette and record the results in the table, repeating the titration as many times as you consider necessary to achieve consistent results.

34

Results: Titration number

1

2

3

Final burette reading / cm3 Initial burette reading / cm3 Volume of P used / cm3 Best titration results () Summary: Tick () the best titration results. Using these results, the average volume of P required was …………………………. cm3 . Volume of Q used was ……………….………… cm3 Hence, …………. cm3 of NaOH required ................ cm3 of HCl Questions: 1. Calculate the number of moles of sodium hydroxide used in the titration.

2. Write a balanced equation for the neutralisation of HCl by NaOH.

3. How many moles of HCl would be used in neutralising NaOH in the titration?

4. Having known the number of moles of hydrochloric acid from (3) and also the volume of the acid used in the titration, calculate the concentration in mol dm3 of the hydrochloric acid.

35

Activity 9.2

ACIDS, BASES AND NEUTRALIZATION Aim:

To titrate sodium carbonate and hydrochloric acid and to find percentage purity of sodium carbonate.

Materials: Solution P is 0.2 mol dm3 hydrochloric acid Solution Q is made by dissolving 28 g of a mixture of sodium chloride and sodium carbonate in water and made up to 1 dm3 solution. Procedure: You are required to find the percentage purity of sodium carbonate in the mixture. 1. Put solution P in a burette. 2. Pipette a 25.0 cm3 portion of Q into a conical flask. Add a few drops of either screened methyl orange or methyl orange indicator. 3. Titrate the solution Q with the solution P from the burette and record the results in the table, repeating the titration as many times as you consider necessary to achieve consistent results. Results: Titration number

1

2

3

Final burette reading / cm3 Initial burette reading / cm3 Volume of P used / cm3 Best titration results () Summary: In the above titration …………… cm3 of solution P is required to exactly neutralise …………... cm3 of solution Q.

36

Questions: 1. Write a chemical equation to show the reaction involved in the titration.

2. From the titration, calculate the concentration of sodium carbonate in solution Q in mol dm3 .

3. Given that the molecular mass of sodium carbonate, Na 2 CO 3 is 106, calculate the concentration of sodium carbonate in solution Q in g dm 3 .

4. Calculate the percentage purity by mass of sodium carbonate in the original mixture used in the preparation of solution Q.

37

Activity 9.3

ACIDS, BASES AND NEUTRALIZATION Aim:

To show reaction between sodium hydroxide and ammonium chloride

Apparatus: Test tube with rack Test tube holder

Bunsen burner Spatula

Materials: Sodium hydroxide solution Solid ammonium chloride Litmus paper Procedure: 1. Put one spatula full solid ammonium chloride into a test tube and slowly pour in sodium hydroxide solution. Smell the gas. 2. Warm the mixture gently by using a small Bunsen flame. 3. Test the gas by using moist red and blue litmus paper. Results: 1. When aqueous sodium hydroxide is added to solid ammonium chloride and warmed gently a gas is formed. Describe the smell of the gas. ……………………………………………………………………………………………………………… 2. Test the gas with moist blue litmus paper. Describe your observation. ……………………………………………………………………………………………………………… 3. Test the gas with moist red litmus paper. Describe your observation. ……………………………………………………………………………………………………………… Conclusion: 1. When an ammonium salt reacts with an alkali …………………… gas is formed, which turn moist ………….. litmus paper ………………. . 2. Write an equation for the reaction between sodium hydroxide and ammonium chloride. ………………………………………………………………………………………………………………

38

Activity 10.1

SALTS Aim:

To prepare copper(II) sulphate crystals by reacting sulphuric acid with copper(II) oxide or copper(II) carbonate

Apparatus: Beaker Conical flask Stirrer Filter funnel

Filter paper Evaporating basin Spatula

Materials: Dilute sulphuric acid Copper(II) oxide or copper(II) carbonate Procedure: 1. Put about 100 cm3 of dilute sulphuric acid in a beaker and heat it gently. 2. Add copper(II) oxide or copper(II) carbonate to the hot sulphuric acid, a little at a time until in excess. 3. Filter out the excess copper(II) oxide or copper(II) carbonate by using filter funnel and filter paper. 4. Collect the filtrate in an evaporating basin and evaporate the filtrate until saturated. 5. Leave the saturated filtrate to cool and crystallise.

39

Questions: 1. Name the base or carbonate used in this reaction? ……………………………………………………………………………………………………………… 2. What is the formula of this compound? ……………………………………………………………………………………………………………… 3. Is this compound soluble or insoluble? ……………………………………………………………………………………………………………… 4. Name the above reaction. ……………………………………………………………………………………………………………… 5. Write chemical equation for the above reaction. ……………………………………………………………………………………………………………… 6. Name the salt formed from the above reaction. ……………………………………………………………………………………………………………… 7. Name the filtrate. ……………………………………………………………………………………………………………… 8. What is the residue left on the filter paper? ……………………………………………………………………………………………………………… 9. Why must the copper(II) oxide be added in excess? ……………………………………………………………………………………………………………… 10. Can the above reaction be carried out using titration method? Why? ………………………………………………………………………………………………………………

40

Activity 10.2

SALTS Aim:

To prepare insoluble salts

Apparatus: Test tubes with rack Materials: Silver nitrate solution Potassium iodide solution

Dilute hydrochloric acid solution Dilute nitric acid solution

Procedure: Carry out the following tests and record your observations in the table. Test no. 1

Test

observations

(a) To a portion of silver nitrate solution, add dilute hydrochloric acid until a change is seen. (b) Leave the mixture to stand for a few minutes and observe.

2

(a) To a portion of potassium iodide solution, add lead(II) nitrate solution until a change is seen. (b) Leave the mixture to stand for a few minutes and observe.

41

Questions: 1. Name the above reaction. ……………………………………………………………………………………………………………… 2. Write the symbol and ionic equations for the reactions between silver nitrate and hydrochloric acid. (a) Symbol equation ……………………………………………………………………………………………………….. (b) Ionic equation ……………………………………………………………………………………………………….. 3. Write the symbol and ionic equations for the reactions between potassium iodide and lead(II) nitrate. (a) Symbol equation ……………………………………………………………………………………………………….. (b) Ionic equation ………………………………………………………………………………………………………..

42

Activity 10.3

SALTS Aim: To investigate solubility of salts in water Apparatus: Test tubes with rack Stoppers Spatula Materials: Distilled Water Solid salts: Silver chloride Sodium chloride Lead(II) chloride Barium sulphate Copper(II) sulphate

Lead(II) sulphate Sodium carbonate Calcium carbonate Potassium nitrate Silver nitrate

Procedure: 1. Add half a spatula of a salt to a test tube. 2. Half-fill the test tube with distilled water. Then stopper the tube and shake well. 3. Record your observation in the table provided. Results: Solubility (soluble or insoluble)

Salt Silver chloride Sodium chloride Lead(II) chloride Barium sulphate Copper(II) sulphate Lead(II) sulphate Sodium carbonate Calcium carbonate Potassium nitrate Silver nitrate 43

Questions: State whether you think the following substances are soluble or insoluble in water: (a) Zinc nitrate

……………………………

(b) Potassium nitrate

……………………………

(c) Copper(II) carbonate

……………………………

(d) Sodium nitrate

……………………………

(e) Ammonium chloride

……………………………

44

Activity 10.4

SALTS Aim: To determine the solubility of salts in g cm 3 . Apparatus: Beakers Glass rod Spatula Balance

Filter funnel Filter paper Measuring cylinder

Materials: Water Solid salts: sodium chloride and copper(II) sulphate Procedure: 1. Label each beaker with the name of the salts. Measure 100 cm3 water into each beaker. 2. Record the mass of these beakers containing water. 3. Add a spatula of sodium chloride salt and stir till dissolved. 4. Repeat step 3 until no more salt can dissolve in the 100 cm3 of water. 5. Filter this salt solution and collect the filtrate. 6. Record the mass of this salt solution (filtrate). 7. Subtract this new mass with the previous mass. 8. The difference will represent the solubility of the salt in g 100cm 3 . 9. This solubility can then be converted to g dm 3 . 10. Repeat the experiment with copper(II) sulphate salt. 11. Record your results in the table.

45

Results: Sodium chloride salt Mass of beaker + 100 cm3 water / g

(a)

Mass of salt solution (filtrate) / g

(b)

Solubility in g 100cm 3

(a) – (b)

Solubility in g dm 3

Conclusion:

46

Copper(II) sulphate salt

Activity 11.1

QUALITATIVE ANALYSIS Aim: To identify the following cations: Al3  , NH 4 , Ca 2  , Cu 2  , Fe 2  , Fe 3 and Zn 2  Apparatus: Test tubes with rack Test tube holder Bunsen burner Materials: Aqueous sodium hydroxide solution Aqueous ammonia solution Aqueous solution of the following salts: Aluminium sulphate – A Ammonium chloride – B Calcium chloride – C Copper(II) sulphate – D

Iron(II) sulphate – E Iron(III) chloride – F Zinc nitrate – G

Procedure: A, B, C, D, E, F and G are unknown solutions containing different cations. Carry out the following test and record your observations in the table. You should test and name, where possible, any gases evolved. Test no. 1

Test

Observations

(a) To a portion of A add aqueous sodium hydroxide until a change is seen. (b) Add an excess of aqueous sodium hydroxide to the mixture from (a).

(c) To a portion of A add aqueous ammonia until a change is seen. (d) Add an excess of aqueous ammonia to the mixture from (c).

47

Test no. 2

Test

Observations

(a) To a portion of B add equal volume of aqueous sodium hydroxide. (b) Then warm gently (c) To a portion of B add equal volume of aqueous ammonia.

3

(a) To a portion of C add aqueous sodium hydroxide until a change is seen. (b) Add an excess of aqueous sodium hydroxide to the mixture from (a). (c) To a portion of C add equal volume of aqueous ammonia.

4

(a) To a portion of D add aqueous sodium hydroxide until a change is seen. (b) Add an excess of aqueous sodium hydroxide to the mixture from (a).

(c) To a portion of D add aqueous ammonia until a change is seen.

(d) Add an excess of aqueous ammonia to the mixture from (c).

48

Test no. 5

Test

Observations

(a) To a portion of E add aqueous sodium hydroxide until a change is seen. (b) Add an excess of aqueous sodium hydroxide to the mixture from (a). (c) To a portion of E add aqueous ammonia until a change is seen. (d) Add an excess of aqueous ammonia to the mixture from (c).

6

(a) To a portion of F add aqueous sodium hydroxide until a change is seen. (b) Add an excess of aqueous sodium hydroxide to the mixture from (a). (c) To a portion of F add aqueous ammonia until a change is seen. (d) Add an excess of aqueous ammonia to the mixture from (c).

7

(a) To a portion of G add aqueous sodium hydroxide until a change is seen. (b) Add an excess of aqueous sodium hydroxide to the mixture from (a). (c) To a portion of G add aqueous ammonia until a change is seen. (d) Add an excess of aqueous ammonia to the mixture from (c). 49

Questions: Give the names and formulae of the cations present in: Name

Formula

Solution A Solution B Solution C Solution D Solution E Solution F Solution G

50

Activity 11.2

QUALITATIVE ANALYSIS Aim: To identify the following anions: CO 23 , Cl  , I , NO 3 and SO 24 Apparatus: Test tubes with rack Test tube holder Bunsen burner Materials: Aqueous solution of the anions above: Sodium carbonate – A Sodium chloride – B Potassium iodide – C Silver nitrate – D Sodium sulphate - E

Aqueous barium chloride Aqueous lead(II) nitrate Aqueous sodium hydroxide Aluminum foil Dilute nitric acid

Students are required to carry out tests to identify the anions in solutions A, B, C, D and E. State either positive (  ) or negative ( X ) for each test using the reagents below Solution

Dilute nitric acid

AgNO 3 (aq)

BaCl 2 (aq)

A B C D E

51

Pb(NO 3 ) 2 (aq)

NaOH + Al and heat

52

E

D

C

B

A

Solution

Reagent giving positive test Procedure

Write the positive test for the anions A, B, C, D and E in the spaces provided Observations

Anion present

Activity 11.3

QUALITATIVE ANALYSIS Aim:

To test for gases

Apparatus: Bunsen burner Litmus paper Spatula Filter paper

Test tube holder Plastic syringe Test tubes with rack Wooden splint

Materials: Acidified potassium dichromate (VI) Distilled water Ammonium chloride solution Hydrogen peroxide solution Bleach solution Limewater

Dilute hydrochloric acid Solid calcium hydroxide Manganese(IV) oxide Sodium sulphite Calcium carbonate Magnesium ribbon

Procedure: 1. Generate the gas as described in the table. 2. Note its colour and odour and record your observations in the table 3. Carry out specific tests described in column 3 and record your observations. Note: Use a test tube holder when heating anything in a test tube.

53

Test No. 1

Gas

observations

Hydrogen Put one magnesium ribbon into a test tube.

Specific Test and observations Wooden splint test Insert a lighted wooden splint into the mouth of the test tube.

Then add 2 – 3 cm 3 of dilute hydrochloric acid into the test tube. 2

Oxygen

Wooden splint test Insert a glowing wooden splint into the mouth of the test tube.

Put 2 – 3 cm 3 of hydrogen peroxide into a test tube. Then add a small amount of manganese(IV) oxide into the test tube. 3

Carbon dioxide Put calcium carbonate into a test tube.

Limewater test Collect the gas using plastic syringe then pass the gas into test tube containing limewater then shake.

Then add 2 – 3 cm 3 of dilute hydrochloric acid into the test tube. 4

Sulphur dioxide Put sodium sulphite into a test tube.

Litmus test Hold moist blue litmus paper in the gas.

Then add 2 – 3 cm 3 of dilute hydrochloric acid into the test tube and heat.

Potassium dichromate(VI) test Hold a piece of filter paper dipped in acidified potassium dichromate(VI) in the gas. Potassium manganate(VII) test Hold a piece of filter paper dipped in acidified potassium manganate(VII) in the gas.

54

Test No. 5

Gas

observations

Chlorine

Specific Test and observations Litmus test Hold moist blue litmus paper in the gas.

3

Add about 2 cm of dilute hydrochloric acid into a test tube. Then add about 2 cm 3 of bleach solution. 6

Ammonia Put ammonium chloride solution into a test tube.

Litmus test Hold moist red litmus paper in the gas.

Then add one spatula of solid calcium hydroxide and warm gently. 7

Water vapour One third fill a boiling tube with water

Hold a piece of blue cobalt chloride paper in the gas.

Then heat gently.

8

Nitrogen dioxide (N 2006, P 3)

Litmus test Hold moist blue litmus paper in the gas.

Put one spatula of sodium nitrite in a test tube. Then add about 2 cm3 of dilute hydrochloric acid.

55

Activity 12.1

METALS AND EXTRACTION Aim:

To compare the reactivity of metals by displacement reaction.

Apparatus: Test tubes with rack Sandpaper Spatula Materials: Magnesium ribbon Zinc foils Iron fillings Copper foils

1 mol dm3 Zinc sulphate solution 1 mol dm3 Iron (II) sulphate solution 1 mol dm3 Copper (II) sulphate solution

3

1 mol dm Magnesium sulphate solution

Procedure: 1. Clean a strip of magnesium ribbon with sandpaper. 2. Half fill a test tube with zinc sulphate solution. 3. Immerse the magnesium ribbon in the zinc sulphate solution. 4. Leave the mixture aside for a few minutes, and observe. 5. Repeat the experiment above using the materials stated in the table.

56

Results: Observations with Magnesium

Zinc

Iron

Copper

Magnesium sulphate solution

Zinc sulphate solution

Iron(II) sulphate solution

Copper(II) sulphate solution

Questions: 1. From your results arrange the four metals in order of decreasing reactivity. Most reactive

Least reactive

................................... , ................................... , ................................... , ................................... . 2. Complete these displacement equations: CuSO4 (aq) + Fe (s)

 

__________

+

_________

FeSO4 (aq) + Zn (s)

 

__________

+

_________

ZnSO4 (aq) + Mg (s)

 

__________

+

_________

3. Aluminium is more reactive than zinc but less reactive than magnesium. Will aluminium displace iron from a solution of iron (II) salt? Explain. ……………………………………………………………………………………………………………… ……………………………………………………………………………………………………………… 57

Activity 12.2

METALS AND EXTRACTION Aim:

To show Thermit reaction (reduction of metal oxide) – demonstration only.

Apparatus: Sand bath Paper cone made from filter paper Spatula Materials: Iron(III) oxide Aluminium powder Magnesium powder

Magnesium ribbon Barium peroxide (or potassium chlorate)

Procedure: 1. Mix a few grams of iron(III) oxide powder with an equal amount of aluminium powder and put the mixture in a paper cone mounted on a sand bath as shown in the diagram. 2. Mix a little barium peroxide or potassium chlorate with magnesium powder and pour the mixture into the paper cone containing the mixture of iron(II) oxide and aluminium powder. This mixture is to set off a preliminary reaction. 3. Using a long piece of clean magnesium ribbon as a fuse, stick one end into the base of the paper cone and burn the other end. [Caution: Immediately move far away from the burning magnesium ribbon.]

Magnesium ribbon as fuse

Mixture of barium peroxide and magnesium powder

Paper cone Mixture of iron(III) oxide and aluminium powder Sand bath

58

Results: 1. Describe the reaction. ……………………………………………………………………………………………………………… ……………………………………………………………………………………………………………… ………………………………………………………………………………………………………………

2. Write the equation to show the reaction between iron(III) oxide and aluminium. ……………………………………………………………………………………………………………...

59

Activity 12.3

METALS AND EXTRACTION Aim:

To show the action of heat on the carbonates.

Apparatus: Test tubes with rack Test tube holders

Bunsen burner Plastic syringe (or delivery tube)

Materials: Zinc carbonate Copper(II) carbonate Lime water Procedure: 1. Put some powdered zinc carbonate into a test-tube and heat it strongly. 2. Pass the gas given out through lime water and observe for the formation of a new substance. 3. Repeat the experiment using copper(II) carbonate and observe for the colour change and the formation of a new substance. Results: (a) Zinc carbonate is …………………… (colour). When strongly heated zinc carbonate gives out a …………………… (colour) gas which turns lime water …………………… (b) The residue is …………………… (colour) when hot and ……………………(colour) when cold. (c) The gas is …………………… and the residue is …………………… (d) Write word equation for the above reaction. ……………………………………………………………………………………………………………… (e) Write symbol equation for the above reaction. ……………………………………………………………………………………………………………… (f) Copper(II) carbonate is ……………………(colour). When strongly heated copper(II) carbonate gives out a ……………………(colour) gas which turns lime water …………………… (g) The residue is ……………………(colour) (h) The gas is …………………… and the residue is …………………… (i) Write word equation for the above reaction. ……………………………………………………………………………………………………………… (j) Write symbol equation for the above reaction. ……………………………………………………………………………………………………………… 60

Activity 12.4

METALS AND EXTRACTION Aim:

To determine the conditions for rusting

Apparatus: Boiling tubes Test tube rack Materials: Iron nails Boiled water

Cooking oil Tap water

Procedure: 1. Prepare four boiling tubes in a rack, labelled them as A, B, C and D. 2. Into each boiling tube place a clean iron nail. 3. Pour water from the tap into boiling tube A until the nail is fully submerged. 4. Pour some hot water into boiling tube B until the nail is fully submerged and pour a layer of cooking oil to cover the surface of the water. 5. Place a little anhydrous cobalt chloride into boiling tube C and cork the mouth of the boiling tube with a rubber bung. 6. Pour some cooking oil into boiling tube D until the nail is fully submerged. 7. Leave the four boiling tubes in the rack for a few days and then observe. Results: (a) Does the tap water in boiling tube A contain air? Why? ……………………………………………………………………………………………………………… (b) Does the water in boiling tube B contain air? Why? ……………………………………………………………………………………………………………… (c) Does the air in the boiling tube C contain water? Why? ……………………………………………………………………………………………………………… (d) Does boiling tube D contain any air or water? ……………………………………………………………………………………………………………… (e) Which iron nail becomes rusty? ……………………………………………………………………………………………………………… (f) Why does the iron nail become rusty? ……………………………………………………………………………………………………………… 61

Activity 12.5

METALS AND EXTRACTION Aim:

To show sacrificial protection of metal.

Apparatus: 3 Petri dishes Cotton wool Materials: Salt solution Iron nails Magnesium ribbon Copper foil Procedure: 1. Fill the three Petri dishes with some salt solution and then make a cushion of cotton wool and place it in each Petri dish so that it is soaked in the salt solution. 2. Label the Petri dishes as A, B, and C. 3. In Petri dish A, place a clean iron nail on the cushion of cotton wool. In Petri dish B, place a clean iron nail wound with magnesium ribbon and in Petri dish C, place a clean iron nail wound with copper strip. 4. Leave the three Petri dishes for a few days and then observe.

cotton wool soaked in salt solution

A

B Magnesium ribbon

62

C Copper foil

Results: 1. What happen to the iron nail in A? ……………………………………………………………………………………………………………... 2. Which iron nail shows no rusting? Why? ……………………………………………………………………………………………………………… 3. What is the difference between iron nail in A from that in C? ……………………………………………………………………………………………………………… 4. Explain your observation in 3. ………………………………………………………………………………………………………………

63

Activity 12.6

METALS AND EXTRACTION Aim:

To reduce lead(II) oxide by carbon.

Apparatus: Bunsen burner Blow pipe Materials: Lead(II) oxide Carbon block Procedure: 1. Wet the middle of carbon block using tap water. 2. Place a little amount of lead(II) oxide on a carbon block. 3. Heat and blow air using blow pipe over the oxide. 4. Observe any changes on the carbon block.

lead(II) oxide carbon block

blow pipe

Bunsen burner

Observations: …………………………………………………………………………………………………………………... …………………………………………………………………………………………………………………...

64

Activity 13.1

THE PERIODIC TABLE Aim:

To show reactivity of Group I metals with water.

Apparatus: Beaker Scalpel

Pair of forceps White tile

Materials: Lithium Potassium Sodium

Litmus paper Phenolphthalein

Procedure: 1. Put a few drops of phenolphthalein in a 250 cm3 beaker containing water. 2. Cut a very small piece of lithium and drop it in water, and observe. 3. Repeat the above experiment with the following metals: sodium and potassium. 4. Record your observations in the table.

Lithium or sodium or potassium White tile

Results: Metal

Lumps of lithium or sodium or potassium

Effect of solution on red litmus paper

Observations

Lithium

Sodium

Potassium 65

Effect on phenolphthalein indicator

Questions: 1. Place the three metals in the order of decreasing reactivity. Most reactive

Least reactive

…………………………… , …………………………… , …………………………… From the results, it shows that when a metal reacts with water, the solution formed is ……………………………. The compound formed is an ……………………………. The gas given off in the reaction is …………………………… 2. Write a balanced chemical equation for the reaction between lithium and water. ……………………………………………………………………………………………………………… 3. Write a balanced chemical equation for the reaction between sodium and water. ……………………………………………………………………………………………………………… 4. Write a balanced chemical equation for the reaction between potassium and water. ………………………………………………………………………………………………………………

66

Activity 14.1

ENERGY FROM CHEMICALS Aim: To find  H using 0.1 mol dm3 HCl and 0.1 mol dm3 NaOH solutions. Apparatus: Thermometer Plastic cup

Two 100 cm3 measuring cylinders 250 cm3 beakers Materials: 0.1 mol dm3 HCl – solution Q 0.1 mol dm3 NaOH – solution P

Procedure: In this question you are required to determine the heat of neutralisation between a strong acid HY and a strong alkali MOH. You are provided with the following. (a)

Solution P is 0.1 mol dm3 alkali MOH.

(b)

Solution Q is 0.1 mol dm3 acid HY.

1. Using a measuring cylinder measure 100 cm3 of Q into a plastic cup and then measure the temperature. 2. Using another measuring cylinder measure 100 cm3 of P into a 250 cm3 beaker and similarly measure the temperature. 3. Record the temperature of the two solutions and calculate the mean temperature. 4. Slowly pour the solution Q into the beaker containing solution P. 5. Stir the mixture with the thermometer and record the maximum temperature reached.

67

Results: Initial temperature of the acid Q

= …………......

(a)

Initial temperature of the alkali P

= ……………..

( b)

Mean initial temperature [ (a  b)  2 ]

= ……………..

(c )

Maximum final temperature of mixture

= ……………..

(d)

Temperature increase

= ……………..

(e )

[d – c]

(Assume that the density of HCl and NaOH is 1 g cm 3 ) Questions: 1. Write an equation to show the neutralisation between P and Q. ……………………………………………………………………………………………………………… 2. Write an ionic equation for the above reaction. ……………………………………………………………………………………………………………… 3. Ignoring the heat lost to the plastic container and the beaker, calculate the heat given out by the reaction. (The specific heat of solution is 4.2 J g °C). [Hint: Use the formula H = m x c x θ] ……………………………………………………………………………………………………………… ……………………………………………………………………………………………………………… 4. Calculate the number of mole of the acid HY or the alkali MOH involved in the reaction. ……………………………………………………………………………………………………………… ……………………………………………………………………………………………………………… 5. Calculate the heat of neutralisation,  H, between MOH and HY. ……………………………………………………………………………………………………………… ………………………………………………………………………………………………………………

68

Activity 14.2

ENERGY FROM CHEMICALS Aim:

To investigate heat of solution of salts.

Apparatus: Plastic cup Thermometer Spatula Materials: Sodium hydroxide Ammonium chloride Procedure: 1. Measure out 50 cm3 of water in a measuring cylinder and pour it into a plastic cup. 2. Measure the initial temperature of the water. 3. Add a spatula full of solid sodium hydroxide into the water in the cup. 4. Stir to dissolve the solid. 5. Record the highest temperature obtained. 6. Repeat step 1- 4 with ammonium chloride and for step 5 record the lowest temperature obtained. [Repeat with other salts; nitrates, sulphates, carbonates etc.] Result: Reaction

Initial temperature/ C

Final temperature/ C

Sodium hydroxide + water Ammonium chloride + water

Conclusion:

69

Change in temperature/ C

Exothermic or Endothermic

Activity 14.3

ENERGY FROM CHEMICALS Aim:

To set up Daniel cell.

Apparatus: Beaker Voltmeter Materials: Copper plate Zinc plate

1.0 mol dm3 copper(II) sulphate solution Voltmeter Connecting wires

1.0 mol dm3 zinc sulphate solution

Procedure: 1. Set up copper metal as the positive terminal and zinc metal as the negative terminal. 2. Immerse the zinc metal in zinc sulphate solution and the copper metal in copper(II) sulphate solution. 3. Connect the two solutions using a salt bridge as shown. salt bridge

copper plate

zinc plate

copper(II) sulphate solution

zinc sulphate solution voltmeter

Or connect the two solutions using a porous pot as shown below.

zinc

zinc plate

copper plate copper can

porous pot zinc sulphate solution

copper(II) sulphate solution

zinc sulphate solution porous pot

Results: The electrode potential of Daniel cell is ………………. V 70

copper(II) sulphate solution

Activity 15.1

ELECTROLYSIS Aim:

To demonstrate electrolysis of molten lead(II) bromide.

Apparatus: Crucible Spatula Graphite electrodes Power pack

Crocodile clips Tripod stand Clay pipe triangle Bunsen burner

Materials: Lead(II) bromide Procedure: 1. A crucible is half filled with lead(II) bromide solid. 2. The solid lead(II) bromide is heated until it melts to a molten state. 3. Two carbon electrodes are dipped in the molten lead(II) bromide and are then connected to power pack using crocodile clips. 4. Electric current is allowed to flow through for 15 minutes and the changes that occur at the cathode and anode are recorded.

Results: Observations

Inference

At the anode

At the cathode

71

Activity 15.2

ELECTROLYSIS Aim:

To demonstrate electrolysis of dilute sodium chloride solution.

Apparatus: Power pack Carbon electrodes Crocodile clips

100 cm3 beaker

Materials: Aqueous 0.5 mol dm3 sodium chloride solution Procedure: 1. Aqueous sodium chloride is put into a beaker 2. Insert two carbon electrodes into the aqueous sodium chloride and connect them to the power pack. 3. The switch is turned on and electric current is allowed to flow for 15 minutes and observe. 4. Collect the gas and test. 5. Record your observations in the table.

carbon electrodes aqueous sodium chloride

Results: Electrolyte

Dilute sodium chloride solution

Observations

Test for the gas

At the cathode:

Splint test

At the anode:

Splint test

72

Activity 15.3

ELECTROLYSIS Title:

To Demonstrate electrolysis of concentrated sodium chloride solution.

Apparatus: Power pack Carbon electrodes Crocodile clips

100 cm3 beaker

Materials: Concentrated sodium chloride solution (1 mol dm 3 ) Procedure: 1. Concentrated sodium chloride is put into a beaker. 2. Insert two carbon electrodes into the concentrated sodium chloride solution and connect them to the power pack. 3. The switch is turned on and electric current is allowed to flow for 15 minutes and observe. 4. Collect the gas and test. 5. Record your observations in the table.

carbon electrodes concentrated sodium chloride solution

Results: Electrolyte

Observations At the cathode:

Test for the gas Splint test

concentrated sodium chloride solution At the anode:

Litmus test

73

Activity 15.4

ELECTROLYSIS Aim:

To demonstrate electrolysis of copper(II) sulphate using carbon electrodes.

Apparatus: Power pack Carbon electrodes Crocodile clips

100 cm3 beaker

Materials: Aqueous 0.5 mol dm3 copper(II) sulphate solution Procedure: 1. Copper(II) sulphate solution is put into a beaker. 2. Insert two carbon electrodes into the copper(II) sulphate solution and connect them to the power pack. 3. The switch is turned on and electric current is allowed to flow for 15 minutes and observe. 4. Record your observations in the table. Results: Electrolyte

Observations At the cathode:

Copper(II) sulphate solution

At the anode:

Questions: 1. What would you expect the colour of the copper(II) sulphate solution to be if the electrolysis is carried out for a long period? Why? ……………………………………………………………………………………………………………… ……………………………………………………………………………………………………………… 2. Describe a test for the product formed at the anode. ……………………………………………………………………………………………………………… 74

Activity 15.5

ELECTROLYSIS Aim:

To demonstrate electrolysis of copper(II) sulphate using copper electrodes.

Apparatus: Power pack Copper electrodes Crocodile clips

100 cm3 beaker

Materials: Aqueous 1.0 mol dm3 copper(II) sulphate solution Procedure: 1. Copper(II) sulphate solution is put into a beaker. 2. Insert copper electrodes into the copper(II) sulphate solution and connect them to the power pack. 3. The switch is turned on and electric current is allowed to flow for 15 minutes and observe. 4. Record your observations in the table. Results: Electrode

Observations At the cathode:

Copper(II) sulphate solution

At the anode:

Questions: 1. What would you expect the colour of the copper(II) sulphate solution to be if the electrolysis is carried out for a long period? Why? ……………………………………………………………………………………………………………… ………………………………………………………………………………………………………………

75

Activity 15.6

ELECTROLYSIS Aim:

To demonstrate electroplating of spatula with copper.

Apparatus: Power pack Beaker 250 cm Crocodile clips

1.0 mol dm3 copper(II) sulphate Copper plate Metal spatula

3

Materials: 1.0 mol dm3 copper(II) sulphate Procedure: 1. Pour about 200 cm3 of 1.0 mol dm3 copper(II) sulphate solution into a beaker. 2. A piece of copper plate is connected to the positive terminal. This plate act as the anode. 3. The metal spatula is connected to the negative terminal. This metal spatula acts as the cathode. 4. Immerse both the metal spatula and the copper plate in the copper(II) sulphate solution. (Make sure that they do not come into contact) 5. The solution is electrolysed for 30 minutes using a small current (0.5 A ). Results: Observations

Half-equations

At the cathode:

At the cathode:

At the anode:

At the anode:

76

Activity 16.1

SPEED OF REACTIONS Aim:

To show the effect of concentration on the speed of reaction.

Apparatus: Beaker Stop-watch

Filter paper marked “X” Measuring cylinders

Materials: 2.0 mol dm3 hydrochloric acid 0.25 mol dm3 sodium thiosulphate, Na 2 S 2 O 3 solution Distilled water Procedure 1. Measure 5 cm3 of sodium thiosulphate solution into a 250 cm3 beaker and add 45 cm3 of distilled water. 2. Place the beaker over the filter paper marked “X”. 3. By using a separate measuring cylinder, add 10 cm3 of 2.0 mol dm3 HCl and at once start the stop-watch. 4. Swirl the beaker a few times and then put the beaker back on the paper over the mark “X”. 5. Observe the mark “X” from above through the solution mixture in the beaker. 6. As more precipitate is formed the mark “X” will eventually disappears from sight. 7. Stop the stop-watch when the mark “X” just disappears from sight. Record the time taken in the table below. 8. Repeat the experiment by changing the volume of sodium thiosulphate and distilled water. 9. Record the time taken in the table provided.

77

Results: Volume of sodium thiosulphate solution,

Volume of distilled water

Volume of hydrochloric acid

cm3

cm3

cm3

Reaction time s

5 10 15 20 25 Draw the graph of volume of sodium thiosulphate in cm3 against time in seconds. Conclusion: As the concentration of sodium thiosulphate increases ………………………………………………….. ………………………………………………………………………………………………………………...… …………………………………………………………………………………………………………………... …………………………………………………………………………………………………………………...

78

Activity 16.2

SPEED OF REACTIONS Aim:

To show the effect of temperature on the speed of reaction.

Apparatus: Beaker Stop-watch Filter paper marked “X”

Measuring cylinder Thermometer

Materials: 2.0 mol dm3 hydrochloric acid Distilled water 0.25 mol dm3 sodium thiosulphate, Na 2 S 2 O 3 solution Procedure: 1. Measure 20 cm3 of sodium thiosulphate solution into a beaker and add 60 cm3 of distilled water. Record the temperature of the solution mixture. 2. Place the beaker over the filter paper marked “X”. 3. Add 20 cm3 of 2.0 mol dm3 HCl and at once start the stop-watch. 4. Swirl the beaker a few times and put the beaker back on the filter paper over the mark “X”. 5. Observe the mark “X” through the solution mixture. 6. Stop the stop-watch when the mark “X” just disappears from sight. Record the time taken in the table provided. 7. Repeat the experiment at different temperatures by heating the mixtures of 20 cm3 of sodium thiosulphate and 60 cm3 of distilled water to 40°C, 50°C, 60°C and 70°C.

79

Results: Temperature

Reaction time in seconds

Rate ( 1 t )

Room temperature 40°C 50°C 60°C 70°C Draw the graph of rate against temperature. Conclusion: As the temperature increases ……………………………………………………………………………….. ………………………………………………………………………………………………………………...… …………………………………………………………………………………………………………………... …………………………………………………………………………………………………………………...

80

Activity 16.3

SPEED OF REACTIONS Aim:

To show the effect of particle size using calcium carbonate (lump and powder) with hydrochloric acid.

Apparatus: Conical flask connected to gas syringe Stop-clock Electronic balance Measuring cylinder 50 cm3 Materials: Hydrochloric acid 0.25 mol dm3 Calcium carbonate (powder and lump forms) Procedure: 1. Use an electronic balance to weigh exactly 0.5 g of marble chips and put them in a conical flask connected to a gas syringe. 2. Use a measuring cylinder to measure exactly 30.0 cm3 of 0.25 mol dm3 hydrochloric acid and pour into the conical flask containing the marble chips. 3. Immediately cork the conical flask to the gas syringe and at the same time start the stop-clock. 4. Read the volume of the carbon dioxide collected in the gas syringe for every 10 seconds until the reaction stops. 5. Create your own table to tabulate the readings. 6. Repeat the above experiment by using marble chips of (a) smaller size and (b) powdered calcium carbonate. 7. Draw on the same graph paper the volume of carbon dioxide gas against time for the three experiments. Result:

Conclusion: As the particle size increases ……………………………………………………………………………….. ………………………………………………………………………………………………………………...… 81

Activity 16.4

SPEED OF REACTIONS Aim:

To show the decomposition of hydrogen peroxide using manganese(IV) oxide.

Apparatus: Delivery tube Rubber bung Rubber tubing connector Retort stand and clamp

50 cm3 measuring cylinder Clock Spatula Conical flask Gas syringe Materials:

0.2 mol dm3 or ‘2 volume’ hydrogen peroxide (20 volume diluted 10 x) Powdered manganese(IV) oxide Procedure: 1. Set up the apparatus shown in the diagram.

hydrogen peroxide solution manganese(IV) oxide

2. Measure 50 cm3 of hydrogen peroxide solution in the conical flask. 3. Add a little amount of manganese (IV) oxide into the hydrogen peroxide solution. 4. Immediately cork rubber bung to the conical flask and start timing. 5. Gently swirl the flask while recording the volume of gas collected in the gas syringe every minute. Do this for 10 minutes or until the gas syringe is full. Take care not to undo the rubber tube connector. 6. Record your results in table provided.

82

Results: Time/min

0

1

2

3

4

5

6

Volume of gas /cm3 Plot a graph of volume of gas evolved (vertical axis) against time. Conclusion:

83

7

8

9

10

Activity 17.1

REVERSIBLE REACTIONS Aim:

To show reversible reactions.

Apparatus: Test tubes with rack Materials: Potassium chromate(VI) solution Aqueous sodium hydroxide Aqueous copper (II) sulphate Concentrated hydrochloric acid Concentrated ammonium hydroxide Dilute sulphuric acid Procedure: 1. (a) Pour potassium chromate(VI) solution in a test-tube until it is about one-fifth full and slowly adds dilute sulphuric acid until a change is seen. Observation: ……………………………………………………………………………………………………… ……………………………………………………………………………………………………… Write ionic equation for the reaction in (a). ……………………………………………………………………………………………………… (b) To the mixture from (a) slowly add dilute sodium hydroxide solution and observe the change. Observation: ……………………………………………………………………………………………………….. .………………………………………………………………………………………………………. Explain in terms of reversibility the observation in the experiment (b). ………………………………………………………………………………………………………… …………………………………………………………………………………………………………

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2. (a) Pour copper (II) sulphate solution into a test-tube until it is about one-fifth full and slowly add concentrated hydrochloric acid until a change is seen. Observation: …………………….…………………………………………………………………………………. ..……………………………………………………………………………………………………… Write the ionic equation for the reaction in (a). ……………………………………………………………………………………………………….. (b) To the mixture from (a) add concentrated ammonia solution a little at a time until a change is seen. Observation: ……………………………………………………………………………………………………… ……………………………………………………………………………………………………… 3. State Le Chatelier’s principle. …………………………………………………………………………………………………………...... ......................................................................................................................................................

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Activity 17.2

REVERSIBLE REACTIONS Aim:

To prepare fertilizer using nitric acid (The manufacture of fertilizer from ammonia).

Apparatus: Burette Pipette Pipette filler 2 conical flasks Retort stand

Evaporating dish Bunsen burner Tripod stand and wire gauze White tile

Materials: Dilute nitric acid Aqueous ammonium hydroxide Methyl orange indicator Procedure: 1. Place nitric acid in burette. 2. Pipette 25.0 cm3 of aqueous ammonium hydroxide into a conical flask. Add a few drops of the indicator. 3. Titrate the acid with the alkali until neutralise. 4. Repeat titration, this time without the indicator. 5. Heat the mixture (ammonium nitrate solution) until it is saturated. 6. Transfer some mixture into an evaporating dish and leave to crystallise. Conclusion:

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Activity 18.1

REDOX Aim:

To show the colour change in oxidizing agents - acidified potassium manganate(VII) and acidified potassium dichromate(VI) solution.

Apparatus: Measuring cylinder Beaker Conical flask

Test-tube Burette

Material: Potassium manganate(VII) solution Potassium dichromate(VI) solution

Dilute sulphuric acid Sodium sulphite solution

Procedure: 1. Measure 20 cm3 of potassium manganate(VII) solution into a conical flask and add about 5 cm3 of dilute sulphuric acid.

2. From the burette, slowly add sodium sulphite solution (Na2SO3) into the conical flask until the purple colour of potassium manganate(VII) disappears. 3. Repeat the experiment above using potassium dichromate(VI) solution. Observations: (a) Potassium manganate (VII) solution is …………………… (colour) due to the presence of …………………… ions. (b) As sodium sulphite solution is added, the colour changes to …………………… (colour) due to the formation of …………………… ion. (c) Potassium dichromate(VI) solution is …………………… (colour) due to the presence of …………………… ions. (d) As sodium sulphite solution is added the colour changes to …………………… (colour) due to the formation of …………………… ion. Conclusion: Potassium manganate(VII) and potassium dichromate(VI) are oxidizing agents. They are …………………… by the reducing agent, sodium sulphite.

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Activity 18.2

REDOX Aim:

To show the colour change of iodide ion in redox reaction.

Apparatus: Boiling tube Test tube Materials: Hydrogen peroxide ‘20-volume’, 1.0 mol dm3 potassium iodide solution Starch solution Procedure: 1. Pour about 5 cm3 of potassium iodide solution into a boiling tube and slowly add equal volume of ‘20-volume’ hydrogen peroxide, and observe. 2. Leave the mixture for a few minutes and observe for the formation of any solid substance. 3. Transfer some solution from the boiling tube into a test tube and add a few drops of starch solution. Results: (a) Potassium iodide solution changed from ……………………………. (colour) to ……………………………. (colour) when added with hydrogen peroxide. (b) When left to stand for a few minutes a ……………………………. (colour) solid substance was seen at the bottom of the boiling tube. (c) When starch was added, the solution turned ……………………………. (colour) showing that ……………………………. was formed. Questions: 1. Write an ionic equation to show oxidation of iodide ion. ……………………………………………………………………………………………………………… 2. Write an ionic equation for the reduction of hydrogen peroxide. ……………………………………………………………………………………………………………… 3. Write ionic equation to show redox reaction between potassium iodide and hydrogen peroxide. ………………………………………………………………………………………………………………

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Activity 19.2

ATMOSPHERE AND ENVIRONMENT Aim:

To demonstrate water treatment using alum (Potassium aluminium sulphate).

Apparatus: 500 cm3 beakers Spatula Materials: Sample of muddy water Alum Procedure: 1. Fill two beakers with muddy water until three quarter full. 2. Add a spoonful of alum in one but not the other. 3. Leave the two beakers for about 10 minutes and compare the appearance of the water in the two beakers.

alum

muddy water

Observations: …………………………………………………………………………………………………………………... …………………………………………………………………………………………………………………...

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Activity 20.1

ORGANIC CHEMISTRY Aim:

To show incomplete combustion of hydrocarbon.

Apparatus: Bunsen burner Wooden splint

Evaporating dish

Materials: Benzene Toluene Kerosene Procedure: 1. Place a little of each hydrocarbon into an evaporating dish. 2. Burn the hydrocarbon with a lighted wooden splint and observe for the products. (Note: Burning should be carried out outside the lab as its burning will produce dense smoke and carbon monoxide). 3. Observe the colour of the flame and record it in the table below

Lighted wooden splint

Evaporating dish Benzene

Results: Hydrocarbon

Formula

Benzene

C 6H 6

Toluene

C 6H 5 CH 3

Kerosene

C15H 32

Mr

90

% of carbon

Colour of the flame

Questions: 1. Do you think these hydrocarbons make good fuel? Explain. ……………………………………………………………………………………………………………… ……………………………………………………………………………………………………………… ……………………………………………………………………………………………………………… 2. Which of the three hydrocarbons produces the least smoke? Relate your observations in terms of percentage of carbon in the compound. ……………………………………………………………………………………………………………… ……………………………………………………………………………………………………………… ……………………………………………………………………………………………………………… ………………………………………………………………………………………………………………

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Activity 20.3

ORGANIC CHEMISTRY Aim:

To test for alkenes with bromine

Apparatus: Test tube Materials: Liquid bromine Hexene Procedure: 1. Add about 2 cm 3 of liquid bromine in a test tube. 2. Then, add about 2 cm 3 of hexene to the liquid bromine. 3. Shake the mixture gently. 4. Observe the colour changes that take place in the test tube. Results: 1. The brown colour of liquid bromine is ……………………………... 2. State the molecular formula of hexene: ……………………………. 3. Write the equation for the addition of bromine to hexene. ……………………………………………………………………………………………………………… 4. Name the product of addition of bromine to hexene. ………………………………………………………………………………………………………………

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Activity 20.4

ORGANIC CHEMISTRY Aim:

To compare the flammability and the colour of the flame produced by different alcohols and to show the variation of physical properties of the first four alcohols.

Apparatus: Evaporating dishes Wooden splint

Bunsen burner

Materials: Methanol Ethanol Propanol

Butanol Cobalt(II) chloride paper

Procedure: 1. Pour a little of each of the following alcohols (methanol, ethanol, propanol and butanol) into four different evaporating dishes. 2. Burn the alcohols with the lighted wooden splint. 3. Compare the colour of the flame and the ease of it burning. 4. Observe what is left on the evaporating dishes. Test it with blue cobalt(II) chloride paper. Lighted wooden splint

Lighted wooden splint

Evaporating dish

Evaporating dish

Methanol

Ethanol

Lighted wooden splint

Lighted wooden splint

Evaporating dish

Evaporating dish

Propanol

Butanol

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Questions: 1. Note the differences of the colour of the flames from the burning of methanol to butanol. Explain this variation. ……………………………………………………………………………………………………………… ……………………………………………………………………………………………………………… ……………………………………………………………………………………………………………… 2. Which alcohol is the most flammable? ……………………………………………………………………………………………………………… 3. Work out the molecular masses of the four alcohols. Alcohol

Molecular mass

Methanol Ethanol Propanol Butanol

4. How is the flammability of the alcohols related to their molecular masses? ……………………………………………………………………………………………………………… ………………………………………………………………………………………………………………

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Activity 20.5

ORGANIC CHEMISTRY Aim:

To compare fluidity of the alcohols.

Apparatus: Test tubes with rack Materials: Methanol Ethanol

Propanol Butanol

Procedure: 1. Half fill the four test tubes with four different alcohols and compare their fluidity. 2. Place the order of fluidity among the four alcohols. Results: most fluid 1

2

3

least fluid 4

Questions: As the molecular mass increases the alcohols become ………………………………. fluid.

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Activity 20.6

ORGANIC CHEMISTRY Aim:

To show oxidation of ethanol to ethanoic acid.

Apparatus: Bunsen burner Boiling tube Test tube

Retort stand with clamp Beaker Delivery tube

Materials: Ethanol Potassium dichromate(VI) solution Dilute sulphuric acid

Cold water Blue litmus paper

Procedure: 1. The boiling tube is filled with approximately 5 cm3 of potassium dichromate(VI) solution. 2. About 5 cm3 of dilute sulphuric acid is added to the potassium dichromate(VI) solution. 3. About 5 cm3 of ethanol is added to the acidified potassium dichromate(VI) solution. 4. A rubber stopper fitted with a delivery tube is inserted into the boiling tube. The delivery tube is inserted into a test tube placed in a beaker half-filled with cold water. 5. The mixture of ethanol and acidified potassium dichromate(VI) is boiled slowly. The distillate is collected in the test tube. 6. Observe the colour change of acidified potassium dichromate(VI) to show that oxidation of ethanol to ethanoic acid has taken place. 7. The colour and the odour of the distillate are recorded. 8. The distillate is tested with a piece of blue litmus paper.

Boiling tube

Test tube

Mixture of ethanol and acidified potassium dichromate(VI)

Distillate Cold water

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Results: Test on the distillate

Observations

Colour Smell Action on blue litmus paper Questions: 1. When ethanol is boiled with acidified potassium dichromate(VI) solution, it is oxidised to ethanoic acid which has the smell of …………………………….. 2. The colour of acidified potassium dichromate changes from ……………………………. to ………………………………..

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Activity 20.7

ORGANIC CHEMISTRY Aim:

To show acidic properties of carboxylic acid.

Apparatus: Evaporating dish Tripod stand Bunsen burner Spatula

Wooden splints Test tubes Glass rod Delivery tube

Materials: Ethanoic acid Aqueous sodium hydroxide Sodium carbonate powder

Magnesium ribbon Limewater Red litmus paper

Procedure: (a) Reaction of ethanoic acid with sodium hydroxide 1. About 10 cm3 of ethanoic acid is poured into an evaporating dish. 2. Drop a piece of red litmus paper into the evaporating dish to act as indicator. 3. Slowly add sodium hydroxide solution to the ethanoic acid until the red litmus paper just turns blue. 4. Remove the litmus paper and heat the solution mixture until dryness.

Observations: Describe the product left behind the evaporating dish after the solution mixture is evaporated to dryness. …………………………………………………………………………………………………………………... .............................................................................................................................................................

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(b) Reaction between ethanoic acid with magnesium ribbon 1. The test tube is filled with about 5 cm3 of ethanoic acid. 2. A piece of magnesium ribbon (about 3 cm long) is added to ethanoic acid. 3. Test the gas given off with the lighted wooden splint. Observations: …………………………………………………………………………………………………………………...

(c) Reaction between ethanoic acid with sodium carbonate 1. About 5 cm3 of ethanoic acid is added to a test tube. 2. A spatula of sodium carbonate powder is added to the ethanoic acid. 3. The gas released is passed into limewater. Observations: …………………………………………………………………………………………………………………...

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REFERENCES: 1. T.Y. Toon, L.W. Leng & T.O Tin. Success Chemistry SPM. Selangor Darul Ehsan: Oxford Fajar Sdn. Bhd, 2007. 2. Practical Chemsitry For ‘O’ Level. Singapore: Dyna Publisher Pte. Ltd. 3. C.N. Prescott. Chemistry A Course for ‘O’ Level. Practical Workbook. Volume 1. Singapore: Federal Publications (S) Pte Ltd, 1994. 4. C.N. Prescott. Chemistry A Course for ‘O’ Level. Practical Workbook. Volume 2. Singapore: Federal Publications (S) Pte Ltd, 1994. 5. L.J. Rasanayagam. Practical Chemistry. A course for ‘O’ Level. Volume 1. Singapore: Federal Publications (S) Pte Ltd, 1981. 6. L.J. Rasanayagam. Practical Chemistry. A course for ‘O’ Level. Volume 2. Singapore: Federal Publications (S) Pte Ltd, 1981. 7. J.G.R. Briggs. Chemistry ‘O’ Level Practical. Volume 1 (2nd Ed.). Singapore: Pearson Education Asia Pte Ltd, 1998. 8. J.G.R. Briggs. Chemistry ‘O’ Level Practical. Volume 2 (2nd Ed.). Singapore: Pearson Education Asia Pte Ltd, 1998. 9. T.Y. Toon & C.L. Kwong. Chemistry Matters for GCE ‘O’ Level. Practical Workbook. Singapore: Times Media Private Limited, 2002. 10. L.J. Rasanayagam & R.M. Kok. GCE ‘O’ Level Chemistry Matters. Practical Book. Singapore: Marshall Cavendish Education, 2007. 11. R.M. Heyworth & J.G.R. Briggs. Chemistry Insights ‘O’ Level (2nd Ed.) Practical Workbook. Singapore: Pearson Education South Asia Pte Ltd, 2007.

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