Exptans Bk2(e)

Chemistry A Modern View Experiment Workbook 2　Suggested answers Contents PART V ACIDS AND ALKALIS Chapter 15

Acids

15.1 To investigate properties of dilute acids

3

15.2 To study the role of water in exhibiting characteristic properties of acid

6

15.3 To investigate the corrosive nature of concentrated acids

Chapter 16

10

Alkalis

16.1 To investigate the action of dilute alkalis on ammonium compounds

11

16.2 To investigate the action of dilute alkalis on aqueous metal ions to form metal hydroxide precipitates

12

16.3 Action of concentrated sodium hydroxide solution on meat (T)

Chapter 17

Indicators and pH

17.1 To find pH values of some common substances

Chapter 18

14

15

Strength of acids and alkalis (Extension)

18.1 To compare the relative strength of acids and of alkalis

17

Chapter 19 Neutralization and salts 19.1 To investigate the temperature change associated with a neutralization reaction

22

19.2 To prepare sodium sulphate crystals from an acid-alkali titration

Chapter 21 Simple volumetric work (Extension) ©Aristo Educational Press Ltd. 2003

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23

21.1 To prepare standard ethanedioic acid solutions

24

21.2 To find the molarity of a given hydrochloric acid

25

21.3 To determine the concentration of ethanoic acid in commercial vinegar

26

Chapter 22 Rate of reaction (Extension) 22.1 To investigate the effect of concentration on the rate of reaction

27

22.2 To investigate the effect of surface area on rate of reaction

28

22.3 Effect of temperature on the rate of reaction

33

PART VI CHEMICAL CELLS AND ELECTROLYSIS Chapter 23

Chemical cells in daily life

23.1 To compare the service life of a zinc-carbon cell with an alkaline manganese cell of the same size (T)

Chapter 24

38

Simple chemical cells

24.1 To make simple chemical cells and construct part of the Electrochemical Series

Chapter 26

39

Cell reactions (Extension)

26.1 To construct a simple chemical cell using inert electrodes

41

26.2 To examine and compare the internal structures of a well-used and a new zinccarbon cell

Chapter 27

42

Electrolysis (Extension)

27.1 Electrolysis of dilute sulphuric acid (T)

43

27.2 Effect of concentration on preferential discharge of ions

44

27.3 Effect of electrodes on products of electrolysis

47

27.4 Electroplating

48

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PART VII PRODUCTS FROM IMPORTANT PROCESSES Chapter 28

Chlorine and hypochlorite

28.1 To investigate the oxidizing property of chlorine water

51

28.2 To make chlorine bleach

52

28.3 Action of chlorine bleach on coloured substances

56

28.4 Action of acid on chlorine bleach

57

Chapter 29

Sulphuric acid and sulphur dioxide

29.1 To investigate properties of concentrated sulphuric acid (S/T)

58

29.2 To dilute concentrated sulphuric acid (S/T)

60

29.3 To prepare sulphur dioxide and test for its properties

61

29.4 To bleach coloured papers and flower petals with sulphur dioxide (S/T)

65

MICROSCALE EXPERIMENT M1

Electrolysis using a microscale Hoffman apparatus

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Experiment 15.1

To investigate properties of dilute acids

Test

Observations 2 M hydrochloric acid

1.

■ Test

2.

■ (a)

with a blue litmus paper.

Clean a 2 cm length of magnesium ribbon with sandpaper. Add the cleaned ribbon to the acid. Quickly place the test tube upright in a test tube rack. Immediately cover the mouth of the tube with an inverted

2 M ethanoic acid

turns red

turns red

rapid effervescence of a colourless gas; the magnesium dissolves quickly to give a colourless solution

effervescence of a colourless gas; the magnesium dissolves to give a colourless solution

a ‘pop’ sound is heard

a ‘pop’ sound is heard

hydrogen

hydrogen

rapid effervescence of a colourless gas; the calcium granules dissolve quickly to give a colourless solution

effervescence of a colourless gas; the calcium granules dissolve to give a colourless solution

a ‘pop’ sound is heard

a ‘pop’ sound is heard

hydrogen

hydrogen

rubber stopper (Figure 15.1). Leave the tube to stand, until the magnesium has dissolved completely. ■ (b)

Remove the stopper and quickly put a burning splint into the mouth of the tube (Figure 15.2).

(c) Name the gas evolved. 3.

■ (a)

Repeat Step 2(a), using 2 calcium granules instead of a magnesium ribbon. (Caution! Handle calcium with forceps, never with bare fingers.)

■ (b)

Test any gas evolved with a burning splint.

(c) Name the gas evolved. (d) Write a general word equation for the reactions involved in Steps 2(a) and 3(a).

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acid + metal → salt + hydrogen

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Test

Observations 2 M hydrochloric acid

4.

■ (a) Add

1/4 spatula measure of

2 M ethanoic acid

the black copper(II)

the black copper(II)

copper(II) oxide to the acid. Warm

oxide solid dissolves to

oxide solid dissolves to

the mixture gently (Figure 15.3).

give a blue solution

give a blue solution

(Caution! Wear safety spectacles.) (b) Write a general word equation for

acid + metal oxide → salt + water

the reactions involved here. 5.

the calcium hydroxide

the calcium hydroxide

calcium hydroxide to the acid.

dissolves to give a

dissolves to give a

Warm the mixture gently.

colourless solution

colourless solution

■

(a) Add 1/4 spatula measure of

(Caution! Wear safety spectacles.) (b) Write a general word equation for

acid + metal hydroxide → salt + water

the reactions involved here. 6.

■ (a) Add

3 spatula measures of

rapid

effervescence of a

anhydrous sodium carbonate to the

effervescence of a

colourless gas; the

acid. Pass any gas evolved through

colourless gas; the

carbonate dissolves to

limewater (Figure 15.4).

carbonate dissolves

form a colourless

quickly to form a

solution; limewater

colourless solution;

turns milky

limewater turns milky (b) Name the gas evolved.

carbon dioxide

carbon dioxide

(c) Write a general word equation for

acid + carbonate → salt + carbon dioxide + water

the reactions involved here.

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7.

■ (a)

Repeat Step 6(a), using sodium

rapid

effervescence of a

hydrogencarbonate instead of

effervescence of a

colourless gas; the

sodium carbonate. Pass any gas

colourless gas; the

hydrogencarbonate

evolved through limewater.

hydrogencarbonate

dissolves to form a

dissolves quickly to

colourless solution;

form a colourless

limewater turns milky

solution; limewater turns milky (b) Name the gas evolved.

carbon dioxide

(c) Write a general word equation for

acid + hydrogencarbonate

carbon dioxide

the reactions involved here. → salt + carbon dioxide + water

8.

Hydrogen gas. Hydrogen ion, H+(aq). Yes.

9.

c.

No.

d.

No. Hydrochloric acid reacts with the carbonate to give a salt, carbon dioxide and water. The H+(aq) ions are removed as water. Thus the characteristic acidic properties are lost.

e.

10. a.

Na2CO3(s) + 2HCl(aq) → 2NaCl(aq) + CO2(g) + H2O(l)

red

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salt; hydrogen salt; water salt; carbon dioxide; water b.

hydrogen ions H+(aq)

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Experiment 15.2 sample laboratory report Title:

To study the role of water in exhibiting characteristic properties of acid

Purpose To study the role of water in exhibiting characteristic properties of acid.

Apparatus and chemicals used • • • • •

Watch glass (dry) Dry test tubes (150 × 18 mm) in rack Stopper that fits 24 mm diameter test tube Matches/lighter Blue litmus paper (dry)

• • • •

Magnesium ribbon (2 cm length) Citric acid crystals (dry), 2 g Deionized water Wooden splint

Chemical reactions involved Mg(s) + 2H+(aq) 2H2(g) + O2(g)

Mg2+(aq) + H2(g) 2H2O(l)

('pop' sound test)

Procedure (A)

To compare the action of solid citric acid and its aqueous solution on dry blue litmus paper

1.

(a) (b) (c)

Half a spatula measure of dry solid citric acid was added to a dry watch glass. A dry blue litmus paper was dipped into the solid acid (Figure 1a) Any colour change of the dry blue litmus paper was recorded.

2.

(a) (b) (c)

1 cm3 of water was added to the solid citric acid. A dry blue litmus paper was dipped into the aqueous solution (Figure 1b). Any colour change of the dry blue litmus paper was recorded.

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(B)

To compare the action of solid citric acid and its aqueous solution on magnesium

3.

(a) (b) (c)

A piece of magnesium ribbon was put into a dry test tube (Figure 2a). 3 spatula measures of solid citric acid were added to the test tube (Figure 2b). Any observation occurred was recorded

4.

(a) (b)

Water was added to the test tube (from Step 3) to a depth of 3 cm. The tube was shaken to dissolve the citric acid crystals. It was quickly placed upright in a test tube rack. The mouth of the tube was immediately covered with an inverted rubber stopper. The tube was allowed to stand for 5 minutes (Figure 3a). Any observation occurred was recorded.

(c)

(d) (e)

After 5 minutes, the stopper was removed and a burning splint was quickly put into the mouth of the tube (Figure 3b). Any observation of the experiment was recorded.

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Observation 1. 2. 3. 4.

(Reference to Step 1): No observable change. (Reference to Step 2): The blue litmus paper changed to red. (Reference to Step 3): No observable change. (Reference to Step 4): The citric acid gradually dissolved to form a colourless solution. At the same time, colourless gas bubbles were evolved from the surface of the magnesium ribbon. The magnesium ribbon gradually became smaller. The bottom of the test tube became warm. The gas gave out a 'pop' sound with a burning splint, which was identified as hydrogen.

Interpretation 1. 2. 3. 4.

No observable change because there is no hydrogen ions in the absence of water. Thus, solid citric acid does not show acidic property. After the addition of water, hydrogen ions that can turn the blue litmus paper red, are formed. No observable change because there is no hydrogen ions in the absence of water. Thus, solid citric acid does not show any acidic property. Citric acid ionizes in water to give H+(aq) ions. It reacts with magnesium to give hydrogen gas. Mg(s) + 2H+(aq) Mg2+(aq) + H2(g) It shows acidic property.

Discussion 1. 2.

3. 4.

It is essential that the apparatus and chemicals used must be dry at the start, so that the effect of adding water can be compared correctly. It is not a good practice to cover the test tube with student 's thumb because there may be some acid on the rim of the test tube. If a right-sized stopper is used to cover the test tube, do not leave it unattended when there is effervescence in the test tube. Pressure will build up inside the tube and the stopper may shoot out. This will accidentally hurt students' eyes. The stopper used should have a bigger size (suitable for test tube with 24 mm diameter). Unreacted magnesium should never be disposed of into the sink. It should be collected by the teacher and the laboratory technician will take care of it. Magnesium metal surface should be cleaned with sand paper to remove any oxide layer together with the grease. It should be noted that holding the metal by hand would make the metal greasy.

Conclusion Without water, an acid does not show the usual acidic properties. Water must be present for an acid ©Aristo Educational Press Ltd. 2003

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to give H+(aq) ions, which are responsible for the typical acidic properties.

Answers to questions for further thought 1. It changes blue litmus paper to red colour. It conducts electricity. It reacts with metal oxides and hydroxides to give salts and water. It reacts with carbonates and hydrogencarbonates to give salts, carbon dioxide and water. 2. (a) When it is dissolved in water, the solid acid ionizes to form H+(aq) ions, which react with sodium hydrogencarbonate. There is effervescence, carbon dioxide gas being given off. (b) H+(aq) + HCO3−(aq) H2O(l) + CO2(g) (c) It should be stored in a dry cool place. 3. In the presence of water, acidic gas will dissolve and ionize to form H+(aq) ions. Thus, it can change the blue litmus paper to red colour. 4. (a) No colour change. It is because pure ethanoic acid liquid contains no water. (b) From blue to red. In the presence of water, ethanoic acid ionizes to form H+(aq) ions. Thus, it can show acidic properties.

©Aristo Educational Press Ltd. 2003

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Experiment 15.3

To investigate the corrosive nature of concentrated acids

1.

d.

2.

Yes.

A ‘pop’ sound is heard. Hydrogen.

Concentrated hydrochloric acid. Zn(s) + 2HCl(aq) → ZnCl2(aq) + H2(g) 3.

c.

No signs of reaction no matter the acid is cold or hot. Yes. A dilute aqueous solution of a typical acid does not react with copper and those metals below copper in the metal reactivity series.

4.

Brown fumes are evolved quickly to fill the tube. 1; 4; 1; 2; 2 No. An aqueous solution of a typical acid reacts with those metals, which are above copper in the reactivity series. In such reactions, the gas liberated is hydrogen.

5.

a.

zinc; the same; faster; acidic; acidity; corrosive; acidity

b.

no; oxidizes; very corrosive; oxidizing

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Experiment 16.1

To investigate the action of dilute alkalis on ammonium compounds

4.

A characteristic pungent choking smell of ammonia. Ammonia.

5.

b.

It turns from red to blue. (NH4)2SO4(aq) + 2NaOH(aq) → Na2SO4(aq) + 2NH3(g) + 2H2O(l)

6.

Ammonia; ammonium sulphate; sodium hydroxide; ammonia; pungent, choking; red; blue

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Experiment 16.2

To investigate the action of dilute alkalis on aqueous metal ions to form metal hydroxide precipitates

Test

Observations 0.5 M NaOH(aq)

1.

Add a piece of red litmus paper to 2 cm3 of an alkali solution in a test tube (Figure 16.4).

2.

(a)

Using a teat pipette, add 10 drops of sodium hydroxide solution to 6 cm3 of water in a test tube (Figure 16.5). Rub a little of the diluted solution between your fingers.

turns blue

it has a slippery (soapy) feel

1 M NH3 (aq) turns blue

it has a slippery (soapy) feel

(Caution! (1) Wash your hand immediately afterwards with plenty of water. (2) Omit this step if there is a wound in your fingers. This is to avoid possible irritation.) (b)

Rub a little of the 1 M ammonia solution (undiluted) between your fingers.

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6. Metal ions

Addition of a little NaOH(aq)

Addition of excess NaOH(aq)

Addition of a little NH3(aq)

Colour of precipitate formed (if any)

Does the precipitate dissolve in excess NaOH(aq)?

Colour of precipitate

Pb2+(aq)

white

yes

white

Cu2+(aq)

blue

no

blue

Fe2+(aq)

dirty green

no

dirty green

Fe3+(aq)

reddish brown

no

reddish brown

K+(aq)

(no precipitate)

(not applicable)

(no precipitate)

white

yes

white

Zn2+(aq)

7.

formed (if any)

Yes. Hydroxide ion, OH-(aq). NaOH(s) + H2O(l) → Na+(aq) + OH−(aq) NH3(aq) + H2O(l)

8.

NH4+(aq) + OH−(aq)

a.

Pb2+(aq), Cu2+(aq), Fe2+(aq), Fe3 +(aq), Zn2+(aq)

b.

Pb2+(aq) + 2OH−(aq) → Pb(OH)2(s); Cu2+(aq) + 2OH−(aq) → Cu(OH)2(s); Fe2+(aq) + 2OH−(aq) → Fe(OH)2(s); Fe3+(aq) + 3OH−(aq) → Fe(OH)3(s); Zn2+(aq) + 2OH−(aq) → Zn(OH)2(s)

9.

Pb2+(aq), Zn2+(aq)

10. a.

red; blue

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soapy metal hydroxides Pb(OH)2(s), Cu(OH)2(s), Fe(OH)2(s), Fe(OH)3(s), Zn(OH)2(s) b.

hydroxide ion OH−(aq)

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Experiment 16.3

Action of concentrated sodium hydroxide solution on meat (T)

1.

Pale pink or slightly yellow (answer varies). Yes.

5.

The chicken foot becomes smaller. It turns reddish yellow and somewhat ‘translucent’. The solution becomes pale red. The chicken foot becomes white. There is no other apparent change.

6.

The chicken foot crumbles (breaks up) when stirred. A mixture of skin, meat and white oily mass floats. Small pieces of bones sink to the bottom.

7.

There is no apparent change.

8.

Concentrated sodium hydroxide solution is very corrosive on meat.

9.

He should wash the affected part with plenty of water (for at least a few minutes), then report to the teacher.

10. corrosive; eat; chicken feet

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Experiment 17.1

To find pH values of some common substances

4.

(1) (2) (3) (4) (5) (6) (7) (8) (9) (10) (11) (12)

4 10 2 3 9 10 7 answer varies 1 14 12 7

5.

(a) distilled water, sodium chloride solution (b) soft drink (7-up), vinegar, lemon juice, 1 M hydrochloric acid (c) soap solution, Philips Milk of Magnesia/ window cleaner, limewater, 1 M sodium hydroxide

12. pH scale

0

1

2

3

4

5

6

7

8

9

10

11

12

Indicator

Methyl orange

red

orange

Litmus

red

Phenolphthalein

Colourless

yellow purple

blue very pale pink

red

13. Red / yellow. Not sure. Acidic / not sure. 14. Red / blue. Not sure. Acidic / alkaline. ©Aristo Educational Press Ltd. 2003

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13

14

15. Colourless / red. Not sure. Not sure / alkaline. 16. No, we could only get a rough idea of the pH, acidity or alkalinity of a solution. 17. By using universal indicator or pH meter.

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Experiment 18.1 sample laboratory report Title: To compare the relative strength of acids and of alkalis Purpose To compare the relative strength of (A) acids and (B) alkalis.

Apparatus and chemicals used Approach 1: (A)

• • • • • • •

To compare the relative strength of acids and of alkalis by measuring electrical conductivities

For comparing the relative strength of acids

Beaker (100 cm3) • Hydrochloric acid (0.1 M) Power supply (0 − 24 V) / 6 V d.c. supply • Ethanoic acid (0.1 M) Electrode rod holder 2 carbon rods 3 connecting wires, with crocodile clips at both ends Digital multimeter (or milliammeter) Measuring cylinder (100 cm3)

(B)

For comparing the relative strength of alkalis

• Same as in (A) • Ammonia solution (0.1 M)

Approach 2: (A)

• • • •

To compare the relative strength of acids and of alkalis by measuring pH values

For comparing the relative strength of acids

2 test tubes (150 × 18 mm) in rack Glass rod Measuring cylinder (100 cm3) pH paper (pH range 1 − 14) with colour chart

(B)

• Sodium hydroxide solution (0.1 M)

• Hydrochloric acid (0.1 M) • Ethanoic acid (0.1 M)

For comparing the relative strength of alkalis

• Same as in (A) • Ammonia solution (0.1 M)

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• Sodium hydroxide solution (0.1 M)

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Chemical reactions involved HCl(aq) H+(aq) 1 Cl−(aq) CH3COOH(aq) CH3COO−(aq) + H+(aq) NaOH(s) + water Na+(aq) + OH−(aq) NH3(aq) + H2O(l) NH4+(aq) + OH−(aq)

Approach 1:

To compare the relative strength of acids and of alkalis by measuring electrical conductivities

Procedure (A)

To compare the relative strength of acids by measuring electrical conductivities.

1. 2.

80 cm3 of 0.1 M hydrochloric acid was put into a small beaker (100 cm3). The apparatus was set-up as shown in Figure 1.

1

3.

2

5. 6.

The range selector of the multimeter was adjusted to a suitable range for measuring d.c. current. The reading was noted immediately. The electrode rod holder was taken out of the beaker. The carbon rods were washed with running tap water. The beaker was emptied and washed well with water. Steps 1 to 5 were repeated. 80 cm3 of 0.1 M ethanoic acid was used instead.

(B)

To compare the relative strength of alkalis by measuring electrical conductivities.

7. 8.

Steps 1 to 5 were repeated. 80 cm3 of 0.1 M sodium hydroxide solution was used instead. Steps 1 to 5 were repeated. 80 cm3 of 0.1 M ammonia solution was used instead.

4.

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Approach 2:

To compare the relative strength of acids and of alkalis by measuring pH values

(A)

To compare the relative strength of acids by measuring pH values.

9.

(a)

A glass rod was dipped into a little 0.1 M hydrochloric acid in a test tube to get a drop of the solution (Figure 2a).

11.

(b) The drop of acid was applied to a piece of pH paper (Figure 2b). The colour of the pH paper was matched with the colour chart. The pH value was then recorded. Steps 9 to 10 were repeated using 0.1 M ethanoic acid instead.

(B)

To compare the relative strength of alkalis by measuring pH values.

12. 13.

Steps 9 to 10 were repeated using 0.1 M sodium hydroxide solution instead. Steps 9 to 10 were repeated using 0.1 M ammonia solution instead.Observation

10.

Approach 1: 1.

2.

To compare the relative strength of acids and of alkalis by measuring electrical conductivities (Reference to Steps 1 to 6): Equal concentration of hydrochloric acid (0.1 M) and ethanoic acid (0.1 M) show different electrical conductivities. Hydrochloric acid shows a higher electrical conductivity than that of ethanoic acid. (Reference to Steps 7 and 8): Equal concentration of sodium hydroxide solution (0.1 M) and ammonia solution (0.1 M) show different electrical conductivities. Sodium hydroxide solution shows a higher electrical conductivity than that of ammonia solution.

Approach 2:

To compare the relative strength of acids and of alkalis by measuring pH

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1.

2.

values (Reference to Steps 9 to 11): Equal concentration of hydrochloric acid (0.1 M) and ethanoic acid (0.1 M) show different pH values. Hydrochloric acid shows a lower pH value than that of ethanoic acid. (Reference to Steps 12 and 13): Equal concentration of sodium hydroxide solution (0.1 M) and ammonia solution (0.1 M) show different pH values. Sodium hydroxide solution shows a higher pH value than that of ammonia solution.

Interpretation Approach 1: 1.

2.

To compare the relative strength of acids and of alkalis by measuring electrical conductivities Equal concentration of hydrochloric acid and ethanoic acid means that they have the same concentration of solute in the solution. However, hydrochloric acid is a strong acid, which will ionize completely in aqueous solution. Ethanoic acid is a weak acid which will only slightly/partially/incompletely ionize in aqueous solution. As a result, hydrochloric acid will have a higher ionic concentration and conduct electricity better. HCl(aq) H+(aq) + Cl−(aq) (complete ionization) + CH3COOH(aq) CH3COO2(aq) + H (aq) (slight ionization) Equal concentration of sodium hydroxide solution and ammonia solution means that they have the same concentration of solute in the solution. However, sodium hydroxide solution is a strong alkali which will dissociate completely in aqueous solution. Ammonia solution is a weak alkali which will only slightly/partially/incompletely ionize in aqueous solution. As a result, sodium hydroxide solution will have a higher ionic concentration and conduct electricity better. NaOH(s) + water Na+(aq) + OH−(aq) (complete dissociation) − + NH3(aq) + H2O(l) NH4 (aq) + OH (aq) (slight ionization)

Approach 2: 3.

4.

To compare the relative strength of acids and of alkalis by measuring pH values Equal concentration of hydrochloric acid and ethanoic acid means that they have the same concentration of solute in the solution. However, hydrochloric acid is a strong acid, which will ionize, completely in aqueous solution. Ethanoic acid is a weak acid which will only slightly/partially/ incompletely ionize in aqueous solution. As a result, hydrochloric acid will have a higher concentration of hydrogen ions, and thus a lower pH value. HCl(aq) H−(aq) + Cl−(aq) (complete ionization) − + CH3COOH(aq) CH3COO (aq) + H (aq) (slight ionization) Equal concentration of sodium hydroxide solution and ammonia solution means that they have the same concentration of solute in the solution. However, sodium hydroxide solution is

©Aristo Educational Press Ltd. 2003

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a strong alkali which will dissociate completely in aqueous solution. Ammonia solution is a weak alkali which will only slightly/partially/incompletely ionize in aqueous solution. As a result, sodium hydroxide solution will have a higher concentration of hydroxide ions, and thus a higher pH value. NaOH(s) + water Na+(aq) + OH−(aq) (complete dissociation) NH3(aq) + H2O(l) NH4+(aq) + OH−(aq) (slight ionization)

Discussion 1.

2. 3. 4.

For a fair comparison, the strong acid and weak acid used should be of equal concentration. The difference in electrical conductivities is thus a result of the difference in the degree of ionization of the two acids in aqueous solution. The digital multimeter is an ideal instrument for measuring electric current here. The readings are quite accurate and can be easily read on the LCD screen. Always remember to connect the negative pole of the power supply to the negative (black) terminal of the multimeter. Electrolysis of solution will take place, liberating gases at electrodes. The gas bubbles on the surface of electrodes increase the resistance of the circuit, causing the current reading to drop with time.

Conclusion 1. 2.

Hydrochloric acid is a stronger acid than ethanoic acid of the same concentrations? Sodium hydroxide solution is a stronger alkali than ammonia solution of the same concentrations?

Answers to questions for further thought 1. 2.

3.

Citric acid, ethanoic acid, ascorbic acid, etc. No, it is because for a fair comparison, all factors such as temperature and concentration of solute in aqueous solution should be the same, except that different acids or alkalis are used. So when we are measuring the electrical conductivities, the difference is a direct reflection of the degree of ionization (or dissociation) of the acids or alkalis. Thus, the strength of acids and of alkalis can be compared. (a) water, ammonia (the most plentiful), ammonium ions, hydroxide ions, hydrogen ions (b) water, sodium ions, hydroxide ions (the most plentiful), hydrogen ions

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Experiment 19.1

To investigate the temperature change associated with a neutralization reaction

4.

Given out. Exothermic

7.

Roughly the same. H+(aq); OH−(aq); Twice; given out; heat up; rises; roughly the same

11. The temperature rise is about half of that in Part B. 12. Exothermic. 13. Yes. 14. very little heat; bad; heat; heat; conduction; heat losses; bad; heat

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Experiment 19.2

To prepare sodium sulphate crystals from an acid-alkali titration

5.

Yellow.

9.

To obtain the salt free from any indicator.

10.

b.

12.

Colourless crystals.

If all the water were driven away, the sodium sulphate obtained would be a powder, not crystals.

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Experiment 21.1 4.

To prepare standard ethanedioic acid solutions

9.51 (12.0 + 16.0 × 2 + 1.0) × 2 + 2 × (1.0 × 2 + 16.0) 126.0 250.0 9.51/126.0 250.0/1000 0.302

11. 0.302 0.302

25.0 250.0

0.0302

Yes. This is because its concentration (molarity) is accurately known.

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Experiment 21.2 5.

Yellow.

6.

b.

9.

0.102 26.1

To find the molarity of a given hydrochloric acid

The end point has been reached.

10. NaOH(aq) + HCl(aq) → NaCl(aq) + H2O(l) 0.102 0.102

25.0 2.55 × 10−3 1000

1:1 26.1; 2.55 × 10−3 2.55 × 10−3 26.1/1000

0.0977

©Aristo Educational Press Ltd. 2003

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Experiment 21.3

To determine the concentration of ethanoic acid in commercial vinegar

2.

c.

6.

0.102

Colourless.

25.2 7.

CH3COOH(aq) + NaOH(aq) → CH3COONa(aq) + H2O(l) 25.2; 0.102 25.2 2.57 × 10−3 1000

0.102

1:1 2.57 × 10−3 a.

2.57 × 10−3 25.0/1000

b. 0.103

0.103

250.0 1.03 25.0

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Experiment 22.1

To investigate the effect of concentration on the rate of reaction

8.

The curve is steep at first as the reaction rate is the fastest, but becomes less steep with time. Finally, it becomes horizontal, indicating the finish of the reaction. Quick and accurate results can be obtained by using data-logger. Besides, continuous monitoring is possible. However, using data-logger is more complicated than traditional methods and it is less easy to be handled.

9.

increase; increase

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Experiment 22.2 sample laboratory report Title:

To investigate the effect of surface area on rate of reaction

Purpose To investigate the effect of surface area on rate of reaction by using data-logger.

Apparatus and chemicals used • • • • • • • •

Suction flask (250 cm3) Dry lumps of calcium carbonate, 0.5 g • 3) Measuring cylinder (25 cm Dry powdered calcium carbonate, 0.5 g • Plastic sample bottle, large enough to hold • Dilute HCl(aq) (1 M) 3 about 25 cm of 1 M hydrochloric acid Stand, boss and clamp Data-logger connected to a computer with pre-installed software, absolute pressure sensor Rubber tubing (8 inches) for connecting the suction flask and the data logger Electronic balance (accurate to 0.01 g) Vacuum sealant and adhesive

Chemical reactions involved CaCO3(s) + 2HCl(aq)

CaCl2(aq) + CO2(g) + H2O(l)

Procedure 1.

0.5 g of lumps of calcium carbonate was weighed out and put inside a dry suction flask. (Figure 1)

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2.

25 cm3 of 1 M hydrochloric acid was measured by a measuring cylinder and transferred into a plastic sample bottle (Figure 2).

3. 4.

The plastic sample bottle with hydrochloric acid was put into the suction flask carefully. The suction flask was stoppered carefully and sealed to avoid air leakage. A piece of rubber tubing was used to connect the data-logger and the suction flask for pressure measurement. The data logging software was started and set properly for data collection. The suction flask was inverted to mix the reactants thoroughly and the data collection process was started immediately. See Figures 3 and 4.

5. 6.

pressure sensor

data-logger

to computer

Figure 4

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7. 8. 9.

When the reaction finished, the data collection process was stopped. The results (data and graph) were saved and printed out to hand in together with the report. Steps 1 − 8 were repeated, but 0.5 g of powdered calcium carbonate was used instead.

Observation 1. 2.

(Reference to Step 6): Effervescence occurred in the reaction mixture. (Reference to Step 6): The pressure inside the suction flask increased steadily and very slowly with lumps of calcium carbonate added (more than 2.5 minutes) but very rapidly (less than 0.5 minute) with powdered calcium carbonate added. 3. The results were shown below: With lumps of calcium carbonate: (Specimen results)

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With powdered calcium carbonate: (Specimen results) Pressure vs Time

3

Time (s)

Interpretation 1. 2.

Effervescence occurred in the reaction mixtures, due to the formation of carbon dioxide gas. CaCO3(s) + 2HCl(aq) CaCl2(aq) + CO2(g) + H2O(l) As more and more carbon dioxide was produced, the pressure inside the suction flask also increased.

Discussion 1.

2. 3.

In case the calcium carbonate lumps are covered with powder, they should be washed beforehand in dilute HCl and then in water to remove the powder. They should then be blotted. Since particle size of calcium carbonate solids would change as reaction proceeds, only the initial rates of the two experiments should be compared. Theoretically, the two graphs should become flat at the same value on the Y-axis, i.e. same final pressure. But in practice, there may be some deviations as a result of experimental errors

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(e.g. masses used not the same). Actually, the final pressure cannot be reached as it takes quite a long time for the calcium carbonate lumps to react completely.

Conclusion The reaction rate increases with larger surface area of solid reactant.

Answers to questions for further thought 1. 2.

3.

Hydrochloric acid HCl. The hydrochloric acid should be in excess so that the concentration of it does not change a lot. Reaction rate was the greatest at the start and decreased gradually to zero. At the start, the concentration of the reactants, were the highest (hence fastest reaction rate). As the reaction proceeded, the concentrations of the reactants decreased (hence slower reaction rate), until one of the reactants (calcium carbonate CaCO3 in this case) was used up. Reaction rate became zero as the reaction stopped. Chewing can break the food into small pieces, so that there can be a large increase in surface area of the food for the digestive juices to work on, and thus facilitates digestion.

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Experiment 22.3 sample laboratory report Title:

Effect of temperature on the rate of reaction

Purpose To investigate the effect of temperature on the rate of reaction.

Apparatus and chemicals used • • • • • • • • • • •

Beakers (100 cm3) Measuring cylinder (50 cm3) Measuring cylinder (10 cm3) Stopwatch White tile Black/blue 'Wytebord' markers Bunsen burner Tripod and gauze Heat-resistant mat Thermometer (−0  110 °C, stirring rod type) Tissue paper or blotting paper (to wipe beakers dry)

• Dilute HCl(aq) (2.0 M) (about 40 cm3) • Na2S2O3 solution (~0.05 M) (about 80 cm3) • Deionized water

Chemical reaction involved S2O32−(aq) + 2H+(aq)

SO2(g) + S(s) + H2O(l)

Procedure 1. 2.

A thick cross with the size just smaller than the base of a small beaker was marked with a black/blue 'Wytebord' marker on a white tile placed on the bench. (a) 5 cm3 of sodium thiosulphate solution was mixed with 45 cm3 of water in a 50 cm3 measuring cylinder. (b) The mixture was added to the small beaker placed on the marked white tile. (c) 5.0 cm3 of hydrochloric acid was poured into the beaker quickly. The stopwatch was started at the same time (Figure 1a). The mixture was stirred gently with the thermometer and the temperature of the reaction mixture was taken (Figure 1b).

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(d)

3.

The cross was observed by looking vertically down through the solution. The time taken for the cross to be just 'blotted out' completely was recorded (Figure 1c).

Another 5 cm3 of thiosulphate solution was mixed with 45 cm3 of water in a 50 cm3 measuring cylinder. The mixture was added to another clean small beaker. The beaker was heated until the temperature of the reaction mixture was just above 35°C (Figure 2a).

(a)

The beaker was placed on the marked white tile. 5.0 cm3 of hydrochloric acid was poured into the beaker quickly. The stopwatch was started at the same time. The mixture was stirred gently with the thermometer and the temperature of the mixture was taken once it became

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(b)

The time taken for the cross to be just 'blotted out' completely was recorded.

4.

Step 3 was repeated at about 45°C, 55°C, 60°C, 65°C and 70°C.

5.

The results obtained in the experiment were tabulated and a graph of time against

1

temperature was plotted.

Observation 1.

2.

3.

At the start of the experiment, the cross was seen clearly. As reaction proceeded, sulphur concentration continued to rise. The cross became fainter but could still be seen. After a certain time, sulphur concentration became just sufficient to 'blot out' the cross completely. The results were tabulated as below: Temperature of reaction mixture (°C)

Time for cross to be just 'blotted out' (s)

22 35 45 55 60 65 69

154.0 63.5 41.3 27.4 23.0 18.0 14.9

1 Time

(Specimen results)

was worked out in each case:

Temperature of reaction mixture (°C)

Time for cross to be just 'blotted out' (s)

1/Time (s−1)

22 35 45 55 60 65 69

154.0 63.5 41.3 27.4 23.0 18.0 14.9

6.5 × 10−3 15.7 × 10−3 24.2 × 10−3 36.5 × 10−3 43.5 × 10−3 55.6 × 10−3 67.1 × 10−3

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1

A graph of Time against temperature was plotted below:

70

60

50

40

30

20

10

0

10

20

30

40

50

60

70

Interpretation 1.

Sodium thiosulphate solution and hydrochloric acid reacted according to the equation: S2O32−(aq) + 2H+(aq) SO2(g) + S(s) + H2O(l) As sulphur formed, the mixture became cloudy white, and then cloudy yellow. The degree of

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2.

cloudiness of the reaction mixture at any moment indicated the extent of the reaction. In this experiment, the beaker containing the reaction mixture was placed over a cross marked on a piece of white tile. The cross was observed directly from above. The time (t) taken for the cross just to be 'blotted out' was recorded. This corresponded to the time needed for the formation of a certain definite amount of sulphur. As temperature of the reactants increased, the time to 'blot out' the cross became shorter. Since reaction rate is inversely proportional to the time taken to 'blot out' the cross, a shorter time means a faster reaction. From the graph, it could be seen that reaction rate increased with an increase in temperature.

Discussion 1.

2.

3. 4.

For a more accurate comparison, the same cross mark must be used throughout the experiment. So it is necessary to keep the cross dry and clean. The same person is to observe the cross at the same distance from above the cross. He should also be the timer. Small amount of SO2 with choking smell is produced during the experiment. Although the gas is poisonous, it does not matter much at such low concentration in a well-ventilated laboratory. The beaker inside should be cleaned and dried, otherwise, it will dilute the solution in the beaker. When heating Na2S2O3 solution in the beaker, the thermometer may be used as a stirring rod (provided that care is taken). It is advisable to place the thermometer in the warm thiosulphate solution before pouring the acid in. The temperature will only drop by 1 to several degrees (depending on the starting temperature of the thiosulphate solution), and thus the thermometer can quickly adjust to register the temperature after mixing. This practice is particularly advantageous for experiments carried out at high temperatures (e.g. 70°C)  the reaction rate is so fast that we cannot wait too long for the mercury thread of thermometer to go up from room temperature to the final high temperature.

Conclusion Reaction rate increases with increasing temperature of the reaction mixture.

Answers to questions for further thought 1.

2.

No. If the reaction rate is directly proportional to temperature, the shape of the graph should be a straight line, which is not the case in this experiment. From the graph, it seems that reaction rate rises exponentially with temperature. It is a rule of thumb that for every 10°C rise in temperature, reaction rate will be doubled. To investigate the effect of a certain factor (temperature in this case) on reaction rate, all other factors should be made the same at the start.

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3.

Fresh food can be kept fresh for a longer time in a refrigerator. Low temperatures would slow down reactions that deteriorate the food.

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Experiment 23.1

To compare the service life of a zinc-carbon cell with an alkaline manganese cell of the same size (T)

2.

c

3.

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4.

(a) 13.5 52.5 (b) Alkaline manganese cell. About 4 times. (c) Alkaline manganese cell.

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Experiment 24.1

To make simple chemical cells and construct part of the Electrochemical Series

5. Metal couple in cell

Voltage of cell (volt)

Mg/Cu

+1.41

Fe/Cu

+0.20

Zn/Cu

+0.79

Cu/Cu

0

Ag/Cu

−0.22

8. Metal couple in cell

Voltage of cell (volt)

Mg/Cu

+1.41

Fe/Cu

+0.20

Zn/Cu

+0.83

Cu/Cu

0

Ag/Cu

−0.23

11. Metal couple in cell

Voltage of cell (volt)

Mg/Cu

+1.84

Fe/Cu

+0.49

Zn/Cu

+0.92

Cu/Cu

0

Ag/Cu

−0.17

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12. a.

Mg/Cu cell, Fe/Cu cell, Zn/Cu cell.

b. From the other metal to copper. c. The other metal. d. Mg, Zn, Fe, Cu. 13. a. Ag/Cu cell. b. From copper to silver. c. Copper. 14. Mg, Zn, Fe, Cu, Ag. 15.

a.

Yes.

b.

Metals react by losing electrons. The higher the tendency of a metal to form ions (the higher its position in the E.C.S.), the more reactive it would be (the higher its position in the metal reactivity series).

16. electrodes; electrolyte; positive ions; greater; voltage

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Experiment 26.1

To construct a simple chemical cell using inert electrodes

2.

b.

0.98

3.

b.

Yes. Yellow. Iodine.

c.

2I−(aq) → I2(aq) + 2e− Oxidation.

4.

a.

MnO4−(aq) + 8H+(aq) + 5e− → Mn2+(aq) + 4H2O(l) Reduction.

b.

2MnO4−(aq) + 16H+(aq) + 10I−(aq) → 2Mn2+(aq) + 5I2(aq) + 8H2O(l)

c.

Y.

d.

From X to Y. X. Y. Cathode.

5.

Zero volt. To complete the circuit by allowing ions to move between the two half-cells.

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Experiment 26.2

To examine and compare the internal structures of a well-used and a new zinc-carbon cell

1.

2.

a.

The zinc cup of the used cell is thinner than that of the new cell. This shows that some zinc metal has dissolved in the used cell.

b.

A strong smell of ammonia is detected in the used cell, but not in the new cell. This shows that some ammonia has formed in the used cell.

c.

In the used cell, a large area around the graphite rod looks wet. In the new cell, the ammonium chloride paste looks only very slightly wet. This shows that some liquid (probably water) has formed in the used cell.

a.

Zn(s) → Zn2+(aq) + 2e−

b.

2NH4+(aq) + 2e− → 2NH3(g) + H2(g) 2MnO2(s) + H2(g) → Mn2O3(s) + H2O(l)

3.

(1) The metal casing makes the cell leakproof, preventing the electrolyte from leaking out through the worn-out zinc cup. (Any electrolyte leaking out would damage the electrical appliance in which the cell is put.) (2)

The cell with a metal casing has a longer life.

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Experiment 27.1 3.

b.

Electrolysis of dilute sulphuric acid (T)

Small colourless gas bubbles are evolved continuously. Very small colourless gas bubbles are evolved continuously.

4.

c.

2 : 1 (approx.)

6.

The glowing splint is relighted. Oxygen. 4OH−(aq) → O2(g) + 2H2O(l) + 4e−

7.

A ‘pop’ sound is heard. Hydrogen. 2H+(aq) + 2e− → H2(g)

8.

2H2O(l) → 2H2(g) + O2(g) Water. No. Distilled water is a very poor conductor of electricity. Although sulphuric acid itself is not electrolysed, it can increase the electrical conductivity of water.

9.

hydrogen H+(aq); hydrogen gas; hydroxide OH−(aq); oxygen gas; decrease; remains unchanged; increases

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Experiment 27.2

Effect of concentration on preferential discharge of ions

4.

At cathode Colourless gas bubbles are evolved. The solution around the cathode turns purple (in a few seconds). At anode Colourless gas bubbles are evolved. The solution around the anode turns red (in a few seconds).

5.

c. A 'pop' sound is heard. The gas is hydrogen.

8.

At cathode Colourless gas bubbles are evolved. The solution around the cathode turns purple (in a few seconds). The solution becomes bleached after some time. At anode

9.

Colourless gas bubbles are evolved. The solution around the anode turns red (in a few seconds). The solution becomes bleached within a short time (less than half a minute).

A 'pop' sound is heard. Hydrogen.

10. Bleaching solution or 'swimming pool'. Chlorine.

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11.

0.1 M NaCl

2 M NaCl

Na+(aq), H+(aq)

Na+(aq), H+(aq)

alkaline

alkaline

hydrogen

hydrogen

2H+(aq) + 2e− → H2(g)

2H+(aq) + 2e− → H2(g)

Cl−(aq), OH−(aq)

Cl−(aq), OH−(aq)

C

Cations present

A T H O D E

Is the solution around the electrode acidic, neutral or alkaline?

A

Anions present

N O D E

Is the solution around the electrode acidic, neutral or alkaline?

acidic

acidic

Main product

oxygen

chlorine

4OH−(aq) → O2(g) + 2H2O(l) + 4e−

2Cl−(aq) → Cl2(g) + 2e−

Main product Ionic half-equation

Ionic half-equation 12.

hydrogen H+(aq); hydrogen gas; hydroxide OH−(aq); oxygen gas; chloride Cl−(aq) ; concentration effect; chlorine gas

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15.

C

Cations present

A T H O D E

Is the solution around the electrode acidic, neutral or alkaline?

A

Anions present

N O D E

Colour of solution around the electrode

2 M NaI

Na+(aq), H+(aq)

Na+(aq), H+(aq)

alkaline

alkaline

hydrogen

hydrogen

2H+(aq) + 2e− → H2(g)

2H+(aq) + 2e− → H2(g)

Br−(aq), OH−(aq)

I−(aq), OH−(aq)

brown colour

deep brown colour

Main product Ionic half-equation

(bleached after some time) Main product Ionic half-equation

16.

2 M NaBr

bromine

iodine

2Br−(aq) → Br2(aq) + 2e−

2I−(aq) → I2(aq) + 2e−

hydrogen

hydroxide OH−(aq); reducing; halide; oxygen gas; halide; halide ions; hydroxide OH−(aq); halogen

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Experiment 27.3 3.

Effect of electrodes on products of electrolysis

A reddish brown solid is deposited. Cu2+(aq) + 2e− → Cu(s) Colourless gas bubbles are evolved. 4OH−(aq) → O2(g) + 2H2O(l) + 4e−

4.

The cathode is now electrode Y (graphite); the anode is electrode X (copper coated on graphite).

5.

Reddish brown copper is deposited.

Cu2+(aq) + 2e− → Cu(s) The reddish brown solid (copper) dissolves gradually. (Some copper may fall to the bottom of the beaker.) When there is little or no copper on the anode, colourless gas bubbles are liberated there. At first, Cu(s) → Cu2+(aq) + 2e−; when there is little or no copper left, 4OH−(aq) → O2(g) + 2H2O(l) + 4e−. 6.

copper; oxygen; copper dissolves

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Experiment 27.4 sample laboratory report Title: Electroplating Purpose To design and perform an experiment to electroplate a metal object with nickel.

Apparatus and chemicals used • 6 V d.c. supply • 3 connecting wires with crocodile clips at both ends • 6 V bulb in holder • Sand paper (5 cm × 5 cm), 2 pieces • Beaker (100 cm3) • Crucible tongs • Electrode foil holder • Forceps (placed beside propanone, kept in the fume cupboard)

• Propanone (2 cm3), kept inside the fume cupboard • Nickel foil (5 cm × 2 cm) • Copper foil (5 cm × 2 cm) /brass key/ 50¢ coin • Nickel plating solution (an aqueous nickel(II) salt solution), 60 cm3 • Tissue paper/ cotton wool

Chemical reactions involved At the nickel anode: At the cathode (object to be electroplated):

Ni(s) Ni2+(aq) + 2e− Ni2+(aq) + 2e− Ni(s)

Procedure 1.

(a)

A copper foil was cleaned with sandpaper (Figure 1a).

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(b) (c)

2.

(d) (e) (a)

(b)

(c)

The foil was washed in running tap water (Figure 1b). The foil was degreased (grease was removed) by rubbing the foil with a piece of cotton wool (soaked in propanone) held by forceps (Figure 1c). The foil was washed again under tap water (Figure 1d). The copper foil was put on a piece of tissue paper (Figure 1e) to keep the foil clean. 60 cm3 of the nickel plating solution was poured into a beaker.

The circuit was connected as shown in Figure 2. The copper was made the cathode (connected to the negative terminal of the battery); nickel was made the anode (connected to the positive terminal). A current was allowed to flow for 10 minutes. When half the time had passed, the cathode was taken out. The sides were reversed and it was put back into the plating

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3.

(a) (b) (c) (d)

solution. (This was to ensure even plating.) The cathode was removed (the copper foil was newly coated with nickel). It was washed in running tap water. It was allowed to dry on a tissue paper. The appearance of the electroplated copper was noted.

Observation 1. 2. 3.

(Reference to Step 1): The copper foil looked shiny and clean. (Reference to Step 2): The electroplated copper became silvery grey in appearance. (Reference to Step 3): A silvery grey coating was electroplated on the copper foil.

Interpretation 1. 2.

Propanone is a very good solvent to dissolve grease, thus the copper metal looks shiny. At the cathode, Ni 2+(aq) ions are discharged to form Ni(s) on the copper surface. Ni2+(aq) + 2e− Ni(s)

Discussion 1.

2. 3.

4.

Alternatively, the copper foil can be degreased by dipping it in concentrated detergent solution for a minute, or rubbing with cotton wool soaked in dilute ammonia solution. Warm dilute sodium hydroxide can also remove grease. However, the solution is corrosive. Stirring of the plating solution during electroplating may result in a good coating. This can be done by using the magnetic stirrer. The light bulb indicated that an electric current was flowing and limited this to a small current density suitable for electroplating. If the bulb is not used, a variable resistor and an ammeter can be used to obtain good results. The metal foils can be conveniently clamped by an electrode foil holder which consists of a plastic strip fitted with two crocodile clips. The advantages are: (a) The metal foils will not touch each other and the circuit will not be shorted. (b) The foils can be kept at a fixed distance apart. (c) The crocodile clips will not be wetted by the solution so easily. So that the clips will not rust so easily.

Conclusion When an object is to be electroplated, it is made the cathode of an electrolytic cell. The plating metal is usually made the anode. The plating solution (electrolyte) is a solution of one of the salts ©Aristo Educational Press Ltd. 2003

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of the plating metal. Simple electroplating experiment can be done in the school laboratory conveniently. Answers to questions for further thought 1. (i) The object to be plated should be clean and free from grease. (ii) Keep the electric current small. (iii) The electrolyte should be maintained at a certain constant acidity or alkalinity. 2. Copper, silver, gold, tin or chromium. They are low in the Electrochemical Series. 3. These objects are first sprayed with a layer of powdered graphite or metal and then electroplated in the usual way. 4. The nickel anode ionizes to form Ni2+(aq) ions, replacing those removed from the solution at the cathode. Thus the concentration of the electrolyte can be kept constant at a desired level.

©Aristo Educational Press Ltd. 2003

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Experiment 28.1

To investigate the oxidizing property of chlorine water

1.

d. Halide

(a) Colour of solution after addition of chlorine water

(b) Colour of heptane layer

Bromide

yellow

red-orange

Iodide

brown

purple

4.

oxidizing; bromine; yellow; iodine; brown; heptane; red-orange; purple

5.

Cl2(aq) + 2Br−(aq) → 2Cl−(aq) + Br2(aq); Cl2(aq) + 2I−(aq) → 2Cl−(aq) + I2(aq) Redox (or displacement).

6.

oxidizing

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Experiment 28.2 sample laboratory report Title: To make chlorine bleach Purpose To make chlorine bleach in the school laboratory.

Apparatus and chemicals used • • • • •

Electrolysis cell Rubber bands Rimless test tubes Stopper to fit small test tube Wooden piece (longer than diameter of wide glass tube) • Connecting wires fitted with crocodile clips • 6 V battery

• • • • •

Stand, boss and clamp Dropper Boiling tube (150 × 24 mm) in rack Stopper to fit 24 mm diameter test tube Red litmus paper • Blue litmus paper • Brine (saturated NaCl solution)

Chemical reactions involved At the cathode: 2H+(aq) + 2e− H2(g) − At the anode: 2Cl (aq) Cl2(g) + 2e− Cl2(g) + 2NaOH(aq) NaOCl(aq) + NaCl(aq) + H2O(l) sodium hypochlorite

Procedure In this experiment, brine (saturated solution of sodium chloride) was electrolysed. The chlorine gas collected from the anode was allowed to react with the resulting alkaline solution. 1. Brine was added to an electrolysis cell, until the liquid level was slightly above the electrodes. See Figure 1.

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2.

(a)

(b)

3.

4.

A dropper was used to fill two small rimless test tubes with brine (Figure 2). The tubes were inverted over the electrodes (Figure 3). (Before the tubes were inverted, an empty beaker should be placed under the electrolysis cell. This would prevent possible spillage of liquid on the bench.)

(c) The electrolysis cell was clamped. (d) The electrodes were connected to a 6 V d.c. source. Any observation around the electrodes and in the small test tubes A and B was recorded. (a) Electrolysis was allowed to continue. When the small

©Aristo Educational Press Ltd. 2003

tubeA

Figure 3

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5. 6.

tube above the anode (tube B) was full of gas, the power supply was disconnected. (b) The small tubes were taken out of the solution and tube B was quickly stoppered. (c) The rubber bands were loosen to release tube B. A piece of red litmus paper was used to test the resulting solution in the electrolysis cell. The change in colour of the litmus paper was recorded. (a) The resulting solution from the electrolysis cell was poured into a boiling tube to a depth of 3 cm. (b) Tube B was inverted. The stopper was quickly removed and tube B was dropped into the boiling tube (Figure 4a).

(Caution! It must be very careful when tube B was dropped into the boiling tube. It is afraid that the glass might be broken and the solution might be spilled out.) (c) (d) (e) 7.

The boiling tube was stoppered tightly immediately (Figure 4b). The boiling tube was inverted gently and the contents were swirled (Figure 4c) The boiling tube was inverted back and it was swirled again. Steps (d) and (e) were repeated for about 5 minutes. A chlorine bleach had been prepared. It should contain sodium hypochlorite and show bleaching action. (a) A piece of blue litmus paper and a piece of red litmus paper were dropped into the solution in the boiling tube. (b) The changes in colours of the blue and the red litmus papers were recorded.

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Observation 1.

2. 3.

(Reference to Step 3): Colourless gas bubbles were evolved at cathode. The gas was collected and water level in tube A dropped. Greenish-yellow gas bubbles were evolved at anode. The gas was collected and water level in tube B dropped. (Reference to Step 5): The red litmus paper turned blue. (Reference to Step 7): The blue litmus paper was bleached (from blue to white) immediately. The red litmus paper was bleached (from red to white) immediately.

Interpretation 1.

2. 3.

Hydrogen was formed at the cathode and collected in tube A. 2H+(aq) + 2e− H2(g) Chlorine was formed at the anode and collected in tube B. 2Cl−(aq) Cl2(g) + 2e− The resulting solution was alkaline. It should contain sodium hydroxide. The chlorine bleach formed bleached the dyes in the red and blue litmus paper. Cl2(g) + 2NaOH(aq) NaOCl(aq) + NaCl(aq) + H2O(l) OCl−(aq) + dye(aq) Cl2(aq) + (dye + O) coloured

colourlesss

Discussion 1. 2. 3.

The small tubes should not cover the electrodes completely, otherwise electrical resistance would be increased and the rate of electrolysis would be reduced considerably. The time needed to collect a full tube of hydrogen gas is shorter than that of chlorine. This is because hydrogen is insoluble in water, while chlorine is fairly soluble. The resulting solution is mainly a mixture of sodium hydroxide and sodium chloride, with a little chlorine dissolved. It is alkaline. A longer electrolysis results in a higher alkalinity and stronger corrosiveness of the resulting solution.

Conclusion 1. 2.

3.

Chlorine can be prepared by the electrolysis of brine. Chlorine gas formed at the anode can react with the resulting sodium hydroxide solution in the electrolysis cell. Chlorine bleach, containing sodium hypochlorite as the active ingredient, is formed. The chlorine bleach can bleach the dyes in the litmus paper.

Answers to questions for further thought 1.

The theoretical volume ratio of H2 to Cl2 should be 1:1. Smaller volume of chlorine gas is collected because chlorine is fairly soluble in water.

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2.

At room temperature, chlorine can be liquefied under pressure. It is stored and transported in liquid state in steel cylinders. 3. (i) Chlorine bleach should be stored away from sunlight or excess heat. It is because sunlight and excess heat can speed up the decomposition of chlorine bleach. (ii) Chlorine bleach should be kept out of reach of children because it is toxic. (iii) When chlorine bleach is used, the room must be well ventilated because a little chlorine gas is given off, which is toxic. (iv) When using chlorine bleach, wear plastic gloves because it is irritant to skin. (v) Never mix chlorine bleach with acidic substances or other cleaners. It is because acidic substances react with chlorine bleach to liberate the toxic chlorine gas. − OCl (aq) + Cl− (aq) + 2H+(aq) Cl2(g) + H2O(l)

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Experiment 28.3

Action of chlorine bleach on coloured substances

1.

Sodium hypochlorite (usually 5.25%).

3.

A faint choking smell (like that of swimming pool).

4.

b.

It turns white (is bleached) immediately. It turns white (is bleached) immediately.

5.

i.

The colour becomes paler.

ii.

The stain is bleached.

iii. The stain is bleached. iv. 6.

The stain is bleached.

sodium hypochlorite; oxidation

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Experiment 28.4 1.

Action of acid on chlorine bleach

a.

Colourless.

b.

The solution changes from colourless to pale greenish yellow, a stream of small bubbles being evolved.

2.

The gas has a characteristic pungent, choking smell (like that of swimming pool). Chlorine.

3.

b. It becomes pink first (in about 5 seconds) and then white, i.e. bleached (in about 10 seconds). Cl−(aq); OCl−(aq); 2; Cl2(g)

5.

mineral acid; smell; blue litmus paper

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Experiment 29.1

To investigate properties of concentrated sulphuric acid (S/T)

1.

c.

It turns from blue granular crystals to a pale blue powder. conc. CuSO4‧5H2O(l)

2.

H2SO4

CuSO4(s) + 5H2O(l)

b.

The sugar turns brown. A little white steam is given out.

d.

The sugar chars (turns black). A lot of white steam is given out. A black spongy mass rises up the beaker. Sugar charcoal (carbon). 12C(s) + 11H2O(g)

3.

d.

All the words appear (black words on white paper). On heating, water is driven away from dilute sulphuric acid. The dilute sulphuric acid becomes concentrated. The concentrated sulphuric acid formed dehydrates the cellulose in paper, leaving black carbon. Dehydrating property.

4.

c.

No signs of reaction no matter the acid is cold or hot. Yes. A dilute aqueous solution of a typical acid does not react with copper and those metals below copper in the metal reactivity series.

5.

a.

No.

c.

It changes from light orange to green. Yes.

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e.

Blue. Cu2+(aq). Cu(s) + 2H2SO4(l) → CuSO4(s) + SO2(g) + 2H2O(l) No. A typical acid reacts with those metals which are above copper in the reactivity series. In such cases, the gas liberated is hydrogen. Redox reaction. An oxidizing agent. Concentrated sulphuric acid is reduced, as the oxidation number of S decreases from +6 in H2SO4(l) to +4 in SO2(g).

6.

dehydrating; oxidizing

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Experiment 29.2 7.

To dilute concentrated sulphuric acid (S/T)

lighter; heat; heating; boiling; steam; acid spray; heat; break

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Experiment 29.3 sample laboratory report Title:

To prepare sulphur dioxide and test for its properties

Purpose (A) To prepare sulphur dioxide gas in the school laboratory. (B) To investigate the properties of sulphur dioxide gas.

Apparatus and chemicals used • • • •

Anti-bumping granules Blue litmus paper/pH paper Dilute sulphuric acid Sodium sulphite solid • Potassium dichromate solution • Boiling tube (fitted with a rubber stopper carrying a bent delivery tube)

• • • •

5 test tubes (4 fitted with rubber stoppers) Test tube rack Bunsen burner and matches Heat-resistant mat • Stand, boss and clamp • Beaker (250 cm3)

Chemical reactions involved SO32−(aq) + 2H+(aq) SO2(g) + H2O(l) SO2(aq) + H2O(l) H2SO3(aq) sulphurous acid

3SO2(aq) + Cr2O72−(aq) + 2H+(aq)

3SO42−(aq) + 2Cr3+(aq) + H2O(l)

Procedure (A)

To prepare sulphur dioxide gas in the school laboratory

1.

(a) 3 spatula measures of sodium sulphite solid was put in a boiling tube. (b) Dilute sulphuric acid was added to the boiling tube to a depth of 3 cm. (c) A few anti-bumping granules were added to the boiling tube. The boiling tube was clamped as shown in Figure 1.

2.

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3.

4.

5.

(a)

The reaction mixture in the boiling tube was warmed gently. The flame was moved about continuously to ensure uniform heating (b) Record the observation when the reaction mixture was heated. (a) The gas liberated was collected in a test tube by download delivery (upward displacement of air). The test tube was stoppered once it was full of the gas. (As the gas generated was misty, one could see when a tube was full of the gas.) The stoppered test tube was put in a test tube rack. (b) 3 more tubes of sulphur dioxide gas were collected. The tubes of gas would be used for Part B of the experiment. Heating was stopped. With a towel to protect the hand, the boiling tube was taken to the fume cupboard, to be cleaned later after cooling.

(B)

To investigate the properties of sulphur dioxide gas

6.

Test for smell (a) A tube of sulphur dioxide gas was taken. The stopper was lifted up slightly to leave a small opening. (b) The gas that escaped out from the tube was smelled carefully. This was done by 'fanning' a little of the gas towards the nose. (c) The stopper was put in place immediately. (d) The smell of sulphur dioxide gas was recorded. Test for water solubility and acidic property (a) A beaker was filled with water to half-full. (b) A tube of sulphur dioxide was inverted and immersed under water in the beaker (Figure 2).

7.

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(c) (d) (e) (f)

8.

9.

The stopper was removed under water. The water was stirred gently with the inverted tube for 4 minutes. A piece of blue litmus paper was dropped into the solution in the beaker. Any change in water level inside the tube and the colour change of the litmus paper were recorded.

Test for bleaching property (a) Inside the fume cupboard, the stopper of a tube of sulphur dioxide was removed. A piece of moist blue litmus paper was quickly put into the tube. (b) The stopper was immediately put in place again. (c) The stoppered tube was taken back to the students' bench. (d) The tube was placed in the test tube rack and allowed to stand for 10 minutes. (e) Any colour change of the blue litmus paper was recorded. Test for reducing property (a) 1 cm3 of potassium dichromate solution and 1 cm3 of dilute sulphuric acid were put in a test tube. (The resulting solution is called 'acidified potassium dichromate solution'.) (b) Inside the fume cupboard, the acidified potassium dichromate solution was added quickly to a tube of sulphur dioxide gas. (c) The stopper of the tube was immediately replaced and the tube was shaken a few times. (d) Any colour change of the acidified potassium dichromate solution was recorded.

Observation 1. 2. 3. 4. 5.

(Reference to Step 3): Effervescence occurred in the reaction mixture. A misty gas was formed. (Reference to Step 6): Sulphur dioxide gas had an irritating choking smell of burning sulphur. (Reference to Step 7): Water level rose well up inside the test tube. The blue litmus paper turned red. (Reference to Step 8): The blue litmus paper turned red and then white (or very pale red). (Reference to Step 9): The orange colour of acidified potassium dichromate solution changed

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to dark green.

Interpretation 1.

2.

Effervescence occurred in the reaction mixture, due to sulphur dioxide gas formed. SO32−(aq) + 2H+(aq) SO2(g) + H2O(l) The sulphur dioxide gas liberated was misty because it contained moisture in it. Water level rose well up inside the test tube. This shows that a lot of sulphur dioxide had dissolved in water, hence the gas is very soluble in water. The litmus paper turned red, showing that the solution was acidic. Sulphur dioxide dissolves in water to form sulphurous acid: SO2(aq) + H2O(l) H2SO3(aq) sulphurous acid

3.

4.

The blue litmus paper turned red first. This shows that sulphur dioxide is an acidic gas. The litmus paper then turned white (or very pale red), showing that sulphur dioxide is a bleaching agent. The orange colour of acidified potassium dichromate solution changed to dark green, showing that sulphur dioxide gas is a reducing agent. The orange dichromate ions were reduced to dark green chromium(III) ions: 3SO2 (aq) + Cr2O72−(aq) + 2H+(aq) SO42−(aq) + 2Cr3+(aq) + H2O(l) orange

dark green

Discussion 1. 2. 3.

Sulphur dioxide gas is much denser than air. Thus it can be collected by downward delivery, as in this experiment. However, the gas collected would have air mixed with it. Inevitably, some sulphur dioxide escaped into the laboratory in this experiment. It would be better if the whole experiment had been carried out inside the fume cupboard. The sulphur dioxide gas collected in this experiment (by downward delivery) was moist and mixed with air. To prepare pure dry sulphur dioxide, the gas generated should be first passed into concentrated sulphuric acid for drying; the dried gas can then be collected using a gas syringe. Using a gas syringe can also minimize the sulphur dioxide gas escaped into the laboratory.

Conclusion 1. 2. 3. 4. 5.

Sulphur dioxide can be prepared by warming a sulphite with a dilute acid. Sulphur dioxide has an irritating choking smell of burning sulphur. Sulphur dioxide is acidic. Sulphur dioxide is a bleaching agent. Sulphur dioxide is a reducing agent. It turns acidified potassium dichromate solution from

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orange to dark green.

Answers to questions for further thought 1. 2. 3.

Sulphur dioxide cannot be collected by displacement of water, as it is very soluble in water. It is not advisable to test for a gas by its smell because the gas may be poisonous. A chemical test would be preferred. Sulphur dioxide emitted into the atmosphere is formed by the burning of sulphur-containing fuels. Almost all of the gas comes from industrial sources ¾ electric power stations, factories and incinerators. Sulphur dioxide dissolves in rainwater to form sulphurous acid, one of the causes of acid rain: SO2(aq) + H2O(l) H2SO3(aq) sulphurous acid

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Experiment 29.4

To bleach coloured papers and flower petals with sulphur dioxide (S/T)

1.

A colourless gas (or a misty gas, due to presence of impurity).

3.

d.

The blue litmus paper becomes white (or very pale red). The red litmus paper becomes white (or very pale red). The blue (or red) flower petals become paler in colour.

4.

Yes. Reduction.

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Experiment M1

Electrolysis using a microscale Hoffman apparatus

11. Small colourless bubbles are evolved continuously. Very small colourless gas bubbles are evolved continuously 12. Test the gas evolved at the anode with a glowing splint, the glowing splint is relighted, indicating that the gas is oxygen. Test the gas evolved at the cathode with a burning splint, the burning splint gives a 'pop' sound, indicating that the gas is hydrogen.

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