Exptans Bk1(e)
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Chemistry A Modern View Experiment Workbook 1 Suggested answers Contents PART I
INTRODUCING CHEMISTRY
Chapter 1 1.1
What is Chemistry?
Observation in chemistry
Chapter 2
3
The fundamentals of chemistry
2.1
To observe Brownian movement of smoke particles (T/S)
4
2.2
To prepare a compound by direct combination of elements and to compare properties of the compound and its constituent elements (S/T)
PART II
PLANET EARTH
Chapter 3 3.1
5
The atmosphere
Test for oxygen
Chapter 4
7
Oceans
4.1
Extraction of common salt from sea water
8
4.2
Isolation of pure water from sea water (T)
9
4.3
Tests to show the presence of sodium and chloride ions in common salt
10
4.4
To show the presence of water in a given sample
12
Chapter 5
Rocks and minerals
5.1
Action of heat, water and acids on calcium carbonate
13
5.2
Chemical tests for the presence of calcium carbonate
14
PART III THE MICROSCOPIC WORLD Chapter 6
Atomic structure
6.1
To inspect samples of some common substances
17
6.2
To find which elements conduct electricity
19
©Aristo Educational Press Ltd. 2003
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Chapter 7 7.1
The Periodic Table
To investigate which elements show similar chemical properties (T/S)
Chapter 8
21
Chemical bonding: Ionic bonding
8.1
To find out which compounds conduct electricity
23
8.2
Effect of electricity on molten lead(II) bromide (T)
25
8.3
Colours of ions
26
8.4
Migration of coloured ions (S/T)
28
8.5
To build a lattice model of sodium chloride
29
Chapter 9 9.1
Chemical bonding: Covalent bonding
To build models of diamond and quartz
Chapter 10
30
Structures and properties
10.1 To compare the properties of sodium chloride, dry ice, wax and quartz
PART IV
31
METALS
Chapter 11
Occurrence and extraction of metals
11.1 To extract metals from metal oxides (T/S)
Chapter 12
32
Reactivity of metals
12.1 To arrange five metals in order of reactivity
33
12.2 Displacement reactions of metals in aqueous solutions
35
Chapter 13
Reacting masses
13.1 To determine the empirical formula of magnesium oxide (T/S) (Extension) 36
Chapter 14
Corrosion of metals and their protection
14.1 To investigate factors that influence rusting
37
14.2 To prevent rusting
40
©Aristo Educational Press Ltd. 2003
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Experiment 1.1
Observation in chemistry
1.
The alcohol burns with a pale blue flame having a yellow tip. The alcohol soon dries and the burning also stops. There is no blackening of/no soot left on the watch glass.
2.
b.
The potassium dichromate solution changes from orange to dark green.
3.
b. d.
A blue precipitate is formed. The precipitate dissolves to form a deep blue solution.
4.
There is no visible change, but the test tube becomes warm.
5.
c.
The white solid dissolves to form a colourless solution. The test tube becomes cool.
6.
d.
There is no visible change, but an irritating smell of ammonia can be detected.
7.
c.
The clear solution gradually turns cloudy white and then cloudy yellow. A choking smell can be detected.
8.
c.
There is effervescence — small colourless gas bubbles are evolved from the magnesium surface, with a hissing sound. The ribbon gradually dissolves to form a colourless solution, a steamy fume being evolved at the same time. The test tube becomes hot. There is a flash of yellow flame and a ‘pop’ sound is heard. The mouth of the tube becomes misty.
d.
Guidelines about observation in chemistry 2, 3, 7 1, 4, 8; 5 3(a) & (b), 7 1, 8(d)
©Aristo Educational Press Ltd. 2003
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1, 6, 7, 8(a) & (b) 8
©Aristo Educational Press Ltd. 2003
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Experiment 2.1
To observe Brownian movement of smoke particles (T/S)
3.
c. d.
There are many pin-point bright dots moving in a random way. The bright dots are reflections of light from smoke particles suspended in air. The smoke particles are hit continually by much smaller, rapidly moving air particles. A smoke particle is hit more strongly on one side and then another. It therefore first moves this way and then that way in a random zigzag path.
©Aristo Educational Press Ltd. 2003
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Experiment 2.2
To prepare a compound by direct combination of elements and to compare properties of the compound and its constituent elements (S/T)
1.
Greyish black.
2.
Yellow.
3.
f.
Dark brown / greyish black.
5.
b.
It turns from white to shiny dark brown.
6.
b. d.
Bubbles of a colourless gas are evolved. (A bad smell can also be detected.) A ‘pop’ sound is heard. A flash of yellow flame is seen. The mouth of the test tube becomes misty. Hydrogen.
7.
b.
Some sulphur powder floats and some suspends in water. There is no other visible change. No.
8.
Black powder.
9.
Colourless gas.
10. e.
Colourless (may be misty).
11. c.
Only in tube ‘x’.
12. d.
No. Only carbon dioxide.
©Aristo Educational Press Ltd. 2003
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13. a. Compound iron(II) sulphide
Constituent elements iron (powder)
sulphur
Appearance
dark brown solid
greyish black solid
yellow solid
Action of dilute hydrochloric acid
reacts to give hydrogen sulphide gas
reacts to give hydrogen gas
no reaction
Compound carbon dioxide
Constituent elements carbon
oxygen
Appearance
colourless gas
black powder
colourless gas
Action of limewater
turns limewater milky
no reaction
no reaction
b.
A compound has properties entirely different from those of its constituent elements.
©Aristo Educational Press Ltd. 2003
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Experiment 3.1
Test for oxygen
1.
It is colourless.
2.
No, it does not rise up. No, oxygen is not soluble in water.
3.
No. No, because there may be other gases that are odourless and insoluble in water.
4.
a. c.
5
The limewater remains clear. No change. No, most gases do not react with limewater to give observable change.
6.
The splint relights. Yes, oxygen can give a positive result with the glowing splint test. Oxygen is the only gas that can relight a glowing splint.
7.
The splint burns more brightly. Yes, only oxygen can make a burning splint burn more brightly.
8.
Testing with a glowing splint is the most suitable. From glowing to relighting is an obvious change that is easy to see. Testing with a burning splint is not as good as with a glowing splint. This is because the change from a smaller fire to a bigger one is not so definite, and the judgement may be subjective. Testing with limewater, pH paper or judging from colour, smell, solubility are not suitable as they all give negative test result that are similar to some other gases like nitrogen.
Yellow. Remains yellow. No, there is no colour change. No. There may be other gases that are neutral / do not have any colour change with pH paper.
©Aristo Educational Press Ltd. 2003
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Experiment 4.1
Extraction of common salt from sea water
3.
A circular disc of white, powdery solid remains. (Some concentric white circular stains appear.) It is common salt (other salts would be present besides NaCl). No.
4.
Some crystals form. It is pure common salt.
5.
Yes.
6.
evaporation saturated; crystallization
©Aristo Educational Press Ltd. 2003
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Experiment 4.2
Isolation of pure water from sea water (T)
1.
It is blue.
2.
No, no solid residue is left. No, it is not a suitable method. The experiment result shows that the dissolved particles of salt can pass through the filter paper and remain in the filtrate.
3.
b.
c.
Some steam is seen, some of the steam condenses to water droplets. Colourless liquid (water) drops out, the water boils and turns into water vapour and the steam condense back to water when touching the cold glass tube. Yes. The experiment shows that the blue dissolved material is not in the distillate. Salt does not vaporize at the boiling point of water (100℃). A lot of steam is seen, the water drops gradually disappear. A lot of steam comes out. The glass wall of the condenser is heated up by the steam and thus can no longer cool and condense the steam into water droplets.
4.
The steam disappears. The steam quickly disappears, and clear colourless liquid (water) drops down. The glass wall of the condenser is cooled by the running water and can condense steam to water again. The condenser is for cooling the steam and condensing it into water. The cold water entering at the lower end and leaves at the upper end. This can ensure a better cooling effect of the condenser.
6.
a. b.
boiling; condensation solvent; non-volatile; water
©Aristo Educational Press Ltd. 2003
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Experiment 4.3
Tests to show the presence of sodium and chloride ions in common salt
1.
c.
Golden yellow.
2. Compound
Flame colour
Compound
Flame colour
Compound
Flame colour
Sodium chloride
golden yellow
Sodium sulphate
golden yellow
Sodium carbonate
golden yellow
Potassium chloride
lilac
Potassium sulphate
lilac
Potassium carbonate
lilac
Calcium chloride
brick red
Copper(II) chloride
bluish green
Yes, they are all golden yellow in colour. No, they are all different. It can be concluded that flame colour in the flame test depends on the metal (ion) part of the compound, and does not depend on the non-metal (ion) part of the compound. Compound
3.
b. c.
Flame colour
Sodium compound
golden yellow
Potassium compound
lilac
Calcium compound
brick red
Copper compound
bluish green
A thick, turbulent white precipitate appears. No, there is no change. The white precipitate remains.
4. Solution
Effect of adding silver nitrate solution
Effect of further addition of dilute nitric acid
Potassium chloride
White ppt.
No effect
Calcium chloride
White ppt.
No effect
Copper(II) chloride
White ppt.
No effect
Sodium sulphate
Clear solution
No effect
©Aristo Educational Press Ltd. 2003
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Potassium sulphate
Clear solution
No effect
Sodium carbonate
White ppt.
Effervescence, ppt. disappears
Potassium carbonate
White ppt.
Effervescence, ppt. disappears
No, sodium sulphate gives no precipitate. Yes, they all give a white precipitate. No, all sulphates do not give a precipitate. Yes. The precipitate dissolves and a clear solution remains. This is to distinguish between chlorides and carbonates. The precipitate formed by chloride with silver nitrate cannot dissolve in dilute nitric acid. 5.
a. b. c.
golden yellow white precipitate; nitric dissolves
©Aristo Educational Press Ltd. 2003
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Experiment 4.4
To show the presence of water in a given sample
1.
c.
The copper(II) sulphate changes from white to blue.
2. Liquid
Effect on anhydrous copper(II) sulphate
Salt solution
changes from white to blue
Ethanol
no change
Oil
no change
Dry cleaning liquid
no change
Only salt solution gives the same result as water. 3.
a. b.
The cobalt chloride paper changes from blue to pink. Liquid
Effect on cobalt chloride test paper
Salt solution
changes from blue to pink
Ethanol
no change
Oil
no change
Dry cleaning liquid
no change
Only salt solution can give the same result as water. 4.
a. c.
white; blue; blue; pink anhydrous copper(II) sulphate; dry cobalt chloride paper
©Aristo Educational Press Ltd. 2003
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Experiment 5.1
Action of heat, water and acids on calcium carbonate
2.
A white powdery solid remains. The crystalline shape of the chips is lost.
3.
a. b.
4.
Effervescence occurs, a colourless gas is evolved. The limewater turns milky. calcium carbonate + hydrochloric acid → calcium chloride + carbon dioxide
5.
a. b.
A white suspension forms. The colour of the pH paper does not change. No, because there is still suspension of calcium carbonate powder and the pH of the resulting suspension does not change.
calcium oxide; carbon dioxide hydrochloric acid; carbon dioxide
©Aristo Educational Press Ltd. 2003
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Experiment 5.2 sample laboratory report Title: Chemical tests for calcium carbonate Purpose To test for the presence of calcium carbonate in a solid sample of limestone in the school laboratory.
Apparatus and chemicals used (A)
To test for calcium ions by flame test
• Platinum wire or nichrome wire • Bunsen burner and matches • Mortar and pestle • Watch glass • Test tubes • Test tube rack • Heat-resistant mat (B)
• Limestone (calcium carbonate) • Concentrated hydrochloric acid
To test for carbonate ions:
• Boiling tubes (one of which fitted with a rubber stopper carrying a bent delivery tube) • Test tube rack • Test tubes
• Limestone (calcium carbonate) • Bicarbonate indicator or limewater • Dilute hydrochloric acid
Chemical reaction involved calcium carbonate + hydrochloric acid → calcium chloride + carbon dioxide + water
Procedure (A)
To test for calcium ions by flame test
1. 2.
The solid sample of limestone was crushed into powder by using the mortar and pestle. The nichrome or platinum wire, was dipped into a test tube of concentrated hydrochloric acid and then heated in the hottest part of the Bunsen flame until no characteristic colour shown. After the cleaning, the wire was dipped into the concentrated hydrochloric acid again (Figure 1a) and then into the crushed sample of calcium carbonate (Figure 1b).
3.
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4.
The powdery sample on the wire was heated in the hottest part of the non-luminous flame (Figure 1c).
5.
The colour of the flame was observed and recorded.
(B)
To test for carbonate ions:
6. 7.
5 spatula measures of calcium carbonate were put into a boiling tube (Figure 2a). 5 cm3 of dilute hydrochloric acid was added to the boiling tube containing calcium carbonate (Figure 2b). (a) The boiling tube was quickly covered with a rubber stopper fitted with a bent delivery tube (Figure 2c).
8.
Any gas given out was directed into another boiling tube containing 5 cm3 of limewater or bicarbonate indicator (Figure 2d). Any change in the limewater or bicarbonate indicator was observed and recorded. (b)
9.
©Aristo Educational Press Ltd. 2003
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Observation 1. 2.
(Reference to Step 4): A brick-red flame was observed. (Reference to Step 8): Colourless gas bubbles were given off. Limewater was turned milky OR Bicarbonate indicator was turned from red to yellow.
Interpretation 1. 2.
Calcium ion, gave a brick-red flame colour. Calcium carbonate reacted with acid to give carbon dioxide. calcium carbonate + hydrochloric acid → calcium chloride 1 carbon dioxide 1 water Carbon dioxide turned limewater milky OR Carbon dioxide turned bicarbonate indicator from red to yellow. limewater + carbon dioxide → calcium carbonate + water colourless white precipitate
Discussion 1.
2.
The nichrome/platinum wire used must be sufficiently clean in the flame test to avoid interference. It is suggested that the cleaning process should be repeated a few more times until the wire gives a non-luminous flame. The limewater used for testing carbon dioxide should be freshly prepared. If the limewater used is slightly turbid, it should be filtered to remove the undissolved solid.
Conclusion 1. 2.
Calcium ion gives a brick-red flame in the flame test. Calcium carbonate reacts with dilute acid to give carbon dioxide.
Answers to questions for further thought 1.
2.
Calcium carbonate can be used in manufacturing paper, paints, plastics, adhesives, etc. Calcium carbonate can also be used to lower the acidity in soil and lakes as a result of acid rain. If the marble statue is kept outdoor, it is quite difficult to prevent it from weathering. Acid rain will speed up the weathering because the acidic rainwater will react with calcium carbonate and dissolve it. Even though there is no acid rain, calcium carbonate will react with carbon dioxide and water to form the soluble calcium hydrogencarbonate. calcium carbonate + hydrochloric acid → calcium chloride + carbon dioxide + water calcium carbonate + carbon dioxide + water → calcium hydrogencarbonate Some coating such as wax or lacquer may be put on the statues.
©Aristo Educational Press Ltd. 2003
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3. Lithium ions also give red flame which may be confused with that of calcium ions. The hydrogencarboante ion, which also gives carbon dioxide when being heated or treated with acids, will interfere with the interpretation of the carbonate test.
©Aristo Educational Press Ltd. 2003
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Experiment 6.1
To inspect samples of some common substances
1.
b. Substance
Colour
Physical state
Component element(s)
Common salt
solid
white or colourless
sodium, chlorine
Distilled water
liquid
colourless
hydrogen, oxygen
Aluminium foil
solid
silvery white
aluminium
Sucrose (sugar)
solid
white
carbon, hydrogen, oxygen
Sulphur powder
solid
yellow
sulphur
Copper(II) sulphate crystals
solid
blue
copper, sulphur, hydrogen, oxygen
Sand
solid
greyish yellow
silicon, oxygen
Peanut oil
liquid
yellow
carbon, hydrogen, oxygen
Ethanol
liquid
colourless
carbon, hydrogen, oxygen
Sodium bromide
solid
white
sodium, bromine
Potassium iodide
solid
white
potassium, iodine
Argon (in a light bulb)
gas
colourless
argon
2.
3.
a.
Common salt, aluminium foil, sucrose, sulphur powder, copper(II) sulphate crystals, sand, sodium bromide, potassium iodide.
b.
Distilled water, peanut oil, ethanol.
c.
Argon.
b.
Common salt, distilled water, sucrose, copper(II) sulphate crystals, sand, peanut oil, ethanol, sodium bromide, potassium iodide.
©Aristo Educational Press Ltd. 2003
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4.
No.
5.
solid; solid solid; solid; solid solid; solid; solid solid; solid; solid solid; solid; liquid gas; liquid; solid gas; solid; gas gas; gas; solid solid
6.
solids; mercury; liquid; gases; solids; bromine; liquid; solids
7.
shiny; silvery white (or grey); dull; various
©Aristo Educational Press Ltd. 2003
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Experiment 6.2 2.
To find which elements conduct electricity
Yes. Yes.
3.
Aluminium, iron, lead, magnesium, zinc and graphite.
4.
Yes. Yes.
5.
No. No.
6.
All metals are conductors of electricity. All non-metals (except carbon in the form of graphite) are non-conductors of electricity.
©Aristo Educational Press Ltd. 2003
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Experiment 7.1
To investigate which elements show similar chemical properties (T/S)
2.
3.
4.
e.
The potassium piece melts into a silvery ball and moves about quickly on the water surface with a hissing sound, burning spontaneously with a lilac flame, and finally disappears completely.
f.
The red litmus turns blue.
a.
The sodium piece melts to a silvery ball. It moves about quickly on the water surface with a hissing sound, becoming smaller in size, until finally it disappears completely.
b.
The red litmus turns blue.
a.
The nail sinks to the bottom of the trough. There is no visible change.
b.
The red litmus paper remains red.
5.
Potassium and sodium.
7.
There is a very rapid evolution of colourless gas. The calcium granules dissolve quickly to give a colourless solution. There is a rapid evolution of colourless gas. The magnesium dissolves quickly to give a colourless solution. The copper turnings sink to the bottom. There is no visible change.
8.
Calcium and magnesium.
10.
Very pale greenish yellow solution. Light brown (or yellow) solution. Brown solution.
©Aristo Educational Press Ltd. 2003
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Yellow solid. 11.
Colourless solution. Colourless solution. Colourless solution. A suspension of yellow solid in colourless liquid (no visible change). Tubes ‘1’, ‘2’ and ‘3’ only.
12.
Chlorine, bromine and iodine.
13.
a. b. c.
sodium magnesium bromine and iodine
©Aristo Educational Press Ltd. 2003
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Experiment 8.1
To find out which compounds conduct electricity
2.
No. No.
3.
No.
4.
No. No.
5.
Aqueous sodium chloride and aqueous potassium nitrate.
6.
c.
No.
7. State or form
Solid
Liquid Aqueous (or molten) solution
Constituent elements
M/N or N/N?
carbon, hydrogen
N/N
Compound
8.
Wax
×
×
Sugar
×
×
carbon, hydrogen, oxygen
N/N
Sodium chloride
×
sodium, chlorine
M/N
Potassium nitrate
×
potassium, nitrogen, oxygen
M/N
hydrogen, oxygen
N/N
carbon, hydrogen, oxygen
N/N
Water
×
Ethanol
×
×
Sodium chloride conducts electricity in aqueous solution, but not in solid state.
©Aristo Educational Press Ltd. 2003
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9.
Potassium nitrate conducts electricity in aqueous solution, but not in solid state.
10. Some do, some do not. 11. Some do, some do not. 12
a. b.
metals; non-metals; solid; molten; dissolved; in water; electrolytes non-metals; molten; dissolved in water; non-electrolytes
©Aristo Educational Press Ltd. 2003
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Experiment 8.2
Effect of electricity on molten lead(II) bromide (T)
1.
d.
No. No.
2.
Yes. Yes. Molten lead(II) bromide consists of mobile charged particles (called ions) which can conduct electricity.
3.
b.
Reddish brown. Bromine.
4.
g.
Silvery grey solid. Negative electrode.
h.
Yes. It is lead metal. Lead(II) ions. Positive charge.
5.
solid; molten; charged; ions; mobile; chemical decomposition; electrolytes; positive; ions; negative; ions; electricity; lead; bromine
©Aristo Educational Press Ltd. 2003
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Experiment 8.3
Colours of ions
1. chloride ion
nitrate ion
sulphate ion
chromate ion
dichromate ion
permanganate ion
potassium ion
potassium chloride colourless
potassium nitrate colourless
potassium sulphate colourless
potassium chromate yellow
potassium dichromate orange
potassium permanganate purple
sodium ion
sodium chloride colourless
sodium nitrate colourless
sodium sulphate colourless
sodium chromate yellow
sodium dichromate orange
ammonium ion
ammonium chloride colourless
ammonium nitrate colourless
ammonium sulphate colourless
zinc ion
zinc chloride colourless
zinc nitrate colourless
zinc sulphate colourless
calcium ion
calcium chloride colourless
calcium nitrate colourless
copper(II) ion
copper(II) chloride blue
copper(II) nitrate blue
copper(II) sulphate blue
iron(II) ion
iron(II) chloride green
iron(II) nitrate green
iron(II) sulphate green
iron(III) ion
iron(III) chloride brown
iron(III) nitrate brown
nickel(II) ion
nickel(II) chloride green
nickel(II) nitrate green
Anions
Cations
2.
ammonium dichromate orange
nickel(II) sulphate green
The solutions of potassium chloride, potassium nitrate, potassium sulphate, sodium chloride, sodium nitrate, sodium sulphate, ammonium chloride, ammonium nitrate, ammonium sulphate, zinc chloride, zinc nitrate, zinc sulphate, calcium chloride and calcium nitrate. Colourless. Colourless; Colourless
©Aristo Educational Press Ltd. 2003
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Colourless; Colourless Colourless Colourless; Colourless
3.
Colourless a. Blue b.
Green
c.
Brown
d.
Green
Cations. Since the anions of these compounds are colourless (known from Step 2), it may be deduced that the colours of the solutions are due to the cations. Blue; Green Brown; Green
4.
Purple. Colourless. Purple.
5.
Yellow; Orange
©Aristo Educational Press Ltd. 2003
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Experiment 8.4
Migration of coloured ions (S/T)
1.
b.
The sodium sulphate solution provides more ions to conduct electricity. (If tap water were used, a smaller current and hence slower migration of ions would result.)
2.
b.
Purple colour. Positive electrode. No. Because the purple colour moves towards only one side of the paper.
3.
colourless; purple; mobile; negatively; permanganate; purple; positively; potassium; cannot be seen
5.
e.
Deep green.
6.
c.
Yellow colour. Blue colour.
7.
blue; yellow; negatively; chromate; yellow; positively; copper(II); blue
©Aristo Educational Press Ltd. 2003
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Experiment 8.5 1.
90º.
2.
6
To build a lattice model of sodium chloride
6 3.
Cubic.
4.
Yes. No. No. Giant ionic structure.
©Aristo Educational Press Ltd. 2003
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Experiment 9.1 2.
To build models of diamond and quartz
4 No. Giant covalent structure.
5.
No. Giant covalent structure.
©Aristo Educational Press Ltd. 2003
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Experiment 10.1
To compare the properties of sodium chloride, dry ice, wax and quartz
6. Substance
Physical state
Does it melt easily?
Is it volatile?
Is it soluble in water?
Does its aqueous solution conduct electricity?
solid
no
no
yes
yes
yes
(The solid disappears; a lot of gas bubbles are given out)
no
Sodium Chloride Dry ice
solid (when just (It sublimes at taken out from room vacuum flask), conditions) soon changes to gas
Wax
solid
yes
no
no
Quartz
solid
no
no
no
7.
a.
i. Sodium chloride. ii. Dry ice, wax. iii. Quartz.
b. Substance
Type of structure
Sodium chloride
giant ionic structure
Dry ice
simple molecular structure
Wax
simple molecular structure
Quartz
giant covalent structure
©Aristo Educational Press Ltd. 2003
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8.
a.
They are high-melting solids, non-volatile and soluble in water. They conduct electricity when dissolved in water and when molten.
b.
They have low melting points and are insoluble in water. They do not conduct electricity whether solid or liquid.
c.
They are high-melting solids, insoluble in water.
©Aristo Educational Press Ltd. 2003
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Experiment 11.1 1.
b.
To extract metals from metal oxides (T/S)
brownish black black white yellow
5.
Only on heating silver oxide. Silver oxide. 2Ag2O(s) → 4Ag(s) + O2(g) Silver. Only silver oxide decomposes on heating, while the oxides of the other metals do not. This shows that silver forms the least stable oxide and is thus least reactive.
9. Metal oxide
10.
Signs of formation of a metal (if any)
Copper(II) oxide
reddish brown solid
Zinc oxide
none
Lead(II) oxide
grey solid
Does reduction occur?
×
Only for copper(II) oxide and lead(II) oxide. 2CuO(s) + C(s) → 2Cu(s) + CO2(g); 2PbO(s) + C(s) → 2Pb(s) + CO2(g)
©Aristo Educational Press Ltd. 2003
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11.
Zinc oxide, lead(II) oxide, copper(II) oxide, silver oxide OR zinc oxide, copper(II) oxide, lead(II) oxide, silver oxide.
12.
Zinc, lead, copper, silver OR zinc, copper, lead, silver.
©Aristo Educational Press Ltd. 2003
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Experiment 12.1
To arrange five metals in order of reactivity
1.
b.
After a short time, the calcium burns vigorously with a brick-red flame. A white powder is left.
2.
It does not burn even when red hot. The surface of iron becomes black. It does not burn even when strongly heated. The surface of copper becomes black on strong heating. It burns with a dazzling white flame, forming a white powder.
3.
b.
Calcium granules sink to the bottom of the tube. Colourless gas bubbles are given out at a moderate rate. A milky suspension eventually forms. The tube gets warm.
4.
b.
No reaction. No reaction. Very slow reaction. Tiny gas bubbles are given out very slowly from the metal surface. No reaction.
5.
c.
There is rapid effervescence of a colourless gas. The calcium granule soon dissolves. The tube gets warm quickly.
6.
Evolution of colourless gas bubbles at a moderate rate from the metal surface only heating. No reaction even on heating. There is effervescence of a colourless gas. The ribbon dissolves rapidly. Slow evolution of tiny gas bubbles from the metal surface only on heating. A white solid precipitate is formed.
©Aristo Educational Press Ltd. 2003
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on
7. Metal
Does the metal burn on strong heating?
8.
Calcium
Iron
×
Copper
×
Magnesium
Lead
×
vigorous
vigorous
react with water?
very rapid
×
moderate (on heating)
×
×
×
moderate
react with dilute hydrochloric acid?
very slow
rapid
slow (on heating)
Calcium, magnesium, iron, lead, copper.
©Aristo Educational Press Ltd. 2003
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Experiment 12.2
Displacement reactions of metals in aqueous solutions
8. Solution
Cu2+ (aq)
Mg2+ (aq)
Zn2+ (aq)
Fe2+ (aq)
Ag+ (aq)
×
×
×
Metal Cu Mg
Zn
×
Fe
×
×
Ag
×
×
×
×
9.
Magnesium, zinc, iron, copper, silver.
10.
Consider two metals magnesium and zinc. Magnesium is higher than zinc in the metal reactivity series. This means that magnesium loses electrons more readily than zinc. Thus magnesium atoms lose electrons to become magnesium ions, while zinc ions are forced to gain electrons to become zinc atoms. Thus magnesium displaces zinc metal from zinc sulphate solution. On the other hand, there is no reaction between zinc metal and magnesium sulphate solution. Other displacement reactions can be explained similarly.
11.
Cu(s) + 2Ag+ (aq) → Cu2+(aq) + 2Ag(s)
12.
higher; lower; salt; more; less; less; more
©Aristo Educational Press Ltd. 2003
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Experiment 13.1
To determine the empirical formula of magnesium oxide (T/S) (Extension)
1.
28.094 g
2.
b.
3.
28.303 g
5.
This lets in oxygen from air to react with magnesium. If the lid was not lifted up, the small amount of oxygen inside the crucible would soon be used up in burning.
6.
The white smoke is magnesium oxide. If some escapes, the result will be inaccurate (the mass of oxygen found would be lower than it should be).
9.
White.
Dull grey.
Magnesium oxide. magnesium + oxygen → magnesium oxide 11.
28.439 g
12.
0.345 0.209 0.136 Mg
O
Masses (in g)
0.209
0.136
Number of moles of atoms (mol)
0.209
0.136
24.3
16.0
©Aristo Educational Press Ltd. 2003
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Relative number of moles
= 0.00860
= 0.00850
0.00860
0.00850
0.00850
0.00850
= 1.01
=1
≒ 1
13.
magnesium oxide; one; MgO
©Aristo Educational Press Ltd. 2003
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Experiment 14.1 sample laboratory report Title:
To investigate factors that influence rusting
Purpose To investigate factors that influence rusting.
Apparatus and chemicals used • • • • •
Beaker (250 cm3) Boiling tubes (150 × 18 mm) in rack Thermometer (−10 2 110°C) Small self-adhesive labels Stopwatch sea • 6 clean iron nails (~5 cm long) • Magnesium ribbon (7 cm long)
• Copper wire (7 cm long) • Dilute hydrochloric acid (1 M) • Sea water / salt water (a solution of 30 g sodium chloride per dm3 may be taken as water) • Hot water (about 80°C) • Warm agar solution containing rust indicator
Chemical reactions involved 4Fe(s) + 3O2 (g) + 2nH2O (l) → 2Fe2O3 • nH2O(s)
Procedure To investigate factors affecting rusting
In order to show rusting, a warm rust indicator solution (pale yellow in colour) was used. It contained agar and turned into a gel on cooling. 1. Six iron nails were cleaned, degreased and put into separate test tubes. 2. The six test tubes were labelled and set up as shown in Figure 1: Test tube 1: A clean iron nail was put into the rust indicator solution. (This is the control.) Test tube 2: A length of magnesium ribbon was used to wrap tightly around a clean iron nail and put into the rust indicator solution. Test tube 3: A length of copper wire was used to wrap tightly around a clean iron nail and put into the rust indicator solution. Test tube 4: A clean iron nail was put into the rust indicator solution mixed with 1 cm3 sea water. Test tube 5: A clean iron nail was put into the rust indicator solution mixed with 1 cm3 dilute HCl. Test tube 6: A clean iron nail was put into the rust indicator solution which was immersed into a hot water bath (~80°C). 3. All test tubes were left to stand for 20 minutes. They were then observed carefully, especially for the appearance of blue colour. ©Aristo Educational Press Ltd. 2003
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4.
The rate of rusting could be roughly determined by comparing • the time for the appearance of blue colour around the differently treated iron nails. • the size of the blue patches around the differently treated iron nails.
¡C)
Observation 1. 2. 3.
4. 5.
6.
Test tube 1: Blue colour mainly appeared around the head and tip. Test tube 2: No blue colour appeared. Some colourless gas bubbles were given off around the magnesium ribbon. Test tube 3: Blue colour appeared along the whole length of the wrapped iron nail. When compared to test tube 1, the time for the appearance of blue colour was shorter and the size of the blue patch is bigger. Test tube 4: Blue colour appeared along the iron nail. When compared to test tube 1, the time for the appearance of blue colour was shorter and the size of the blue patch is bigger. Test tube 5: Blue colour appeared quite rapidly along the iron nail. When compared to test tube 1, the time for the appearance of blue colour was shorter and the size of the blue patch is bigger. Some colourless gas bubbles were given off around the iron nail. Test tube 6: Blue colour appeared along the iron nail. When compared to test tube 1, the time for the appearance of blue colour was shorter and the size of the blue patch is bigger.
Interpretation 1.
Test tube 1: Rusting occurred mainly around the head and tip. This was a control experiment.
©Aristo Educational Press Ltd. 2003
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2.
3.
4. 5.
6.
Test tube 2: No rusting occurred. This was because magnesium, being more reactive, lost electrons more easily than iron. Since magnesium lost electrons to iron, iron was prevented from losing electrons and could not form Fe2+(aq) ions. Thus, the rate of rusting was slowed down or the iron nail was even being protected from rusting. Test tube 3: Rusting occurred quite quickly. This was because copper, being less reactive, encouraged the lost of electrons from iron to form Fe2+(aq) ions. Thus, the rate of rusting was increased. Test tube 4: Rusting occurred quite quickly. The presence of soluble salts such as sodium chloride could speed up rusting. Test tube 5: Rusting occurred most easily and quickly. Iron reacted with acid to form Fe2+(aq) ions. Fe(s) + 2H+(aq) → Fe2+(aq) + H2(g) Thus, the rate of rusting was increased in the presence of acid. Test tube 6: Rusting occurred quite quickly. An increase in temperature always increased the rate of chemical reaction including rusting.
Discussion 1.
2. 3.
Clean iron nails should be used. Sand paper may be needed to remove any surface coating. An aqueous solution of detergent or some acetone can be used to degrease the iron nails before the experiment. Magnesium ribbon and the copper wire should be used to wrap around the iron nail tightly. This was to make sure a good contact between the metal and iron nail. Sea water and dilute HCl could be added with stirring to test tubes 5 and 6 respectively before adding the iron nail.
Conclusion 1. 2.
Air and water are essential for rusting to occur. The presence of acidic substances, soluble salts, high temperature, uneven or sharply pointed regions (i.e. head and tip of nails) and the contact with a less reactive metal are all factors speeding up the rate of rusting.
Answers to questions for further thought 1. 2. 3.
Sodium chloride. Cheap, abundant. Rusting should be faster in Hong Kong where the humidity is high.
©Aristo Educational Press Ltd. 2003
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Experiment 14.2 3.
To prevent rusting
A blue colour appears, mainly around the head and tip of the nail. Gas bubbles appear around the magnesium ribbon. No blue colour appears. Nail ‘1’ only.
4.
Rusting occurs in nail ‘1’, but not in nail ‘2’. This is because magnesium, being more reactive than iron, loses electrons more easily. Since magnesium loses electrons to iron, iron is prevented from losing electrons and cannot form Fe2+(aq) ions. Iron is thus protected from rusting.
5.
To prevent rusting, connect the iron piece to a more reactive metal. This method is effective because the other metal will give up electrons in preference, preventing the formation of Fe2+ ions (sacrificial protection).
7.
A blue colour appears, mainly around the head and tip of the nail. No observation. Nail ‘3’ only.
8.
Rusting occurs in nail ‘3’ but not in nail ‘4’. This is because both water and air are kept out from painted iron, thus it is protected from rusting.
9.
Yes, both water and air are kept out from the nail covered with grease.
10.
sacrificial protection; paint
©Aristo Educational Press Ltd. 2003
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INTRODUCING CHEMISTRY
Chapter 1 1.1
What is Chemistry?
Observation in chemistry
Chapter 2
3
The fundamentals of chemistry
2.1
To observe Brownian movement of smoke particles (T/S)
4
2.2
To prepare a compound by direct combination of elements and to compare properties of the compound and its constituent elements (S/T)
PART II
PLANET EARTH
Chapter 3 3.1
5
The atmosphere
Test for oxygen
Chapter 4
7
Oceans
4.1
Extraction of common salt from sea water
8
4.2
Isolation of pure water from sea water (T)
9
4.3
Tests to show the presence of sodium and chloride ions in common salt
10
4.4
To show the presence of water in a given sample
12
Chapter 5
Rocks and minerals
5.1
Action of heat, water and acids on calcium carbonate
13
5.2
Chemical tests for the presence of calcium carbonate
14
PART III THE MICROSCOPIC WORLD Chapter 6
Atomic structure
6.1
To inspect samples of some common substances
17
6.2
To find which elements conduct electricity
19
©Aristo Educational Press Ltd. 2003
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Chapter 7 7.1
The Periodic Table
To investigate which elements show similar chemical properties (T/S)
Chapter 8
21
Chemical bonding: Ionic bonding
8.1
To find out which compounds conduct electricity
23
8.2
Effect of electricity on molten lead(II) bromide (T)
25
8.3
Colours of ions
26
8.4
Migration of coloured ions (S/T)
28
8.5
To build a lattice model of sodium chloride
29
Chapter 9 9.1
Chemical bonding: Covalent bonding
To build models of diamond and quartz
Chapter 10
30
Structures and properties
10.1 To compare the properties of sodium chloride, dry ice, wax and quartz
PART IV
31
METALS
Chapter 11
Occurrence and extraction of metals
11.1 To extract metals from metal oxides (T/S)
Chapter 12
32
Reactivity of metals
12.1 To arrange five metals in order of reactivity
33
12.2 Displacement reactions of metals in aqueous solutions
35
Chapter 13
Reacting masses
13.1 To determine the empirical formula of magnesium oxide (T/S) (Extension) 36
Chapter 14
Corrosion of metals and their protection
14.1 To investigate factors that influence rusting
37
14.2 To prevent rusting
40
©Aristo Educational Press Ltd. 2003
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Experiment 1.1
Observation in chemistry
1.
The alcohol burns with a pale blue flame having a yellow tip. The alcohol soon dries and the burning also stops. There is no blackening of/no soot left on the watch glass.
2.
b.
The potassium dichromate solution changes from orange to dark green.
3.
b. d.
A blue precipitate is formed. The precipitate dissolves to form a deep blue solution.
4.
There is no visible change, but the test tube becomes warm.
5.
c.
The white solid dissolves to form a colourless solution. The test tube becomes cool.
6.
d.
There is no visible change, but an irritating smell of ammonia can be detected.
7.
c.
The clear solution gradually turns cloudy white and then cloudy yellow. A choking smell can be detected.
8.
c.
There is effervescence — small colourless gas bubbles are evolved from the magnesium surface, with a hissing sound. The ribbon gradually dissolves to form a colourless solution, a steamy fume being evolved at the same time. The test tube becomes hot. There is a flash of yellow flame and a ‘pop’ sound is heard. The mouth of the tube becomes misty.
d.
Guidelines about observation in chemistry 2, 3, 7 1, 4, 8; 5 3(a) & (b), 7 1, 8(d)
©Aristo Educational Press Ltd. 2003
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1, 6, 7, 8(a) & (b) 8
©Aristo Educational Press Ltd. 2003
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Experiment 2.1
To observe Brownian movement of smoke particles (T/S)
3.
c. d.
There are many pin-point bright dots moving in a random way. The bright dots are reflections of light from smoke particles suspended in air. The smoke particles are hit continually by much smaller, rapidly moving air particles. A smoke particle is hit more strongly on one side and then another. It therefore first moves this way and then that way in a random zigzag path.
©Aristo Educational Press Ltd. 2003
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Experiment 2.2
To prepare a compound by direct combination of elements and to compare properties of the compound and its constituent elements (S/T)
1.
Greyish black.
2.
Yellow.
3.
f.
Dark brown / greyish black.
5.
b.
It turns from white to shiny dark brown.
6.
b. d.
Bubbles of a colourless gas are evolved. (A bad smell can also be detected.) A ‘pop’ sound is heard. A flash of yellow flame is seen. The mouth of the test tube becomes misty. Hydrogen.
7.
b.
Some sulphur powder floats and some suspends in water. There is no other visible change. No.
8.
Black powder.
9.
Colourless gas.
10. e.
Colourless (may be misty).
11. c.
Only in tube ‘x’.
12. d.
No. Only carbon dioxide.
©Aristo Educational Press Ltd. 2003
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13. a. Compound iron(II) sulphide
Constituent elements iron (powder)
sulphur
Appearance
dark brown solid
greyish black solid
yellow solid
Action of dilute hydrochloric acid
reacts to give hydrogen sulphide gas
reacts to give hydrogen gas
no reaction
Compound carbon dioxide
Constituent elements carbon
oxygen
Appearance
colourless gas
black powder
colourless gas
Action of limewater
turns limewater milky
no reaction
no reaction
b.
A compound has properties entirely different from those of its constituent elements.
©Aristo Educational Press Ltd. 2003
-7-
Experiment 3.1
Test for oxygen
1.
It is colourless.
2.
No, it does not rise up. No, oxygen is not soluble in water.
3.
No. No, because there may be other gases that are odourless and insoluble in water.
4.
a. c.
5
The limewater remains clear. No change. No, most gases do not react with limewater to give observable change.
6.
The splint relights. Yes, oxygen can give a positive result with the glowing splint test. Oxygen is the only gas that can relight a glowing splint.
7.
The splint burns more brightly. Yes, only oxygen can make a burning splint burn more brightly.
8.
Testing with a glowing splint is the most suitable. From glowing to relighting is an obvious change that is easy to see. Testing with a burning splint is not as good as with a glowing splint. This is because the change from a smaller fire to a bigger one is not so definite, and the judgement may be subjective. Testing with limewater, pH paper or judging from colour, smell, solubility are not suitable as they all give negative test result that are similar to some other gases like nitrogen.
Yellow. Remains yellow. No, there is no colour change. No. There may be other gases that are neutral / do not have any colour change with pH paper.
©Aristo Educational Press Ltd. 2003
-8-
Experiment 4.1
Extraction of common salt from sea water
3.
A circular disc of white, powdery solid remains. (Some concentric white circular stains appear.) It is common salt (other salts would be present besides NaCl). No.
4.
Some crystals form. It is pure common salt.
5.
Yes.
6.
evaporation saturated; crystallization
©Aristo Educational Press Ltd. 2003
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Experiment 4.2
Isolation of pure water from sea water (T)
1.
It is blue.
2.
No, no solid residue is left. No, it is not a suitable method. The experiment result shows that the dissolved particles of salt can pass through the filter paper and remain in the filtrate.
3.
b.
c.
Some steam is seen, some of the steam condenses to water droplets. Colourless liquid (water) drops out, the water boils and turns into water vapour and the steam condense back to water when touching the cold glass tube. Yes. The experiment shows that the blue dissolved material is not in the distillate. Salt does not vaporize at the boiling point of water (100℃). A lot of steam is seen, the water drops gradually disappear. A lot of steam comes out. The glass wall of the condenser is heated up by the steam and thus can no longer cool and condense the steam into water droplets.
4.
The steam disappears. The steam quickly disappears, and clear colourless liquid (water) drops down. The glass wall of the condenser is cooled by the running water and can condense steam to water again. The condenser is for cooling the steam and condensing it into water. The cold water entering at the lower end and leaves at the upper end. This can ensure a better cooling effect of the condenser.
6.
a. b.
boiling; condensation solvent; non-volatile; water
©Aristo Educational Press Ltd. 2003
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Experiment 4.3
Tests to show the presence of sodium and chloride ions in common salt
1.
c.
Golden yellow.
2. Compound
Flame colour
Compound
Flame colour
Compound
Flame colour
Sodium chloride
golden yellow
Sodium sulphate
golden yellow
Sodium carbonate
golden yellow
Potassium chloride
lilac
Potassium sulphate
lilac
Potassium carbonate
lilac
Calcium chloride
brick red
Copper(II) chloride
bluish green
Yes, they are all golden yellow in colour. No, they are all different. It can be concluded that flame colour in the flame test depends on the metal (ion) part of the compound, and does not depend on the non-metal (ion) part of the compound. Compound
3.
b. c.
Flame colour
Sodium compound
golden yellow
Potassium compound
lilac
Calcium compound
brick red
Copper compound
bluish green
A thick, turbulent white precipitate appears. No, there is no change. The white precipitate remains.
4. Solution
Effect of adding silver nitrate solution
Effect of further addition of dilute nitric acid
Potassium chloride
White ppt.
No effect
Calcium chloride
White ppt.
No effect
Copper(II) chloride
White ppt.
No effect
Sodium sulphate
Clear solution
No effect
©Aristo Educational Press Ltd. 2003
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Potassium sulphate
Clear solution
No effect
Sodium carbonate
White ppt.
Effervescence, ppt. disappears
Potassium carbonate
White ppt.
Effervescence, ppt. disappears
No, sodium sulphate gives no precipitate. Yes, they all give a white precipitate. No, all sulphates do not give a precipitate. Yes. The precipitate dissolves and a clear solution remains. This is to distinguish between chlorides and carbonates. The precipitate formed by chloride with silver nitrate cannot dissolve in dilute nitric acid. 5.
a. b. c.
golden yellow white precipitate; nitric dissolves
©Aristo Educational Press Ltd. 2003
- 12 -
Experiment 4.4
To show the presence of water in a given sample
1.
c.
The copper(II) sulphate changes from white to blue.
2. Liquid
Effect on anhydrous copper(II) sulphate
Salt solution
changes from white to blue
Ethanol
no change
Oil
no change
Dry cleaning liquid
no change
Only salt solution gives the same result as water. 3.
a. b.
The cobalt chloride paper changes from blue to pink. Liquid
Effect on cobalt chloride test paper
Salt solution
changes from blue to pink
Ethanol
no change
Oil
no change
Dry cleaning liquid
no change
Only salt solution can give the same result as water. 4.
a. c.
white; blue; blue; pink anhydrous copper(II) sulphate; dry cobalt chloride paper
©Aristo Educational Press Ltd. 2003
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Experiment 5.1
Action of heat, water and acids on calcium carbonate
2.
A white powdery solid remains. The crystalline shape of the chips is lost.
3.
a. b.
4.
Effervescence occurs, a colourless gas is evolved. The limewater turns milky. calcium carbonate + hydrochloric acid → calcium chloride + carbon dioxide
5.
a. b.
A white suspension forms. The colour of the pH paper does not change. No, because there is still suspension of calcium carbonate powder and the pH of the resulting suspension does not change.
calcium oxide; carbon dioxide hydrochloric acid; carbon dioxide
©Aristo Educational Press Ltd. 2003
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Experiment 5.2 sample laboratory report Title: Chemical tests for calcium carbonate Purpose To test for the presence of calcium carbonate in a solid sample of limestone in the school laboratory.
Apparatus and chemicals used (A)
To test for calcium ions by flame test
• Platinum wire or nichrome wire • Bunsen burner and matches • Mortar and pestle • Watch glass • Test tubes • Test tube rack • Heat-resistant mat (B)
• Limestone (calcium carbonate) • Concentrated hydrochloric acid
To test for carbonate ions:
• Boiling tubes (one of which fitted with a rubber stopper carrying a bent delivery tube) • Test tube rack • Test tubes
• Limestone (calcium carbonate) • Bicarbonate indicator or limewater • Dilute hydrochloric acid
Chemical reaction involved calcium carbonate + hydrochloric acid → calcium chloride + carbon dioxide + water
Procedure (A)
To test for calcium ions by flame test
1. 2.
The solid sample of limestone was crushed into powder by using the mortar and pestle. The nichrome or platinum wire, was dipped into a test tube of concentrated hydrochloric acid and then heated in the hottest part of the Bunsen flame until no characteristic colour shown. After the cleaning, the wire was dipped into the concentrated hydrochloric acid again (Figure 1a) and then into the crushed sample of calcium carbonate (Figure 1b).
3.
©Aristo Educational Press Ltd. 2003
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4.
The powdery sample on the wire was heated in the hottest part of the non-luminous flame (Figure 1c).
5.
The colour of the flame was observed and recorded.
(B)
To test for carbonate ions:
6. 7.
5 spatula measures of calcium carbonate were put into a boiling tube (Figure 2a). 5 cm3 of dilute hydrochloric acid was added to the boiling tube containing calcium carbonate (Figure 2b). (a) The boiling tube was quickly covered with a rubber stopper fitted with a bent delivery tube (Figure 2c).
8.
Any gas given out was directed into another boiling tube containing 5 cm3 of limewater or bicarbonate indicator (Figure 2d). Any change in the limewater or bicarbonate indicator was observed and recorded. (b)
9.
©Aristo Educational Press Ltd. 2003
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Observation 1. 2.
(Reference to Step 4): A brick-red flame was observed. (Reference to Step 8): Colourless gas bubbles were given off. Limewater was turned milky OR Bicarbonate indicator was turned from red to yellow.
Interpretation 1. 2.
Calcium ion, gave a brick-red flame colour. Calcium carbonate reacted with acid to give carbon dioxide. calcium carbonate + hydrochloric acid → calcium chloride 1 carbon dioxide 1 water Carbon dioxide turned limewater milky OR Carbon dioxide turned bicarbonate indicator from red to yellow. limewater + carbon dioxide → calcium carbonate + water colourless white precipitate
Discussion 1.
2.
The nichrome/platinum wire used must be sufficiently clean in the flame test to avoid interference. It is suggested that the cleaning process should be repeated a few more times until the wire gives a non-luminous flame. The limewater used for testing carbon dioxide should be freshly prepared. If the limewater used is slightly turbid, it should be filtered to remove the undissolved solid.
Conclusion 1. 2.
Calcium ion gives a brick-red flame in the flame test. Calcium carbonate reacts with dilute acid to give carbon dioxide.
Answers to questions for further thought 1.
2.
Calcium carbonate can be used in manufacturing paper, paints, plastics, adhesives, etc. Calcium carbonate can also be used to lower the acidity in soil and lakes as a result of acid rain. If the marble statue is kept outdoor, it is quite difficult to prevent it from weathering. Acid rain will speed up the weathering because the acidic rainwater will react with calcium carbonate and dissolve it. Even though there is no acid rain, calcium carbonate will react with carbon dioxide and water to form the soluble calcium hydrogencarbonate. calcium carbonate + hydrochloric acid → calcium chloride + carbon dioxide + water calcium carbonate + carbon dioxide + water → calcium hydrogencarbonate Some coating such as wax or lacquer may be put on the statues.
©Aristo Educational Press Ltd. 2003
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3. Lithium ions also give red flame which may be confused with that of calcium ions. The hydrogencarboante ion, which also gives carbon dioxide when being heated or treated with acids, will interfere with the interpretation of the carbonate test.
©Aristo Educational Press Ltd. 2003
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Experiment 6.1
To inspect samples of some common substances
1.
b. Substance
Colour
Physical state
Component element(s)
Common salt
solid
white or colourless
sodium, chlorine
Distilled water
liquid
colourless
hydrogen, oxygen
Aluminium foil
solid
silvery white
aluminium
Sucrose (sugar)
solid
white
carbon, hydrogen, oxygen
Sulphur powder
solid
yellow
sulphur
Copper(II) sulphate crystals
solid
blue
copper, sulphur, hydrogen, oxygen
Sand
solid
greyish yellow
silicon, oxygen
Peanut oil
liquid
yellow
carbon, hydrogen, oxygen
Ethanol
liquid
colourless
carbon, hydrogen, oxygen
Sodium bromide
solid
white
sodium, bromine
Potassium iodide
solid
white
potassium, iodine
Argon (in a light bulb)
gas
colourless
argon
2.
3.
a.
Common salt, aluminium foil, sucrose, sulphur powder, copper(II) sulphate crystals, sand, sodium bromide, potassium iodide.
b.
Distilled water, peanut oil, ethanol.
c.
Argon.
b.
Common salt, distilled water, sucrose, copper(II) sulphate crystals, sand, peanut oil, ethanol, sodium bromide, potassium iodide.
©Aristo Educational Press Ltd. 2003
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4.
No.
5.
solid; solid solid; solid; solid solid; solid; solid solid; solid; solid solid; solid; liquid gas; liquid; solid gas; solid; gas gas; gas; solid solid
6.
solids; mercury; liquid; gases; solids; bromine; liquid; solids
7.
shiny; silvery white (or grey); dull; various
©Aristo Educational Press Ltd. 2003
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Experiment 6.2 2.
To find which elements conduct electricity
Yes. Yes.
3.
Aluminium, iron, lead, magnesium, zinc and graphite.
4.
Yes. Yes.
5.
No. No.
6.
All metals are conductors of electricity. All non-metals (except carbon in the form of graphite) are non-conductors of electricity.
©Aristo Educational Press Ltd. 2003
- 21 -
Experiment 7.1
To investigate which elements show similar chemical properties (T/S)
2.
3.
4.
e.
The potassium piece melts into a silvery ball and moves about quickly on the water surface with a hissing sound, burning spontaneously with a lilac flame, and finally disappears completely.
f.
The red litmus turns blue.
a.
The sodium piece melts to a silvery ball. It moves about quickly on the water surface with a hissing sound, becoming smaller in size, until finally it disappears completely.
b.
The red litmus turns blue.
a.
The nail sinks to the bottom of the trough. There is no visible change.
b.
The red litmus paper remains red.
5.
Potassium and sodium.
7.
There is a very rapid evolution of colourless gas. The calcium granules dissolve quickly to give a colourless solution. There is a rapid evolution of colourless gas. The magnesium dissolves quickly to give a colourless solution. The copper turnings sink to the bottom. There is no visible change.
8.
Calcium and magnesium.
10.
Very pale greenish yellow solution. Light brown (or yellow) solution. Brown solution.
©Aristo Educational Press Ltd. 2003
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Yellow solid. 11.
Colourless solution. Colourless solution. Colourless solution. A suspension of yellow solid in colourless liquid (no visible change). Tubes ‘1’, ‘2’ and ‘3’ only.
12.
Chlorine, bromine and iodine.
13.
a. b. c.
sodium magnesium bromine and iodine
©Aristo Educational Press Ltd. 2003
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Experiment 8.1
To find out which compounds conduct electricity
2.
No. No.
3.
No.
4.
No. No.
5.
Aqueous sodium chloride and aqueous potassium nitrate.
6.
c.
No.
7. State or form
Solid
Liquid Aqueous (or molten) solution
Constituent elements
M/N or N/N?
carbon, hydrogen
N/N
Compound
8.
Wax
×
×
Sugar
×
×
carbon, hydrogen, oxygen
N/N
Sodium chloride
×
sodium, chlorine
M/N
Potassium nitrate
×
potassium, nitrogen, oxygen
M/N
hydrogen, oxygen
N/N
carbon, hydrogen, oxygen
N/N
Water
×
Ethanol
×
×
Sodium chloride conducts electricity in aqueous solution, but not in solid state.
©Aristo Educational Press Ltd. 2003
- 24 -
9.
Potassium nitrate conducts electricity in aqueous solution, but not in solid state.
10. Some do, some do not. 11. Some do, some do not. 12
a. b.
metals; non-metals; solid; molten; dissolved; in water; electrolytes non-metals; molten; dissolved in water; non-electrolytes
©Aristo Educational Press Ltd. 2003
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Experiment 8.2
Effect of electricity on molten lead(II) bromide (T)
1.
d.
No. No.
2.
Yes. Yes. Molten lead(II) bromide consists of mobile charged particles (called ions) which can conduct electricity.
3.
b.
Reddish brown. Bromine.
4.
g.
Silvery grey solid. Negative electrode.
h.
Yes. It is lead metal. Lead(II) ions. Positive charge.
5.
solid; molten; charged; ions; mobile; chemical decomposition; electrolytes; positive; ions; negative; ions; electricity; lead; bromine
©Aristo Educational Press Ltd. 2003
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Experiment 8.3
Colours of ions
1. chloride ion
nitrate ion
sulphate ion
chromate ion
dichromate ion
permanganate ion
potassium ion
potassium chloride colourless
potassium nitrate colourless
potassium sulphate colourless
potassium chromate yellow
potassium dichromate orange
potassium permanganate purple
sodium ion
sodium chloride colourless
sodium nitrate colourless
sodium sulphate colourless
sodium chromate yellow
sodium dichromate orange
ammonium ion
ammonium chloride colourless
ammonium nitrate colourless
ammonium sulphate colourless
zinc ion
zinc chloride colourless
zinc nitrate colourless
zinc sulphate colourless
calcium ion
calcium chloride colourless
calcium nitrate colourless
copper(II) ion
copper(II) chloride blue
copper(II) nitrate blue
copper(II) sulphate blue
iron(II) ion
iron(II) chloride green
iron(II) nitrate green
iron(II) sulphate green
iron(III) ion
iron(III) chloride brown
iron(III) nitrate brown
nickel(II) ion
nickel(II) chloride green
nickel(II) nitrate green
Anions
Cations
2.
ammonium dichromate orange
nickel(II) sulphate green
The solutions of potassium chloride, potassium nitrate, potassium sulphate, sodium chloride, sodium nitrate, sodium sulphate, ammonium chloride, ammonium nitrate, ammonium sulphate, zinc chloride, zinc nitrate, zinc sulphate, calcium chloride and calcium nitrate. Colourless. Colourless; Colourless
©Aristo Educational Press Ltd. 2003
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Colourless; Colourless Colourless Colourless; Colourless
3.
Colourless a. Blue b.
Green
c.
Brown
d.
Green
Cations. Since the anions of these compounds are colourless (known from Step 2), it may be deduced that the colours of the solutions are due to the cations. Blue; Green Brown; Green
4.
Purple. Colourless. Purple.
5.
Yellow; Orange
©Aristo Educational Press Ltd. 2003
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Experiment 8.4
Migration of coloured ions (S/T)
1.
b.
The sodium sulphate solution provides more ions to conduct electricity. (If tap water were used, a smaller current and hence slower migration of ions would result.)
2.
b.
Purple colour. Positive electrode. No. Because the purple colour moves towards only one side of the paper.
3.
colourless; purple; mobile; negatively; permanganate; purple; positively; potassium; cannot be seen
5.
e.
Deep green.
6.
c.
Yellow colour. Blue colour.
7.
blue; yellow; negatively; chromate; yellow; positively; copper(II); blue
©Aristo Educational Press Ltd. 2003
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Experiment 8.5 1.
90º.
2.
6
To build a lattice model of sodium chloride
6 3.
Cubic.
4.
Yes. No. No. Giant ionic structure.
©Aristo Educational Press Ltd. 2003
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Experiment 9.1 2.
To build models of diamond and quartz
4 No. Giant covalent structure.
5.
No. Giant covalent structure.
©Aristo Educational Press Ltd. 2003
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Experiment 10.1
To compare the properties of sodium chloride, dry ice, wax and quartz
6. Substance
Physical state
Does it melt easily?
Is it volatile?
Is it soluble in water?
Does its aqueous solution conduct electricity?
solid
no
no
yes
yes
yes
(The solid disappears; a lot of gas bubbles are given out)
no
Sodium Chloride Dry ice
solid (when just (It sublimes at taken out from room vacuum flask), conditions) soon changes to gas
Wax
solid
yes
no
no
Quartz
solid
no
no
no
7.
a.
i. Sodium chloride. ii. Dry ice, wax. iii. Quartz.
b. Substance
Type of structure
Sodium chloride
giant ionic structure
Dry ice
simple molecular structure
Wax
simple molecular structure
Quartz
giant covalent structure
©Aristo Educational Press Ltd. 2003
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8.
a.
They are high-melting solids, non-volatile and soluble in water. They conduct electricity when dissolved in water and when molten.
b.
They have low melting points and are insoluble in water. They do not conduct electricity whether solid or liquid.
c.
They are high-melting solids, insoluble in water.
©Aristo Educational Press Ltd. 2003
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Experiment 11.1 1.
b.
To extract metals from metal oxides (T/S)
brownish black black white yellow
5.
Only on heating silver oxide. Silver oxide. 2Ag2O(s) → 4Ag(s) + O2(g) Silver. Only silver oxide decomposes on heating, while the oxides of the other metals do not. This shows that silver forms the least stable oxide and is thus least reactive.
9. Metal oxide
10.
Signs of formation of a metal (if any)
Copper(II) oxide
reddish brown solid
Zinc oxide
none
Lead(II) oxide
grey solid
Does reduction occur?
×
Only for copper(II) oxide and lead(II) oxide. 2CuO(s) + C(s) → 2Cu(s) + CO2(g); 2PbO(s) + C(s) → 2Pb(s) + CO2(g)
©Aristo Educational Press Ltd. 2003
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11.
Zinc oxide, lead(II) oxide, copper(II) oxide, silver oxide OR zinc oxide, copper(II) oxide, lead(II) oxide, silver oxide.
12.
Zinc, lead, copper, silver OR zinc, copper, lead, silver.
©Aristo Educational Press Ltd. 2003
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Experiment 12.1
To arrange five metals in order of reactivity
1.
b.
After a short time, the calcium burns vigorously with a brick-red flame. A white powder is left.
2.
It does not burn even when red hot. The surface of iron becomes black. It does not burn even when strongly heated. The surface of copper becomes black on strong heating. It burns with a dazzling white flame, forming a white powder.
3.
b.
Calcium granules sink to the bottom of the tube. Colourless gas bubbles are given out at a moderate rate. A milky suspension eventually forms. The tube gets warm.
4.
b.
No reaction. No reaction. Very slow reaction. Tiny gas bubbles are given out very slowly from the metal surface. No reaction.
5.
c.
There is rapid effervescence of a colourless gas. The calcium granule soon dissolves. The tube gets warm quickly.
6.
Evolution of colourless gas bubbles at a moderate rate from the metal surface only heating. No reaction even on heating. There is effervescence of a colourless gas. The ribbon dissolves rapidly. Slow evolution of tiny gas bubbles from the metal surface only on heating. A white solid precipitate is formed.
©Aristo Educational Press Ltd. 2003
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on
7. Metal
Does the metal burn on strong heating?
8.
Calcium
Iron
×
Copper
×
Magnesium
Lead
×
vigorous
vigorous
react with water?
very rapid
×
moderate (on heating)
×
×
×
moderate
react with dilute hydrochloric acid?
very slow
rapid
slow (on heating)
Calcium, magnesium, iron, lead, copper.
©Aristo Educational Press Ltd. 2003
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Experiment 12.2
Displacement reactions of metals in aqueous solutions
8. Solution
Cu2+ (aq)
Mg2+ (aq)
Zn2+ (aq)
Fe2+ (aq)
Ag+ (aq)
×
×
×
Metal Cu Mg
Zn
×
Fe
×
×
Ag
×
×
×
×
9.
Magnesium, zinc, iron, copper, silver.
10.
Consider two metals magnesium and zinc. Magnesium is higher than zinc in the metal reactivity series. This means that magnesium loses electrons more readily than zinc. Thus magnesium atoms lose electrons to become magnesium ions, while zinc ions are forced to gain electrons to become zinc atoms. Thus magnesium displaces zinc metal from zinc sulphate solution. On the other hand, there is no reaction between zinc metal and magnesium sulphate solution. Other displacement reactions can be explained similarly.
11.
Cu(s) + 2Ag+ (aq) → Cu2+(aq) + 2Ag(s)
12.
higher; lower; salt; more; less; less; more
©Aristo Educational Press Ltd. 2003
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Experiment 13.1
To determine the empirical formula of magnesium oxide (T/S) (Extension)
1.
28.094 g
2.
b.
3.
28.303 g
5.
This lets in oxygen from air to react with magnesium. If the lid was not lifted up, the small amount of oxygen inside the crucible would soon be used up in burning.
6.
The white smoke is magnesium oxide. If some escapes, the result will be inaccurate (the mass of oxygen found would be lower than it should be).
9.
White.
Dull grey.
Magnesium oxide. magnesium + oxygen → magnesium oxide 11.
28.439 g
12.
0.345 0.209 0.136 Mg
O
Masses (in g)
0.209
0.136
Number of moles of atoms (mol)
0.209
0.136
24.3
16.0
©Aristo Educational Press Ltd. 2003
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Relative number of moles
= 0.00860
= 0.00850
0.00860
0.00850
0.00850
0.00850
= 1.01
=1
≒ 1
13.
magnesium oxide; one; MgO
©Aristo Educational Press Ltd. 2003
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Experiment 14.1 sample laboratory report Title:
To investigate factors that influence rusting
Purpose To investigate factors that influence rusting.
Apparatus and chemicals used • • • • •
Beaker (250 cm3) Boiling tubes (150 × 18 mm) in rack Thermometer (−10 2 110°C) Small self-adhesive labels Stopwatch sea • 6 clean iron nails (~5 cm long) • Magnesium ribbon (7 cm long)
• Copper wire (7 cm long) • Dilute hydrochloric acid (1 M) • Sea water / salt water (a solution of 30 g sodium chloride per dm3 may be taken as water) • Hot water (about 80°C) • Warm agar solution containing rust indicator
Chemical reactions involved 4Fe(s) + 3O2 (g) + 2nH2O (l) → 2Fe2O3 • nH2O(s)
Procedure To investigate factors affecting rusting
In order to show rusting, a warm rust indicator solution (pale yellow in colour) was used. It contained agar and turned into a gel on cooling. 1. Six iron nails were cleaned, degreased and put into separate test tubes. 2. The six test tubes were labelled and set up as shown in Figure 1: Test tube 1: A clean iron nail was put into the rust indicator solution. (This is the control.) Test tube 2: A length of magnesium ribbon was used to wrap tightly around a clean iron nail and put into the rust indicator solution. Test tube 3: A length of copper wire was used to wrap tightly around a clean iron nail and put into the rust indicator solution. Test tube 4: A clean iron nail was put into the rust indicator solution mixed with 1 cm3 sea water. Test tube 5: A clean iron nail was put into the rust indicator solution mixed with 1 cm3 dilute HCl. Test tube 6: A clean iron nail was put into the rust indicator solution which was immersed into a hot water bath (~80°C). 3. All test tubes were left to stand for 20 minutes. They were then observed carefully, especially for the appearance of blue colour. ©Aristo Educational Press Ltd. 2003
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4.
The rate of rusting could be roughly determined by comparing • the time for the appearance of blue colour around the differently treated iron nails. • the size of the blue patches around the differently treated iron nails.
¡C)
Observation 1. 2. 3.
4. 5.
6.
Test tube 1: Blue colour mainly appeared around the head and tip. Test tube 2: No blue colour appeared. Some colourless gas bubbles were given off around the magnesium ribbon. Test tube 3: Blue colour appeared along the whole length of the wrapped iron nail. When compared to test tube 1, the time for the appearance of blue colour was shorter and the size of the blue patch is bigger. Test tube 4: Blue colour appeared along the iron nail. When compared to test tube 1, the time for the appearance of blue colour was shorter and the size of the blue patch is bigger. Test tube 5: Blue colour appeared quite rapidly along the iron nail. When compared to test tube 1, the time for the appearance of blue colour was shorter and the size of the blue patch is bigger. Some colourless gas bubbles were given off around the iron nail. Test tube 6: Blue colour appeared along the iron nail. When compared to test tube 1, the time for the appearance of blue colour was shorter and the size of the blue patch is bigger.
Interpretation 1.
Test tube 1: Rusting occurred mainly around the head and tip. This was a control experiment.
©Aristo Educational Press Ltd. 2003
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2.
3.
4. 5.
6.
Test tube 2: No rusting occurred. This was because magnesium, being more reactive, lost electrons more easily than iron. Since magnesium lost electrons to iron, iron was prevented from losing electrons and could not form Fe2+(aq) ions. Thus, the rate of rusting was slowed down or the iron nail was even being protected from rusting. Test tube 3: Rusting occurred quite quickly. This was because copper, being less reactive, encouraged the lost of electrons from iron to form Fe2+(aq) ions. Thus, the rate of rusting was increased. Test tube 4: Rusting occurred quite quickly. The presence of soluble salts such as sodium chloride could speed up rusting. Test tube 5: Rusting occurred most easily and quickly. Iron reacted with acid to form Fe2+(aq) ions. Fe(s) + 2H+(aq) → Fe2+(aq) + H2(g) Thus, the rate of rusting was increased in the presence of acid. Test tube 6: Rusting occurred quite quickly. An increase in temperature always increased the rate of chemical reaction including rusting.
Discussion 1.
2. 3.
Clean iron nails should be used. Sand paper may be needed to remove any surface coating. An aqueous solution of detergent or some acetone can be used to degrease the iron nails before the experiment. Magnesium ribbon and the copper wire should be used to wrap around the iron nail tightly. This was to make sure a good contact between the metal and iron nail. Sea water and dilute HCl could be added with stirring to test tubes 5 and 6 respectively before adding the iron nail.
Conclusion 1. 2.
Air and water are essential for rusting to occur. The presence of acidic substances, soluble salts, high temperature, uneven or sharply pointed regions (i.e. head and tip of nails) and the contact with a less reactive metal are all factors speeding up the rate of rusting.
Answers to questions for further thought 1. 2. 3.
Sodium chloride. Cheap, abundant. Rusting should be faster in Hong Kong where the humidity is high.
©Aristo Educational Press Ltd. 2003
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Experiment 14.2 3.
To prevent rusting
A blue colour appears, mainly around the head and tip of the nail. Gas bubbles appear around the magnesium ribbon. No blue colour appears. Nail ‘1’ only.
4.
Rusting occurs in nail ‘1’, but not in nail ‘2’. This is because magnesium, being more reactive than iron, loses electrons more easily. Since magnesium loses electrons to iron, iron is prevented from losing electrons and cannot form Fe2+(aq) ions. Iron is thus protected from rusting.
5.
To prevent rusting, connect the iron piece to a more reactive metal. This method is effective because the other metal will give up electrons in preference, preventing the formation of Fe2+ ions (sacrificial protection).
7.
A blue colour appears, mainly around the head and tip of the nail. No observation. Nail ‘3’ only.
8.
Rusting occurs in nail ‘3’ but not in nail ‘4’. This is because both water and air are kept out from painted iron, thus it is protected from rusting.
9.
Yes, both water and air are kept out from the nail covered with grease.
10.
sacrificial protection; paint
©Aristo Educational Press Ltd. 2003
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