Topic
2
Microscopic World I /
Microscopic World (Combined Science)
Part A Unit-based exercise
Unit 5 Atomic structure
Fill in the blanks
1
atoms
2
solids; liquids; gases
3
metals; metalloids; non-metals
4
low
5
good
6
Boron; silicon; germanium
7
nucleus; electrons; shells
8
protons; neutrons
9
atomic
10 mass
11 isotopes
12 atomic mass
13 electronic arrangement
14 electron
15 orbital
True or false
16 F
Mercury is a metal that exists as a liquid at room temperature and pressure.
17 T
18 T
19 F
28
Molten sulphur contains mobile molecules only. There are no mobile ions or electrons. It does not
conduct electricity.
20 F
The symbol of magnesium is Mg.
21 T The atomic number of an element is the number of protons in an atom of that element.
An atom has equal numbers of protons and electrons. Therefore the atomic number of an element also
equals the number of electrons in an atom of that element.
22 F
The neutral atom of an element must contain equal numbers of protons and electrons (NOT neutrons).
23 F
1
A hydrogen atom (1H) contains no neutron.
24 F
A sodium atom (2113Na) contains 11 protons, 11 electrons and 12 neutrons.
25 T The atomic number of fluorine is 9. A fluorine atom contains 9 protons and 9 electrons.
26 T Isotopes are different atoms of an element which have the same number of protons but a different
number of neutrons.
27 F
Isotopes of an element have different number of neutrons and hence they have different masses.
28 T
29 F
The second electron shell can hold a maximum of 8 electrons.
30 F
The electronic arrangement of a calcium atom is 2,8,8,2. Hence a calcium atom contains 4 occupied
electrons shells.
Multiple choice questions
31 D
Option
Element
Symbol
A
Calcium
Ca
B
Chlorine
Cl
C
Iron
Fe
D
Potassium
K
32 C Carbon, iron and silicon are solids at room temperature and pressure.
33 D
Option
Symbol
Element
State at room temperature and pressure
A
Cl
chlorine
gas
B
N
nitrogen
gas
C
Ne
neon
gas
D
S
sulphur
solid
34 B Beryllium is a solid at room temperature and pressure.
29
35 A Option B — X melts at 1 245 °C and boils at 1 869 °C. It is a solid at 25 °C.
Option C — Y melts at –58 °C and boils at 37 °C. It is a liquid at 25 °C.
Option D — Z melts at 52 °C and boils at 114 °C. It is a solid at 25 °C.
36 B Option A — W melts at –189 °C and boils at –186 °C. It is a gas at –100 °C and 1 atm pressure.
Option B — X melts at –110 °C and boils at –40 °C. It is a liquid at –100 °C and 1 atm pressure.
Option C — Y melts at –7 °C and boils at 60 °C. It is a solid at –100 °C and 1 atm pressure.
Option D — Z melts at –90 °C and boils at 10 °C. It is a solid at –100 °C and 1 atm pressure.
37 C
Substance
Melting point (°C)
Boiling point (°C)
State at room temperature and pressure
W
–50
5
gas
X
4
81
liquid
Y
68
104
solid
Z
–95
69
liquid
X and Z are in liquid state at room temperature and pressure.
38 A Carbon and neon are non-metals. Germanium is a metalloid.
39 A Copper is a metal. Helium and phosphorus are non-metals.
40 D
Option
Symbol
Element
Metal / non-metal
A
Ba
barium
metal
B
Be
beryllium
metal
C
Cs
caesium
metal
D
Kr
krypton
non-metal
41 C Metals are good conductors of electricity and insoluble in water (i.e. Y).
42 D Sillicon is insoluble in water.
43 D Option A — An atom must have equal numbers of protons and electrons (NOT neutrons).
Option B — The mass of one proton is approximately equal to that of 1 840 electrons.
Option C — A neutron carries no charge.
44 A Atomic number of X = 10 = number of protons in an atom
= number of electrons in an atom
Number of neutrons = mass number – atomic number
= 22 – 10
= 12
30
45 B An atom has equal numbers of protons and electrons. Therefore the atom has 28 protons and thus its
atomic number is 28.
Mass number = number of protons + number of neutrons
= 28 + 30
= 58
46 C Atomic number = 23 = number of protons
= number of electrons
Number of neutrons = mass number – atomic number
= 51 – 23 = 28
47 D
Atomic number
Mass number
Number of neutrons
= mass number – atomic number
16
32
16
23
11Na
11
23
12
24
12
24
12
28
14Si
14
28
14
31
15
31
16
Particle
32
16S
12Mg
15P
∴
31
15
P contains the same number of neutrons as
32
16
S.
48 C Isotopes are different atoms of an element which have the same number of protons and electrons, but
a different number of neutrons.
49 D
Atom
Atomic number
Mass number
Number of neutrons
I
17
37
20
II
19
39
20
III
20
40
20
IV
19
41
22
∴ atoms II and IV are isotopes of potassium.
6 x 7.4 + 7 x 92.6
100
= 6.93
50 B Relative atomic mass of lithium =
85 x 72.1 + 87 x 27.9
100
= 85.6
51 B Relative atomic mass of X =
189 x 25 + 190 x 30 + 192 x 45
100
= 190.7
52 C Relative atomic mass of X =
31
53 D Let the relative abundance of
79
X and
X be (100 – y)% and y% respectively.
81
79 x (100 – y) + 81 x y
100
∴ 7 900 – 79y + 81y = 7 990
Relative atomic mass of X =
= 79.9
y = 45
54 D Most elements have more than one isotope and the different isotopes of each element have different
masses. The relative isotopic mass and relative abundance of each isotope of an element in nature
must be considered when calculating the relative atomic mass of the element.
55 B The relative abundance of
X and aX are 65.0% and 35.0% respectively.
69
69 x 65 + a x 35
100
∴ 4 485 + 35a = 6 970
Relative atomic mass of X =
= 69.7
a = 71
56 B An atom of X has 15 electrons. The first 10 electrons fill up the first and second electron shells while
the last 5 electrons go into the third shell.
57 C
Option
Electronic arrangement
of atom
Atomic number
of element
Name of element
Metal / metalloid /
non-metal
A
2,1
3
lithium
metal
B
2,2
4
beryllium
metal
C
2,3
5
boron
metalloid
D
2,4
6
carbon
non-metal
∴ 2,3 represents the electronic arrangement of an atom of a metalloid.
58 C
Option
Symbol
Element
Electronic arrangement of atom
A
Cl
chlorine
2,8,7
B
P
phosphorus
2,8,5
C
S
sulphur
2,8,6
D
Si
silicon
2,8,4
∴ the question shows the electron diagram of an atom of sulphur (symbol S).
59 B (1) Mercury is a metal which is a liquid at room temperature and pressure.
(3) Not all metals are stored in paraffin oil, e.g. magnesium and copper are not stored in paraffin oil.
60 D
61 A (2) Gallium is a metal, not a metalloid.
(3) The crystalline form of silicon (a metalloid) can conduct electricity at room temperature.
32
62 A (2) Isotopes of an element have the same number of protons. Hence they have the same atomic
number.
(3) Isotopes of an element have different number of neutrons and hence different masses.
63 A (1) P is an atom of phosphorus (a non-metal).
(2) Q is an atom of sulphur (a non-metal).
(3) P and Q are atoms of different elements. They are NOT isotopes.
64 B (1) Particles X and Z have different number of protons but the same number of neutrons. Hence they
have different masses.
(2) Particles X and Y have equal numbers of protons and electrons, but a different number of neutrons.
Hence they are isotopes.
(3) Particles Y and Z have different number of electrons. Hence they have different electronic
arrangements.
65 B (1) X is phosphorus. It is a solid at room temperature and pressure.
(2) The electronic arrangement of an atom of X is 2,8,5. Hence there are 5 electrons in the outermost
shell of an atom of X.
(3) Number of neutrons in an atom of X = mass number – atomic number
= 31 – 15
= 16
66 A (1) X is sodium. Hence it is a metal.
(2) The electronic arrangement of an atom of X is 2,8,1. Hence there are 11 electrons and 11 protons
in an atom of X.
(3) X is sodium. Its symbol is Na.
67 A (1) Number of neutrons in a
(2) The atomic number of
60
27
60
27
Co atom = mass number – atomic number
= 60 – 27
= 33
Co is 27. Hence a
60
27
Co atom contains 27 protons.
(3) Isotopes of an element have the same number of electrons.
68 B Both statements are true. However, the state of an element cannot be explained by the fact whether it
is a metal or non-metal.
69 A
70 C An atom is electrically neutral because it has equal numbers of protons and electrons (NOT neutrons).
71 B The atomic number of an element is the number of protons in an atom of the element. A
contains 16 protons.
72 C
54
24
X and
32
16
S atom
54
26
Y represent different elements. They are NOT isotopes.
33
73 A
74 A Most elements have more than one isotope and the different isotopes of each element have different
masses. The relative isotopic mass and relative abundance of each isotope of an element in nature
must be considered when calculating the relative atomic mass of the element.
75 C Isotopes of an element have the same number of protons but a different number of neutrons. Hence
they have different masses.
Unit 6 The periodic table
Fill in the blanks
1
atomic number
2
groups; periods
3
group
4
period
5
metalloids; non-metals
6
alkali
7
paraffin oil
8
increases
9
hydrogen; sodium hydroxide
10 alkaline earth
11 halogens
12 greenish yellow; reddish brown; black
13 decreases
14 noble gases
15 octet
16 Argon
17 positive
18 negative
19 2; 12; 10
20 3; 7; 10
34
True or false
21 F
In the periodic table, all the elements are arranged in order of increasing atomic number.
22 T
23 T Across the second period, the elements change from metals through metalloids to non-metals.
24 F
Atoms of elements in the same period have the same number of occupied electron shells. The atom of
each element in Period 3 has 3 occupied electron shells.
25 F
The electronic arrangement of a sulphur atom is 2,8,6. A sulphur atom has 3 occupied electron shells.
Hence sulphur belongs to Period 3 of the periodic table.
26 T The electronic arrangement of an aluminium atom is 2,8,3. An aluminium atom has 3 outermost shell
electrons. Hence aluminium is a Group III element.
27 T Sodium is a Group I element, an alkali metal.
28 F
Argon belongs to Group 0. It is a noble gas.
29 T Magnesium is a Group II element, an alkaline earth metal.
30 T Neon is a noble gas and belongs to Group 0 of the periodic table.
31 T Potassium is a reactive metal and must be stored in paraffin oil to prevent it from reacting with the air.
32 F
The melting point of the Group I elements decreases down the group.
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Hence the melting point of sodium is lower than that of lithium.
33 F
Beryllium and calcium belong to the same group (Group II). They have similar chemical properties, NOT
the same chemical properties.
34 T
35 F
Iodine vapour is purple in colour.
35
36 T
37 F
35
38 F
A helium atom has a duplet structure in its outermost shell.
Cl and
37
Cl have the same electronic arrangement and hence the same chemical properties.
39 T
40 T An oxygen atom has an electronic arrangement 2,6. It tends to gain 2 electrons in order to obtain the
stable electronic arrangement of a neon atom (2,8). An oxide ion forms.
Multiple choice questions
41 D Option A — Elements in the periodic table are arranged in order of increasing atomic number.
Option B — The vertical columns are called groups.
Option C — The horizontal rows are called periods.
42 C Option A — The number of occupied electron shells in atoms of elements in the same group increases
down the group.
e.g.
Group I element
Electronic arrangement
of atom
Number of occupied
electron shells
Lithium
2,1
2
Sodium
2,8,1
3
Potassium
2,8,8,1
4
Option B — The atomic number of elements in the same group increases down the group.
Option D — Elements in the same group have similar, not the same, chemical properties.
43 C Option A — The reactivity of Period 2 elements changes across the period. Apart from the noble
gases, the most reactive elements are near the edges of the periodic table and the least
reactive ones are in the centre.
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Option B — The atom of each element in Period 2 has two occupied electron shells, NOT 2 outermost
shell electrons.
Option D — Across the period, the elements change from metals through metalloids to non-metals.
44 B The electronic arrangement of a carbon atom is 2,4. A carbon atom has 4 outermost shell electrons.
Hence carbon belongs to Group IV of the periodic table.
45 C The group number of an element equals the number of outermost shell electrons in its atom.
An atom of element X has 6 outermost shell electrons. Hence X belongs to Group VI of the periodic
table.
46 A An atom of X has an electronic arrangement 2,8. X is a noble gas and belongs to Group 0 of the
periodic table.
47 C An atom of element X has an electronic arrangement 2,8,5. It has 5 outermost shell electrons. Hence
X belongs to Group V of the periodic table.
48 B
49 A
Option
Element
Group number
A
Boron
III
B
Bromine
VII
C
Chlorine
VII
D
Silicon
IV
Option
Element
Group
argon
0
neon
0
carbon
IV
chlorine
VII
calcium
II
potassium
I
magnesium
II
sodium
I
A
B
C
D
Argon and neon belong to the same group.
37
50 D Lithium and sodium belong to Group I of the periodic table.
Element
Atomic number
Electronic arrangement of atom
Lithium
3
2,1
Sodium
11
2,8,1
The chemical properties of an element depend on the number of outermost shell electrons in its atom.
Lithium and sodium have the same number of outermost shell electrons in their atoms. Hence they
have similar chemical properties.
51 B
Option
Atomic number of element
Electronic arrangement of atom
Group to which the
element belongs
7
2,5
V
13
2,8,3
III
9
2,7
VII
17
2,8,7
VII
11
2,8,1
I
18
2,8,8
0
14
2,8,4
IV
20
2,8,8,2
II
A
B
C
D
Elements with atomic numbers 9 and 17 belong to the same group. Hence they have similar chemical
properties.
52 D Aluminium belongs to Group III of the periodic table.
Element
Atomic number
B
5
Al
13
Difference in atomic number
8
Electronic arrangement of atom
2,3
2,8,3
18
Ga
31
2,8,18,3
The difference in atomic number between the first three successive Group III elements is either 8 or 18.
This is because the second electron shell can hold 8 electrons while the third electron shell can hold 18
electrons.
Hence the atomic number of an element belonging to the same group as aluminium could be 13 + 18, i.e.
31.
53 B The atomic number of sulphur is 16. The electronic arrangement of an atom of sulphur is 2,8,6. An
atom of sulphur has 3 occupied electron shells. Hence sulphur belongs to Period 3 of the periodic
table.
38
54 A
Option
Element
Atomic
number
Electronic arrangement
of atom
Period to which the
element belongs
argon
18
2,8,8
3
aluminium
13
2,8,3
3
beryllium
4
2,2
2
silicon
14
2,8,4
3
chlorine
17
2,8,7
3
nitrogen
7
2,5
2
phosphorus
15
2,8,5
3
oxygen
8
2,6
2
A
B
C
D
Argon and aluminium belong to the same period.
55 B Option C — Elements in the same group have the same number of outermost shell electrons in their
atoms.
Option D — The reactivity of elements changes across a period. Apart from the noble gases, the most
reactive elements are near the edges of the periodic table and the least reactive ones are
in the centre.
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56 D Option A — Element X is fluorine, a non-metal.
Option B — The electronic arrangement of an atom of X is 2,7. An atom of X has 2 occupied
electron shells. Hence X belongs to Period 2 of the periodic table.
Option C — Fluorine is a gas at room temperature and pressure.
Option D — X belongs to Group VII of the periodic table.
39
57 C Option A — The reactivity of Group I elements increases down the group.
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Option B — Sodium gives a golden yellow flame in flame test.
Option C — The density of sodium is lower than that of water. Hence it floats on water.
Option D — Sodium reacts with water to form hydrogen gas.
58 A The melting point of Group I elements decreases down the group, i.e. decreases with increasing atomic
number.
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59 B Option A — Chlorine is a greenish yellow gas.
Option B — Chlorine belongs to Group VII of the periodic table. Its atom has 7 outermost shell
electrons.
Option C — Sodium chloride is used to manufacture chlorine. Electrolysis of concentrated sodium
chloride solution gives chlorine, hydrogen and sodium hydroxide.
Option D — The reactivity of Group VII elements decreases down the group.
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Hence fluorine is more reactive than chlorine.
60 D
Option
Element
Atomic number
Name of element
Group to which the
element belongs
A
W
4
beryllium
II
B
X
11
sodium
I
C
Y
12
magnesium
II
D
Z
19
potassium
I
Group II elements are less reactive than Group I elements. The reactivity of Group I elements increases
down the group.
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61 A X conducts electricity and reacts readily with dilute hydrochloric acid to give hydrogen. X is probably a
metal.
62 A Option A — The atomic number of calcium is 20. The electronic arrangement of an atom of calcium
is 2,8,8,2. The atom has 4 occupied electron shells. Hence calcium is in Period 4 of the
periodic table.
Option B — Calcium is a Group II element, i.e. an alkaline earth metal.
Option C — Calcium reacts with non-metals to form salts.
Option D — Calcium reacts steadily with cold water, but does NOT catch fire.
63 C The melting point of Group VII elements increases down the group, i.e. increases with atomic number.
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64 C Option A —
Element
Colour
Chlorine
greenish yellow
Bromine
reddish brown
Iodine
black
Chlorine, bromine and iodine are all coloured substances.
Option C — The reactivity of Group VII elements decreases down the group, i.e. decreases with
increasing relative atomic mass.
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Option D — Chlorine, bromine and iodine can react with sodium sulphite solution.
e.g. aqueous bromine + sodium sulphite
sodium sulphate + hydrogen bromide
65 B Argon is used to fill electric light bulbs because it does not react with the metal filament in the light
bulb.
66 C Option C — A helium atom has 2 outermost shell electrons.
Option D — The boiling point of noble gases increases down the group.
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67 A
Noble gas
Atomic number
He
2
Ne
10
Difference in atomic number
8
Electronic arrangement of atom
2
2,8
8
Ar
42
Kr
18
36
2,8,8
18
2,8,18,8
The difference in atomic number between the first four successive noble gases is either 8 or 18. This
is because the second electron shell can hold 8 electrons while the third electron shell can hold 18
electrons.
Therefore if the atomic number of A is x, then the atomic number of B could be x – 8.
68 C The reactivity of elements changes across a period. Apart from the noble gases, the most reactive
elements are near the edges of the periodic table and the least reactive ones are in the centre. Group
I elements are the most reactive metals.
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metal.
69 D Element j is bromine, which exists as a liquid at room temperature and pressure.
70 A Group I elements are metals. They are good conductors of electricity. Their melting points and densities
are much lower than the average values for metals.
71 D
72 C
Option
Description
A
Group I elements have relatively low melting points.
B
Atoms of Group I elements have 1 outermost shell electron.
C
Group I elements are relatively soft.
D
The reactivity of Group I elements increases down the group. Hence rubidium is more reactive
than potassium.
Option
Description
A
Strontium and calcium are Group II elements. They are reactive metals. They are extracted from
their ores by electrolysis.
B
Strontium is a Group II element. Hence its atom has 2 outermost shell electrons.
C
Strontium is a reactive metal. It will react with oxygen in the air to form an oxide layer on the
surface. Hence strontium tarnishes when exposed to the air.
D
The densities of Group II elements are higher than that of water. Hence strontium sinks in water.
43
73 B Option A — Krypton belongs to Group 0 of the periodic table. The atomic number of krypton is 36.
The electronic arrangement of an atom of krypton is 2,8,18,8. Hence an atom of krypton
has an octet structure in its outermost shell.
Option B — The electronic arrangement of an atom of krypton is 2,8,18,8. A krypton atom has 4
occupied electron shells. Hence krypton belongs to Period 4 of the periodic table.
Option C — All Group 0 elements are colourless gases at room temperature and pressure.
Option D — The density of Group 0 elements increases down the group. The density of krypton is
higher than that of air. Hence a balloon full of krypton falls in the air.
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74 B Option B — A magnesium atom has an electronic arrangement 2,8,2. It tends to lose two electrons in
order to obtain the stable electronic arrangement of a neon atom (2,8). A magnesium ion
forms.
Options C and D — Losing electrons does not affect the number of protons and neutrons the
magnesium atom has. Hence its atomic number and mass number remain the
same.
75 D
Option
Species
Electronic arrangement
Number of electrons
Li+
2
2
H
1
1
2,8
10
2,8,8
18
2,8
10
2,8,8
18
2,8
10
2,8
10
A
2–
O
B
Cl–
+
Na
C
2–
S
Ne
D
–
F
Ne and F– have the same number of electrons.
76 C
Number of
Species
protons
neutrons
electrons
26Fe
26
30
26
3+
26
30
23
56
Fe
44
77 D An atom of X loses two electrons to form the X2+ ion with an electronic arrangement 2,8,8. Hence the
electronic arrangement of an atom of X is 2,8,8,2. The atomic number of X is 20. It is calcium.
Option
Description
A, B and C
D
X is a Group II element, an alkaline earth metal.
An atom of X has 4 occupied electron shells. Hence X is a Period 4 element.
78 B An anion carrying 1 negative charge has one more negative charge than positive charge.
The anion has 36 electrons and hence it should have 35 protons.
79 D As M2+ ion has two more positive charges than negative charges, it should have 30 protons. Hence an
atom of M should have 30 protons as well. The atomic number of M is 30. M is zinc (Zn).
80 C Particle X is an oxide ion O2–. Particle Y is an isotope of oxygen.
81 B (1) Across the second period of the periodic table, the elements show a gradual decrease in atomic
size.
Group
I
II
III
IV
Element
Lithium
Beryllium
Boron
Atomic radius (pm)
–12
(1 pm = 10 m)
152
112
83
IV
V
Carbon
Carbon
Nitrogen
(graphite) (diamond)
77
75
VI
VII
Oxygen
Fluorine
73
72
(2) Across the second period, the elements change from metals through metalloids to non-metals.
(3) Across the second period, the melting point of elements rises to Group IV and then falls to low
values.
Group
I
II
III
IV
Element
Lithium
Beryllium
Boron
Melting point (°C)
180
1 280
2 030
IV
V
Carbon
Carbon
Nitrogen
(graphite) (diamond)
3 730
3 500
–210
VI
VII
Oxygen
Fluorine
–218
–220
82 A (1) Elements in the same group show a gradual increase in relative atomic mass.
Group IV element
Symbol
Relative atomic mass
Carbon
C
12.0
Silicon
Si
28.1
Germanium
Ge
72.6
Tin
Sn
118.7
45
(2) The number of occupied electron shells in atoms of Group IV elements increases down the group.
Hence the atomic size of the elements increases down the group.
Group IV element
Symbol
Atomic number
Electronic arrangement of atom
Carbon
C
6
2,4
Silicon
Si
14
2,8,4
Germanium
Ge
32
2,8,18,4
Tin
Sn
50
2,8,18,18,4
(3) All Group IV elements have 4 outermost shell electrons in their atoms.
83 A Element X is argon. An atom of X has an electronic arrangement 2,8,8.
(1) An atom of X has three occupied electron shells. Hence X is in Period 3 of the periodic table.
(2) An atom of X has 8 outermost shell electrons. Hence X is in Group 0 of the period table.
(3)
40
18
X has 22 neutrons.
84 A (1) A lithium atom has 2 occupied electron shells, a sodium atom has 3 while a potassium atom has 4.
Hence the atomic size is in the order lithium < sodium < potassium.
(2) The chemical reactivity of Group I elements increases down the group.
MJUIJVN
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TPEJVN
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Hence the chemical reactivity is in the order lithium < sodium < potassium.
(3) The melting point of Group I elements decreases down the group.
MJUIJVN
NFMUJOHQPJOU
TPEJVN
EFDSFBTJOHEPXO
UIFHSPVQ
QPUBTTJVN
Hence the melting point is in the order lithium > sodium > potassium.
46
85 A (2) The densities of Group II elements are higher than that of water. Hence they sink in water.
(3) Group II elements can react with water. Therefore they are NOT stored in water.
86 D (1) The reactivity of halogens decreases down the group, i.e. decreases with increasing relative atomic
mass.
3FMBUJWFBUPNJDNBTT
GMVPSJOF
DIMPSJOF
CSPNJOF
JPEJOF
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UIFHSPVQ
87 B
Element
Atomic number
Name of element
Metal / non-metal
X
11
sodium
metal
Y
16
sulphur
non-metal
Z
17
chlorine
non-metal
(1) All the three elements belong to the third period of the periodic table.
(2) X is an alkali metal.
88 A (2) Helium is unreactive because it has a duplet structure in the outermost shell of its atom.
89 C (1) A helium atom has 2 outermost shell electrons while other noble gas atoms have 8 electrons in
their outermost shells.
(2)
Noble gas
Symbol
Relative atomic mass
Helium
He
4.0
Neon
Ne
20.2
Argon
Ar
39.9
Krypton
Kr
83.8
Xenon
Xe
131.3
Radon
Rn
222
The relative atomic mass of noble gases increases down the group.
90 D Element X is magnesium.
(1) Magnesium is a reactive metal. It will react with oxygen in the air to form an oxide layer on the
surface. Hence magnesium tarnishes when exposed to the air.
(2) The density of magnesium is higher than that of water. Hence it sinks in water.
47
91 D (1) Caesium is an alkali metal. It reacts with water to give an alkaline solution.
(2) Caesium gives a characteristic flame colour in flame test, just as other alkali metals, such as sodium
and potassium, do.
(3) The melting point of Group I elements decreases down the group.
MJUIJVN
TPEJVN
QPUBTTJVN
SVCJEJVN
DBFTJVN
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UIFHSPVQ
GSBODJVN
The melting point of sodium is less than 100 °C. Hence that of caesium should be less than
100 °C as well.
92 C (1) Strontium is a Group II element. There are 2 outermost shell electrons in a strontium atom. Hence
a strontium atom tends to lose 2 electrons in order to obtain a stable electronic arrangement. An
ion carrying two positive charges is formed.
(2) The reactivity of Group II elements increases down the group.
CFSZMMJVN
NBHOFTJVN
SFBDUJWJUZ
DBMDJVN
EPXOUIFHSPVQ
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TUSPOUJVN
Hence strontium is more reactive than calcium.
93 C (1) The reactivity of halogens decreases down the group.
GMVPSJOF
DIMPSJOF
SFBDUJWJUZ
EFDSFBTJOHEPXO
UIFHSPVQ
CSPNJOF
JPEJOF
Hence the reactivity of halogens is in the order chlorine > bromine > iodine.
48
(2)
Halogen
State at room temperature
Density at 25 °C
Chlorine
gas
0.00321
Bromine
liquid
3.12
Iodine
solid
4.93
Hence the density of the halogens is in the order chlorine < bromine < iodine.
(3)
Halogen
Electronic arrangement of atom
Atomic radius (pm) (1 pm = 10–12 m)
Chlorine
2,8,7
99
Bromine
2,8,18,7
114
Iodine
2,8,18,18,7
133
The number of occupied electron shells in atoms of halogens increases down the group. Hence the
atomic size of halogens is in the order chlorine < bromine < iodine.
94 D (1) There is a gradual increase in the colour intensity of elements down Group VII. Astatine is probably
coloured.
(3) There is a gradual change in state of elements down Group VII, from gas to liquid then to solid.
Astatine is probably a solid at room temperature and pressure.
95 D X is a Group II element as there are 2 outermost shell electrons in its atom.
X is barium, which is below calcium and strontium in Group II of the periodic table.
(1) The densities of Group II elements are higher than that of water. Hence X is denser than water.
(2) X (barium) gives a characteristic flame colour in flame test as other Group II elements, such as
calcium, do.
(3) The reactivity of Group II elements increases down the group.
CFSZMMJVN
NBHOFTJVN
DBMDJVN
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EPXOUIFHSPVQ
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As calcium reacts with dilute hydrochloric acid, X (barium) can probably react with dilute
hydrochloric acid as well.
96 A A hydrogen atom has 1 proton and 1 electron. It forms a H+ ion upon losing 1 electron.
)
[H]+ + e–
Hence a H+ ion has 1 proton only.
49
97 A (1) An oxygen atom has an electronic arrangement 2,6. It gains 2 electrons to form an oxide ion (with
an electronic arrangement 2,8). Hence an oxygen atom and an oxide ion have the same number of
occupied electron shells.
m
Q
O
Q
O
m
F
PYZHFOBUPN
PYJEFJPO
(2) and (3) Gaining electrons does not affect the number of protons and neutrons the oxygen atom
has. Hence its atomic number and mass number remain the same.
98 A
Species
Electronic arrangement of species
O2–
2,8
Li+
2
K+
2,8,8
Hence the species O2– has the same electronic arrangement as a neon atom.
99 B Atom X can form a stable ion X–. It can be deduced that atom X has 7 outermost shell electrons.
Hence X is a halogen.
e.g.
m
Q
O
DIMPSJOFBUPN
m
F
Q
O
DIMPSJEFJPO
(1) and (2) Gaining an electron does not affect the number of protons and neutrons X has. Hence ion X–
and atom X have the same number of neutrons and nuclear charge.
(3) Atom X gains 1 electron to obtain an octet structure in its outermost shell. Ion X– and atom X
have the same number of occupied electron shells.
100 A X is a Group VII element, Y is a noble gas and Z is a Group I element.
(1) Atom of X gains 1 electron to form a stable anion X – . Anion X – has the same electronic
arrangement as atom Y.
Atom of Z loses 1 electron to form a stable cation Z + . Cation Z + has the same electronic
arrangement as atom Y.
Hence anion X– and cation Z+ have the same electronic arrangement.
(2) X and Z are different elements. Atoms of X and Z have different number of protons. Hence anion X–
and cation Z+ have different number of protons.
(3) X, Y and Z belong to different periods of the periodic table.
50
101 B
102 A
Element
Electronic arrangement of atom
Lithium
2,1
Neon
2,8
Atoms of lithium and neon have 2 occupied electron shells. Hence they belong to Period 2 of the
periodic table.
103 B
Element
Electronic arrangement of atom
Nitrogen
2,5
Oxygen
2,6
Atoms of nitrogen and oxygen have 2 occupied electron shells. Hence nitrogen and oxygen belong to
the same period of the periodic table.
104 D Across the second period of the periodic table, the melting point of elements rises to Group IV and
then falls to low values.
.FMUJOHQPJOUž$
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*7
/
7
0
7*
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7**
/F
Elements in the second period of the periodic table have 2 occupied electron shells in their atoms.
51
105 B Across the third period of the periodic table, the atomic size of the elements decreases gradually.
1FSJPE
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/B
.H
"M
4J
4
1
$M "S
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Across the third period of the periodic table, the elements change from metals through metalloids to
non-metals, i.e. the metallic character of the elements decreases.
Group
I
II
III
IV
V
VI
VII
Period 3 element
Sodium
Magnesium
Aluminium
Silicon
Phosphorus
Sulphur
Chlorine
metals
metalloid
non-metals
Type of element
metallic character decreasing
106 A The chemical properties of an element depend on the number of outermost shell electrons in its atom.
Atoms with the same number of outermost shell electrons react in a similar way. Sodium and
potassium belong to the same group. Both of them have 1 outermost shell electron in their atoms.
Therefore sodium and potassium have similar chemical properties.
107 D The density of Group II elements show a gradual increase down the group (except magnesium and
calcium).
Element
Density (g cm–3)
Beryllium
1.85
Magnesium
1.74
Calcium
1.55
Strontium
2.60
Barium
3.51
The reactivity of Group II elements increases down the group.
52
CFSZMMJVN
NBHOFTJVN
SFBDUJWJUZ
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DBMDJVN
EPXOUIFHSPVQ
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108 C The reactivity of halogens decreases down the group.
GMVPSJOF
DIMPSJOF
SFBDUJWJUZ
EFDSFBTJOHEPXO
UIFHSPVQ
CSPNJOF
JPEJOF
The number of occupied electron shells in atoms of halogens increases down the group. Hence the
atomic size also increases down the group.
–12
Halogen
Electronic arrangement of atom
Atomic radius (pm) (1 pm = 10
Fluorine
2,7
72
Chlorine
2,8,7
99
Bromine
2,8,18,7
114
Iodine
2,8,18,18,7
133
m)
109 A
110 C An argon atom has 8 electrons in its outermost shell. It is unreactive. Its chemical properties are
different from that of a chloride ion.
53
Unit 7 Ionic and metallic bonds
Fill in the blanks
1
a) conductors
b) aqueous solution; electrolytes
c) non-conductors
2
ionic
3
negative; positive
4
polyatomic
5
four; two
6
three; two
7
purple; permanganate
8
chromium(III)
9
chemical formula
10 metallic
True or false
11 F
An ionic bond is the strong forces of attraction between oppositely charged ions.
12 F
Ionic bond usually occurs when metal atoms combine with non-metal atoms.
13 T A calcium atom has an electronic arrangement 2,8,8,2. It tends to lose 2 electrons to obtain the
electronic arrangement of a stable argon atom.
A fluorine atom has an electronic arrangement 2,7. It tends to gain 1 electron to obtain the electronic
arrangement of a stable neon atom.
When calcium and fluorine react, the 2 electrons released by the calcium atom are accepted by two
fluorine atoms.
m
'
'
$B
GMVSPJOFBUPN
$B
GMVPSJEFJPO
m
'
DBMDJVNBUPN
DBMDJVNJPO
'
GMVSPJOFBUPN
GMVPSJEFJPO
54
14 F
Element X (atomic number = 11) is sodium. A sodium atom has an electronic arrangement 2,8,1. It
tends to lose 1 electron to obtain the electronic arrangement of a stable neon atom.
Element Y (atomic number = 16) is sulphur. A sulphur atom has an electronic arrangement 2,8,6. It tends
to gain 2 electrons to obtain the electronic arrangement of a stable argon atom.
When sodium and sulphur combine, two sodium atoms are required to release the 2 electrons needed
by the sulphur atom.
/B
/B
m
TPEJVNJPO
TPEJVNBUPN
4
4
/B
TVMQIVSBUPN
/B
TVMQIJEFJPO
TPEJVNJPO
TPEJVNBUPN
Hence the chemical formula of the compound formed is Na2S, i.e. X2Y.
15 T A magnesium atom has an electronic arrangement 2,8,2. It tends to lose 2 electrons to obtain the
electronic arrangement of a stable neon atom.
An oxygen atom has an electronic arrangement 2,6. It tends to gain 2 electrons to obtain the electronic
arrangement of a stable neon atom.
When magnesium and oxygen react, the 2 electrons released by the magnesium atom are accepted by
one oxygen atom.
m
.H
0
.H
0
NBHOFTJVNBUPN
PYZHFOBUPN
NBHOFTJVNJPO
PYJEFJPO
Hence the chemical formula of the compound formed is MgO.
16 T Element X (atomic number = 20) is calcium. A calcium atom has an electronic arrangement 2,8,8,2. It
tends to lose 2 electrons to obtain the electronic arrangement of a stable argon atom.
Element Y (atomic number = 7) is nitrogen. A nitrogen atom has an electronic arrangement 2,5. It tends
to gain 3 electrons to obtain the electronic arrangement of a stable neon atom.
When calcium and nitrogen react, three calcium atoms are required to release the 6 electrons needed by
the two nitrogen atoms.
55
Electron diagram of the compound formed:
m
$B
Ca
/
/
$B
$B
m
/
/
$B
$B
(Only electrons in the outermost shells are shown.)
The chemical formula of the compound formed is Ca3N2, i.e. X3Y2.
17 F
Potassium is a Group I element. Its atom tends to lose 1 electron to obtain the electronic arrangement
of a stable noble gas atom.
Astatine is a Group VII element. Its atom tends to gain 1 electron to obtain the electronic arrangement
of a stable noble gas atom.
When potassium and astatine combine, the electron released by the potassium atom is accepted by the
astatine atom.
,
"U
,
m
"U
(Only electrons in the outermost shells are shown.)
The chemical formula of the compound formed is KAt.
18 F
Iron(II) ion in aqueous solution is pale green in colour.
19 T The orange colour of a potassium dichromate solution comes from the dichromate ions.
20 F
56
Metallic bond is a type of bond in which positive metal ions are held together by a ‘sea’ of mobile
electrons.
Multiple choice questions
21 A Alcohol is made up of carbon, hydrogen and oxygen. Compounds made up of non-metals are nonconductors.
22 C Electrode X is the positive electrode while electrode Y is the negative electrode.
Bromide ions carrying negative charges move towards the positive electrode.
bromide ions – electrons
bromine atoms
bromine molecules
Lead(II) ions carrying positive charges move towards the negative electrode.
lead(II) ions + electrons
lead atoms
Hence a reddish brown gas (bromine) is formed at the positive electrode, i.e. electrode X. A white
shiny solid (lead) is formed at the negative electrode, i.e. electrode Y.
23 A Option A — Molten lead(II) bromide is decomposed into lead and bromine by electricity.
Options B, C and D — In solid state, ions in the compound are held together by strong attraction.
They are not free to move. Hence solid lead(II) bromide does not conduct
electricity.
When lead(II) bromide becomes molten, the lead(II) ions and bromide ions
become mobile. Hence molten lead(II) bromide can conduct electricity.
24 C Potassium (a metal) and oxygen (a non-metal) combine to form an ionic compound.
25 C
Element
Atomic number
Name of element
a
6
carbon
b
9
fluorine
c
10
neon
d
11
sodium
Element b (fluorine, a non-metal) and element d (sodium, a metal) combine to form an ionic
compound.
26 B Element d is a non-metal and element b is a metal. They combine to form an ionic compound.
27 A Metallic bonds are found in metals only, e.g. copper.
57
28 C Electron transfer during the reaction between potassium and oxygen:
,
,
m
QPUBTTJVNBUPN
QPUBTTJVNJPO
0
0
PYJEFJPO
PYZHFOBUPN
,
,
QPUBTTJVNJPO
QPUBTTJVNBUPN
29 B Element X forms a stable X2+ ion. Its atom probably has 2 outermost shell electrons. Thus X is probably
a Group II element.
Element Y forms a stable Y2– ion. Its atom probably has 6 outermost shell electrons. Thus Y is probably
a Group VI element.
Option B — Calcium is a Group II element while sulphur is a Group VI element.
Electron transfer during the reaction between calcium and sulphur:
m
$B
4
$B
4
DBMDJVNBUPN
TVMQIVSBUPN
DBMDJVNJPO
TVMQIJEFJPO
(Only electrons in the outermost shells are shown.)
30 D Element X is magnesium. It belongs to Group II of the periodic table. An atom of X loses 2 electrons
to obtain a stable electronic arrangement.
Elements X and Y form an ionic compound with the chemical formula XY2. That means the 2 electrons
released by an atom of X are accepted by two atoms of Y. Hence it can be deduced that an atom of
Y needs 1 more electron to obtain a stable electronic arrangement.
For atoms of non-metals in Group V, VI and VII, they gain ‘8 – group number’ electrons in order
to obtain stable electronic arrangements. An atom of Y gains 1 electron in order to obtain a stable
electronic arrangement. Hence Y probably belongs to Group VII of the periodic table.
58
Electron transfer during the reaction between X and Y:
m
:
:
9
9
m
:
:
(Only electrons in the outermost shells are shown.)
31 C Element X is sodium. It belongs to Group I of the periodic table. An atom of X loses 1 electron to
obtain a stable electronic arrangement.
Elements X and Y form an ionic compound with the chemical formula X2Y. That means the 2 electrons
released by two atoms of X are accepted by one atom of Y. Hence it can be deduced that an atom of
Y needs 2 more electrons to obtain a stable electronic arrangement.
For atoms of non-metals in Group V, VI and VII, they gain ‘8 – group number’ electrons in order to
obtain stable electronic arrangements. Hence Y probably belongs to Group VI of the periodic table.
Electron transfer during the reaction between X and Y:
9
9
m
:
:
9
9
(Only electrons in the outermost shells are shown.)
32 C An atom of X loses 2 electrons to form a stable X2+ ion. Hence the atom probably has 2 electrons in
its outermost shell.
An atom of Y gains 3 electrons to form a stable Y3– ion. Hence the atom probably has 5 electrons in
its outermost shell.
59
33 B
Option
Compound
Chemical formula
Electron diagram
A
aluminium oxide
Al2O3
The chemical formula is
not in the form of XY.
—
B
magnesium oxide
.H
MgO
m
0
C
lithium fluoride
-J
LiF
m
'
D
sodium chloride
/B
NaCl
m
$M
∴ the positive ions and the negative ions in MgO have the same electronic arrangement.
34 D A magnesium atom has an electronic arrangement 2,8,2. It tends to lose 2 electrons to obtain the
electronic arrangement of a stable neon atom.
Element X (atomic number = 7) is nitrogen. A nitrogen atom has an electronic arrangement 2,5. It
tends to gain 3 electrons to obtain the electronic arrangement of a stable neon atom.
When magnesium and nitrogen react, three magnesium atoms are required to release the 6 electrons
needed by the two nitrogen atoms.
Electron diagram of the compound formed:
.H
m
/
.H
m
/
.H
(Only electrons in the outermost shells are shown.)
The chemical formula of the compound formed is Mg3N2, i.e. Mg3X2.
60
35 C Element X is aluminium. An atom of X has an electronic arrangement 2,8,3. It tends to lose 3 electrons
to obtain the electronic arrangement of a stable neon atom.
Element Y is oxygen. An atom of Y has an electronic arrangement 2,6. It tends to gain 2 electrons to
obtain the electronic arrangement of a stable neon atom.
When X and Y react, two atoms of X are required to release the 6 electrons needed by the three atoms of Y.
Electron diagram of the compound formed:
m
0
"M
m
0
"M
m
0
(Only electrons in the outermost shells are shown.)
The chemical formula of the compound formed is Al2O3, i.e. X2Y3.
36 B An atom of X loses 1 electron to form a stable X+ ion. It can be deduced that the atom of X has 1
electron in its outermost shell. Hence element X is probably a Group I element.
An atom of Y gains 3 electrons to form a stable Y3– ion. It can be deduced that the atom of Y has 5
electrons in its outermost shell. Hence element Y is probably a Group V element.
X could be Li (lithium) and Y could be N (nitrogen), i.e. Option B is the correct answer.
37 D The reactivity of Group I elements increases down the group. Hence potassium is more reactive than
lithium.
MJUIJVN
SFBDUJWJUZ
JODSFBTJOH
TPEJVN
EPXO
UIF
HSPVQ
QPUBTTJVN
61
The reactivity of Group VII elements decreases down the group. Hence fluorine is more reactive than
chlorine.
GMVPSJOF
SFBDUJWJUZ
DIMPSJOF
EFDSFBTJOHEPXO
UIFHSPVQ
CSPNJOF
JPEJOF
Therefore fluorine and potassium (Option D) would react with each other most vigorously.
38 B
Element
Atomic number
Name of element
w
9
fluorine
x
14
silicon
y
18
argon
z
20
calcium
Element w (fluorine) is a reactive non-metal while element z (calcium) is a reactive metal. They react
with each other readily to form an ionic compound.
39 A The chemical formula of ammonium ion is NH4+.
40 C Nickel(II) ion in aqueous solution is green in colour.
41 A
42 C
43 D
Ion
Colour in aqueous solution
Cr3+
green
Cu2+
blue or green
Fe3+
yellow-brown
Zn2+
colourless
44 B The aqueous solution of compound XZ is colourless. Hence X2+(aq) and Z2–(aq) ions are colourless.
The blue colour of aqueous solution of compound WZ is due to the W2+(aq) ions. Hence W2+(aq) ion is
blue in colour.
The orange colour of aqueous solution of compound XY is due to the Y2–(aq) ions. Hence Y2–(aq) ion is
orange in colour.
62
45 D A coloured patch develops near the positive electrode because negative permanganate ions which are
purple in colour move towards the positive electrode.
Positive potassium ions move towards the negative electrode. However, we cannot see the potassium
ions because they are colourless.
46 A The chemical formula of potassium dichromate is K2Cr2O7.
47 B The chemical formula of ammonium sulphate is (NH4)2SO4. The compound consists of 4 elements.
48 C A rubidium ion (Rb+) carries 1 positive charge. A carbonate ion (CO32–) carries 2 negative charges. The
simplest ratio of Rb+ and CO32– in rubidium carbonate is 2 : 1. Hence the chemical formula of the
compound is Rb2CO3.
DIBSHF
3C
m
3C
$0m
2+
49 D A calcium ion (Ca ) carries 2 positive charges while a phosphate ion (PO43–) carries 3 negative charges.
The simplest ratio of Ca2+ to PO43– in calcium phosphate is 3 : 2. Hence the chemical formula of the
compound is Ca3(PO4)2.
DIBSHF
$B
$B
m
$B
10m 10m
50 B M forms a sulphate with the chemical formula M2(SO4)3. The sulphate ion (SO42–) carries 2 negative
charges and the net charge of the compound must be zero. The simplest ratio of ion of M to sulphate
ion in the sulphate is 2 : 3. It can be deduced that the ion of M carries 3 positive charges.
Use the following steps to work out the chemical formula of the chloride of M:
Step
Chloride of M
1 Write down the symbols of ions in the compound.
M
Cl
2 Write down the number of charges of each ion on the top
of each symbol.
3
M
1
Cl
3 Cross multiply the numbers and the symbols.
4 Combine the symbols.
3
M
= M1
1
Cl
= Cl3
MCl3
63
51 D Thorium forms a hydroxide with the chemical formula Th(OH)4. The hydroxide ion (OH–) carries 1
negative charge and the net charge of the compound must be zero. It can be deduced that the
thorium ion carries 4 positive charges.
The permanganate ion (MnO4–) also carries 1 negative charge. Hence the chemical formula of thorium
permanganate is Th(MnO4)4.
52 D Option A — Strontium is a Group II element, an alkaline earth metal.
Option B — The reactivity of Group II elements increases down the group.
CFSZMMJVN
NBHOFTJVN
SFBDUJWJUZ
JODSFBTJOH
DBMDJVN
EPXOUIFHSPVQ
TUSPOUJVN
Hence strontium is more reactive than calcium.
Option C — A strontium ion (Sr2+) carries 2 positive charges while a chloride ion (Cl–) carries 1 negative
charge. The simplest ratio of Sr2+ to Cl– in strontium chloride should be 1 : 2. Hence the
chemical formula of the compound is SrCl2.
DIBSHF
4S
m
$Mm
$Mm
Option D — The densities of Group II elements are higher than 1 g cm–3 while that of sodium is
lower. Hence the density of strontium is higher than that of sodium.
53 A Option A — There is a gradual change in state of elements down Group VII.
Fluorine and chlorine are gases, bromine is a liquid, and iodine and astatine are solids at
room temperature and pressure.
Option B — The reactivity of Group VII elements decreases down the group.
GMVPSJOF
DIMPSJOF
SFBDUJWJUZ
EFDSFBTJOHEPXO
CSPNJOF
UIFHSPVQ
JPEJOF
BTUBUJOF
64
Hence chlorine is more reactive than astatine.
Option C — Astatine is in Period 6 of the periodic table.
2+
–
Option D — A calcium ion (Ca ) carries 2 positive charges while the ion of astatine (At ) carries 1
negative charge. The simplest ratio of Ca2+ to At– in the compound formed between
calcium and astatine should be 1 : 2. Hence the chemical formula of the compound is
CaAt2.
DIBSHF
m
$B
"Um
"Um
54 C Cation X + has an electronic arrangement 2,8. Hence an atom of X should have an electronic
arrangement 2,8,1. Element X is sodium.
Option A — X (sodium) is a solid at room temperature and pressure.
Option B — An atom of X has 3 occupied electron shells. Hence X is in Period 3 of the periodic table.
Option C — X (sodium) reacts vigorously with water to give sodium hydroxide and hydrogen.
+
2–
Option D — The cation of X (X ) carries 1 positive charge while the oxide ion (O ) carries 2 negative
charges. The simplest ratio of X+ to O2– in the oxide of X should be 2 : 1. Hence the
chemical formula of the oxide is X2O.
DIBSHF
9
m
9
0m
55 A (1) Calcium (a metal) and fluorine (a non-metal) combine to form an ionic compound.
56 B Cation X2+ has an electronic arrangement 2,8,8. Hence an atom of X has an electronic arrangement 2,8,8,2.
Element X is calcium.
(1) A calcium ion (Ca2+) carries 2 positive charges while a hydride ion (H–) carries 1 negative charge.
The simplest ratio of Ca2+ to H– in the compound formed between calcium and hydrogen should be
1 : 2. Hence the chemical formula of the compound is CaH2, i.e. XH2.
Electron diagram of the compound formed:
m
)
$B
m
)
(Only electrons in the outermost shells are shown.)
(2) An atom of X has 4 occupied electron shells. Hence X is in Period 4 of the periodic table.
65
57 A (1) The reactivity of Group I elements increases down the group.
MJUIJVN
TPEJVN
QPUBTTJVN
SFBDUJWJUZ
JODSFBTJOH
EPXOUIFHSPVQ
SVCJEJVN
DBFTJVN
As potassium reacts vigorously with water, caesium should also react with water vigorously.
(2) The densities of all Group I elements are quite low.
(3) A caesium ion (Cs+) carries 1 positive charge while a hydroxide ion (OH–) carries 1 negative charge.
The simplest ratio of Cs+ to OH– in caesium hydroxide should be 1 : 1. Hence the chemical formula
of the compound is CsOH.
58 D (1) Oxides of Group I elements are crystalline solids.
(2) Metallic bonding exists in metals.
(3) A rubidium ion (Rb+) carries 1 positive charge while a sulphate ion (SO42–) carries 2 negative
charges. The simplest ratio of Rb+ to SO42– in rubidium sulphate should be 2 : 1. Hence the
chemical formula of the compound is Rb2SO4.
DIBSHF
3C
m
3C
40m
59 B (1) A barium atom has 2 outermost shell electrons. It forms a stable ion by losing 2 electrons. Hence
the ion carries 2 positive charges.
(2) The reactivity of Group II elements increases down the group.
CFSZMMJVN
NBHOFTJVN
DBMDJVN
SFBDUJWJUZ
JODSFBTJOH
EPXOUIFHSPVQ
TUSPOUJVN
CBSJVN
Hence barium is more reactive than calcium.
66
(3) A barium ion (Ba2+) carries 2 positive charges while a sulphide ion (S2–) carries 2 negative charges.
The simplest ratio of Ba2+ to S2– in barium sulphide should be 1 : 1. Hence the chemical formula of
the sulphide is BaS.
2+
2–
60 B (1) A strontium ion (Sr ) carries 2 positive charges while a carbonate ion (CO3 ) carries 2 negative
2+
2–
charges. The simplest ratio of Sr to CO3 in strontium carbonate should be 1 : 1. Hence the
chemical formula of the compound is SrCO3.
(2) In solid state, ions in strontium carbonate are held together by strong attraction. They are not free
to move. Hence solid strontium carbonate does not conduct electricity.
(3) Carbonates of Group II elements are insoluble in water, e.g. calcium carbonate. Hence strontium
carbonate is insoluble in water.
61 B Element X (atomic number = 7) is nitrogen.
(1) X (nitrogen) is a gas at room temperature and pressure.
(2) An atom of X has an electronic arrangement 2,5. It has 2 occupied electron shells. Hence X is in
Period 2 of the periodic table.
(3) A magnesium atom has an electronic arrangement 2,8,2. It tends to lose 2 electrons to obtain the
electronic arrangement of a stable neon atom.
Element X (atomic number = 7) is nitrogen. A nitrogen atom has an electronic arrangement 2,5. It
tends to gain 3 electrons to obtain the electronic arrangement of a stable neon atom.
When magnesium and nitrogen react, three magnesium atoms are required to release the 6
electrons needed by the two nitrogen atoms.
Electron diagram of the compound formed:
.H
m
/
.H
m
/
.H
(Only electrons in the outermost shells are shown.)
The chemical formula of the compound is Mg3N2, i.e. Mg3X2.
67
62 D
Compound
Chemical formula
Electron diagram
Lithium oxide
Li2O
-J
-J
,
K2S
,
/B
Na2O
m
4
Sodium oxide
0
Potassium sulphide
m
/B
m
0
∴ the positive ions and negative ions in potassium sulphide and sodium oxide have the same
electronic arrangement.
63 A
Ion
Chemical formula
Colour in aqueous solution
Chromium(III)
Cr3+
green
Nickel(II)
Ni2+
green
Permanganate
MnO4–
purple
64 C (1) Electrode X is the positive electrode. An orange colour develops near electrode X. This is because
negative dichromate ions move towards the positive electrode.
(2) Electrode Y is the negative electrode. Hydrogen ions (H+) around electrode Y will gain electrons to
form hydrogen gas.
65 A (3) Metallic bonding exists in metals only.
66 A An oxygen atom has an electronic arrangement 2,6. It tends to gain 2 electrons in order to obtain the
stable electronic arrangement of a neon atom (2,8).
m
Q
O
m
F
PYZHFOBUPN
67 C A helium atom has 2 outermost shell electrons.
68
Q
O
PYJEFJPO
68 C A magnesium atom has an electronic arrangement 2,8,2. It has 3 occupied electron shells. Hence
magnesium is in Period 3 of the periodic table.
69 A Element X (atomic number = 20) is calcium. A calcium atom has an electronic arrangement 2,8,8,2. It
tends to lose 2 electrons to obtain the electronic arrangement of a stable argon atom.
Element Y (atomic number = 7) is nitrogen. A nitrogen atom has an electronic arrangement 2,5. It
tends to gain 3 electrons to obtain the electronic arrangement of a stable neon atom.
When calcium and nitrogen react, three calcium atoms are required to release the 6 electrons needed
by the two nitrogen atoms
Electron diagram of the compound formed:
$B
m
/
$B
m
/
$B
(Only electrons in the outermost shells are shown.)
The chemical formula of the compound is Ca3N2, i.e. X3Y2.
70 D Some transition metals (e.g. Fe) can form M3+ ions.
Unit 8 Covalent bonds
Fill in the blanks
1
covalent
2
bond pair
3
lone pair
4
diatomic
5
monoatomic
6
one
7
three; one
8
three
9
dative covalent
10 nitrogen; lone pair
69
True or false
11 T
12 F
Nitrogen exists as diatomic molecules.
/
/
/
/
OJUSPHFONPMFDVMF/≡/
(Only electrons in the outermost shells are shown.)
13 T Carbon and silicon are non-metals. They combine to form a covalent compound (silicon carbide, SiC).
Structure of silicon carbide:
LFZ
DBSCPOBUPN
TJMJDPOBUPN
14 F
Neon is a noble gas. A neon atom has 8 outermost shell electrons. A special stability is obtained when
this happens. Hence neon will NOT form a compound with nitrogen.
15 F
Electron diagram of water:
)
0
)
Total number of electrons in a water molecule
= number of electrons in one oxygen atom + 2 x number of electron in one hydrogen atom
=8 + 2 x 1
= 10
70
16 T In a SiCl4 molecule, one Si atom forms a single bond with each of four Cl atoms.
Electron diagram of SiCl4:
$M
$M
$M
4J
$M
(Only electrons in the outermost shells are shown.)
17 T Covalent bonding exists in both iodine and oxygen.
*
*
0
JPEJOFNPMFDVMF*m*
0
PYZHFONPMFDVMF00
(Only electrons in the outermost shells are shown.)
18 F
Covalent bonding exists in hydrogen chloride while ionic bonding exists in silver chloride.
19 F
A phosphorus atom has an electronic arrangement 2,8,5. It needs 3 more electrons to obtain the
electronic arrangement of a stable argon atom (2,8,8).
A hydrogen atom has an electronic arrangement 1. It needs 1 more electron to obtain the electronic
arrangement of a stable helium atom (2).
In order to obtain stable electronic arrangements, one phosphorus atom forms a single bond with each
of three hydrogen atoms. The chemical formula of the compound is PH3.
Electron diagram of PH3:
)
1
)
)
(Only electrons in the outermost shells are shown.)
71
20 T A dative covalent bond is formed when an ammonia molecule and a hydrogen ion combine to form
an ammonium ion (NH4+). The nitrogen atom in the ammonia molecule supplies its lone pair electrons
to the hydrogen ion.
)
)
)
/
)
)
)
/
)
)
Hence covalent bonds exist in ammonium chloride.
Multiple choice questions
21 D Element X (atomic number = 7) is nitrogen.
A nitrogen atom has an electronic arrangement 2,5. It needs 3 more electrons to obtain the electronic
arrangement of a stable neon atom (2,8). Each nitrogen atom can obtain the electronic arrangement of
a neon atom by sharing three of its electrons with another nitrogen atom.
22 A An atom of an element with atomic number 9 has an electronic arrangement 2,7. It needs 1 more
electron to obtain the electronic arrangement of a stable neon atom (2,8). Each atom can obtain the
electronic arrangement of a neon atom by sharing one of its outermost shell electrons with another
atom.
'
'
(Only electrons in the outermost shells are shown.)
23 D Sulphur and oxygen are non-metals. They combine to form a covalent compound.
24 D Option A — Lithium (a metal) and nitrogen (a non-metal) combine to form an ionic compound.
Option B — Mercury (a metal) and fluorine (a non-metal) combine to form an ionic compound.
Option C — Neon is a noble gas. A neon atom has 8 outermost shell electrons. A special stability is
obtained when this happens. Hence neon will NOT form a compound with nitrogen.
Option D — Fluorine and chlorine are non-metals. They combine to form a covalent compound.
Electron diagram of the compound formed:
$M
'
(Only electrons in the outermost shells are shown.)
72
25 C
Element
Atomic number
Name of element
a
3
lithium
b
14
silicon
c
17
chlorine
d
18
argon
Option A — Element a (lithium) and element c (chlorine) combine to form an ionic compound.
Option B — Argon (element d) is a noble gas. An argon atom has 8 outermost shell electrons. A
special stability is obtained when this happens. Hence argon will NOT form a compound
with lithium (element a).
Option C — Element b (silicon) and element c (chlorine) combine to form a covalent compound.
Option D — Element b (silicon) will NOT form a compound with element d (argon).
26 D The element exists as diatomic molecules. It is a non-metal. Nitrogen in Group V exists as diatomic
molecules.
27 C A nitrogen atom has an electronic arrangement 2,5. It needs 3 more electrons to obtain the electronic
arrangement of a stable neon atom (2,8).
A fluorine atom has an electronic arrangement 2,7. It needs 1 more electron to obtain the electronic
arrangement of a stable neon atom (2,8).
In order to get stable electronic arrangements, one nitrogen atom combines with three fluorine atoms
to form a molecule.
28 D Option A — Helium exists as monoatomic molecules.
Option B — In a fluorine molecule, each fluorine atom shares one of its outermost shell electrons with
another fluorine atom.
'
'
(Only electrons in the outermost shells are shown.)
Option C — In a hydrogen chloride molecule, one hydrogen atom forms a single bond with one
chlorine atom.
)
$M
(Only electrons in the outermost shells are shown.)
73
Option D — A silicon atom has an electronic arrangement 2,8,4. It needs 4 more electrons to obtain
the electronic arrangement of a stable argon atom (2,8,8).
A hydrogen atom has an electronic arrangement of 1. It needs 1 more electron to obtain
the electronic arrangement of a stable helium atom (2).
In order to get stable electronic arrangements, one silicon atom combines with four
hydrogen atoms.
29 C A neon atom has 8 electrons in its outermost shell. A special stability is obtained when this happens.
The atom has very little tendency to share electrons with other neon atoms. Hence neon exists as
atoms.
30 B Element X (atomic number = 16) is sulphur.
A sulphur atom has an electronic arrangement 2,8,6. It needs 2 more electrons to obtain the electronic
arrangement of a stable argon atom (2,8,8).
A chlorine atom has an electronic arrangement 2,8,7. It needs 1 more electron to obtain the electronic
arrangement of a stable argon atom (2,8,8).
In order to obtain stable electronic arrangements, one sulphur atom forms a single bond with each of
two chlorine atoms. The chemical formula of the compound is SCl2, i.e. XCl2.
Electron diagram of SCl2:
$M
4
$M
(Only electrons in the outermost shells are shown.)
31 C In the compound, each atom of X contributes 1 electron for sharing with another atom of X and 1
electron for sharing with an atom of Y.
Each atom of X contributes 2 electrons for sharing as it needs 2 electrons to obtain a stable electronic
arrangement. It can be deduced that an atom of X has 6 outermost shell electrons.
,FZ
:
9
9
:
FMFDUSPOT
JOUIFPVUFSNPTU
TIFMMPGBOBUPNPG9
32 D In the compound, the atom of X contributes 2 electrons for sharing with two atoms of Y. The atom of
X contributes 2 electrons for sharing as it needs 2 electrons to obtain a stable electronic arrangement.
It can be deduced that an atom of X has 6 outermost shell electrons. Element X belongs to Group VI
of the periodic table (i.e. oxygen in this question).
,FZ
:
74
9
:
FMFDUSPOT
JOUIFPVUFSNPTU
TIFMMPGBOBUPNPG9
Each atom of Y contributes 1 electron for sharing as it needs 1 electron to obtain a stable electronic
arrangement. It can be deduced that an atom of Y has 7 outermost shell electrons. Element Y belongs
to Group VII of the periodic table (i.e. fluorine in this question).
,FZ
:
9
:
FMFDUSPOT
JOUIFPVUFSNPTU
TIFMMPGBOBUPNPG:
33 C In the compound, each atom of X contributes 1 electron for sharing with another atom of X and 2
electrons for sharing with two hydrogen atoms.
Each atom of X contributes 3 electrons for sharing as it needs 3 electrons to obtain a stable electronic
arrangement. It can be deduced that an atom of X has 5 outermost shell electrons.
)
34 D
Molecule
)
)
9
9
,FZ
FMFDUSPOT
JOUIFPVUFSNPTU
TIFMMPGBOBUPNPG9
)
Electron diagram (only electrons in the
outermost shells are shown)
)
)
$
$
C2H4
6
)
CO2
)
0
)
H2S
$M
Number of pairs of bond pair electrons
$
0
4
4
)
2
1
PCl3
$M
3
$M
75
35 D
Molecule
Electron diagram (only electrons in the
outermost shells are shown)
Number of pairs of lone pair electrons on
the underlined atom
)
)
CH4
$
)
0
)
HCN
)
)
$
/
/
0
)
1
NH3
)
SCl2
$M
4
2
$M
36 A An atom of X has 5 outermost shell electrons. Hence X probably belongs to Group V of the periodic
table.
An atom of Y has 7 outermost shell electrons. Hence Y probably belongs to Group VII of the periodic
table.
An atom of X needs 3 more electrons to obtain a stable electronic arrangement. An atom of Y needs
1 more electron to obtain a stable electronic arrangement.
In order to obtain stable electronic arrangements, one atom of X forms a single covalent bond with
each of three atoms of Y. A covalent compound forms. The chemical formula of the compound is XY3.
Electron diagram of XY3:
:
9
:
(Only electrons in the outermost shells are shown.)
76
:
37 C An atom of X gains 3 electrons to form an anion X3– with an electronic arrangement 2,8. Hence the
electronic arrangement of an atom of X is 2,5. Element X is nitrogen.
Option A — X (nitrogen) belongs to Group V of the periodic table.
Option B — X (nitrogen) is a gas at room temperature and pressure.
Option C — X (nitrogen) exists as diatomic molecules.
/
/
(Only electrons in the outermost shells are shown.)
Option D — X (nitrogen) and fluorine are non-metals. They combine to form a covalent compound.
38 D
Compound formed between
Electron diagram (only electrons in the outermost shells are shown)
m
'
magnesium and fluorine
.H
m
'
m
-J
lithium and oxygen
0
-J
$M
chlorine and fluorine
chlorine and oxygen
$M
'
0
$M
77
39 B
Option
Substance
Bonding type
aluminium
A
mercury
metallic bonding
sodium
calcium chloride
ionic bonding
hydrogen chloride
covalent bonding
silver chloride
ionic bonding
C
carbon dioxide
nitrogen
oxygen
covalent bonding
D
iodine
methane
sulphur dioxide
covalent bonding
B
40 C In Al2X3, the aluminium ion carries 3 positive charges. The simplest ratio of Al3+ to the ion of X in the
compound is 2 : 3 and the net charge of the compound is zero. Hence the ion of X carries 2 negative
charges.
In Period 3 of the periodic table, sulphur forms an anion carrying 2 negative charges. Hence X is
sulphur.
To obtain stable electronic arrangements, one sulphur atom forms a single bond with each of two
hydrogen atoms. The chemical formula of the compound formed is H2S, i.e. H2X.
Electron diagram of H2S:
)
4
)
(Only electrons in the outermost shells are shown.)
41 D In the compound between element X and chlorine, the atom of X contributes 3 electrons for sharing
with 3 atoms of chlorine. The atom of X contributes 3 electrons for sharing as it needs 3 electrons to
obtain a stable electronic arrangement. Therefore it has 5 outermost shell electrons.
An atom of X forms a stable anion X3– by gaining 3 electrons. Use the following steps to work out the
chemical formula of the compound formed between X and calcium:
Step
1 Write down the symbols of ions in the compound.
Ca
X
2 Write down the number of charges of each ion on the top
of each symbol.
2
Ca
3
X
3 Cross multiply the numbers and the symbols.
4 Combine the symbols.
78
Compound formed between X
and calcium
2
Ca
= Ca3
3
X
= X2
Ca3X2
42 A Option B — Atoms of elements in the same period have the same number of occupied electron shells.
Option C — The melting point of Period 2 elements rises to Group IV and than falls to low values.
.FMUJOHQPJOUž$
m
&MFNFOU
(SPVQOVNCFS
-J
*
#F
**
#
***
$
*7
/
7
0
7*
'
7**
/F
Option D — Lithium and oxygen combine to form an ionic compound.
Fluorine and oxygen combine to form a covalent compound.
Across the second period, from lithium to fluorine, the elements change from metals to
non-metals. The oxides of the elements change from ionic to covalent.
Group
I
II
III
IV
V
VI
VII
Element
lithium
beryllium
boron
carbon
nitrogen
oxygen
fluorine
forms a
covalent
oxide
forms an
ionic oxide
43 D In a piece of metal, metal ions are packed tightly together in a regular pattern to form a giant
structure.
LFZ
NFUBMJPO
NPCJMF
FMFDUSPOT
FMFDUSPO
79
44 B Element X (atomic number = 9) is fluorine.
Element Y (atomic number = 13) is aluminium.
The chemical formula of the compound formed between X and Y is YX3 (or AlF3).
Formula mass of the compound = relative atomic mass of Y + 3 x relative atomic mass of X
= 27.0 + 3 x 19.0
= 84.0
45 D Element X (atomic number = 6) is carbon.
Element Y (atomic number = 16) is sulphur.
The chemical formula of the compound formed between X and Y is XY2 (or CS2).
Relative molecular mass of the compound = relative atomic mass of X + 2 x relative atomic mass of Y
= 12.0 + 2 x 32.1
= 76.2
46 C Phosphorus and oxygen are non-metals. They combine to form a covalent compound.
47 A Elements a, b, e and f are non-metals while elements c and d are metals.
(1) Element a (a non-metal) and element f (a non-metal) combine to form a covalent compound.
(2) Element b (a non-metal) and element d (a metal) combine to form an ionic compound.
(3) Element c (a metal) and element e (a non-metal) combine to form an ionic compound.
48 B
Option
Molecule
Electron diagram (only electrons in the
outermost shells are shown)
With single
bonds only?
)
(1)
CH4
)
$
)
yes
)
80
(2)
CO2
(3)
H2O
0
)
$
0
0
)
no
yes
49 D
Option
Molecule
(1)
CO2
(2)
N2
(3)
HCN
Electron diagram (only electrons in the
outermost shells are shown)
0
$
0
/
yes
/
$
)
With multiple
bond(s)?
yes
yes
/
50 D Chlorine and bromine are in Group VII of the periodic table.
(1) Each chlorine / bromine atom has 7 outermost shell electrons. It gains 1 electron to obtain a stable
electronic arrangement. An ion carrying 1 negative charge is formed.
m
$M#S
Fm
$M#S
(Only electrons in the outermost shells are shown.)
(2) Each chlorine / bromine atom can obtain a stable electronic arrangement by sharing one of its
outermost shell electrons with another chlorine / bromine atom. Hence chlorine / bromine exists as
diatomic molecules.
Electron diagram of a chlorine / bromine molecule:
$M#S
$M#S
(Only electrons in the outermost shells are shown.)
(3) Both chlorine and bromine can react with sodium sulphite solution.
e.g. aqueous bromine + sodium sulphite
sodium sulphate + hydrogen bromide
81
51 B In a methane (CH4) molecule, the carbon atom forms a single bond with each of the four hydrogen
atoms.
Electron diagram of methane:
)
)
$
)
)
(1) The number of bonding electrons contributed by each hydrogen atom in the molecule is 1.
(2) The number of bonding electrons contributed by the carbon atom in the molecule is 4.
(3) Total number of electrons in the molecule
= number of electrons in one carbon atom + 4 x number of electron in one hydrogen atom
= 6 + 4 x 1
= 10
52 B (2) Electron diagram of magnesium bromide:
m
#S
.H
m
#S
(Only electrons in the outermost shells are shown.)
53 A (1) In the compound, each atom of X contributes 1 electron for sharing with another atom of X and 1
electron for sharing with an atom of Y.
Each atom of X contributes 2 electrons for sharing as it needs 2 electrons to obtain a stable
electronic arrangement. It can be deduced that there are 6 electrons in the outermost shell of an
atom of X.
,FZ
:
82
9
9
:
FMFDUSPOT
JOUIFPVUFSNPTU
TIFMMPGBOBUPNPG9
(2) An atom of Y contributes 1 electron for sharing as it needs 1 electron to obtain a stable electronic
arrangement. It can be deduced that there are 7 electrons in the outermost shell of an atom of Y.
,FZ
:
9
9
:
FMFDUSPOT
JOUIFPVUFSNPTU
TIFMMPGBOBUPNPG:
(3) Each atom of X in the compound has 2 lone pairs of electrons.
MPOFQBJSPGFMFDUSPOT
:
9
9
:
MPOFQBJSPGFMFDUSPOT
54 B The electron diagram of ammonium chloride is shown in the question.
)
)
/
m
)
$M
)
(Only electrons in the outermost shells are shown.)
(2) There are 5 electrons in the outermost shell of an atom of Y (nitrogen).
55 A The electron diagram of a metal carbonate is shown in the question.
m
0
9
$
0
0
FMFDUSPOGSPNBO
FYUFSOBMTPVSDF
FMFDUSPOGSPNBO
FYUFSOBMTPVSDF
(Only electrons in the outermost shells are shown.)
(1) An atom of X loses 2 electrons to form a stable X2+ ion. Hence the atom probably has 2 electrons
in its outermost shell.
(2) There are 4 electrons in the outermost shell of an atom of Y (carbon).
(3) There are 6 electrons in the outermost shell of an atom of Z (oxygen).
83
56 A (1) Carbon combines with oxygen to form a compound with the chemical formula CO2.
(2) Lead combines with chlorine to form a compound with the chemical formula PbCl2.
(3) Lithium combines with oxygen to form a compound with the chemical formula Li2O.
57 A In ammonium nitrate (NH4NO3), the cations and anions are held together by ionic bonding, but each
polyatomic ion is a group of atoms held together by covalent bonding.
Electron diagram of an ammonium ion:
Electron diagram of a nitrate ion:
m
0
)
/
)
)
/
0
0
)
(Only electrons in the outermost shells are shown.)
58 D
Element
Atomic number
Name of element
Metal / non-metal
X
9
fluorine
non-metal
Y
12
magnesium
metal
Z
16
sulphur
non-metal
(1) X (a non-metal) and Y (a metal) react to give an ionic compound.
(2) X and Z are non-metals. They form a compound by electron sharing.
(3) Y (magnesium) and Z (sulphur) react to form a compound with the chemical formula YZ.
Electron diagram of the compound formed between Y and Z:
:
(Only electrons in the outermost shells are shown.)
84
m
;
59 D (1) Group II elements are less reactive than Group I elements.
(SPVQ (SPVQ
*
**
Be
r SFBDUJWJUZ
JODSFBTJOH
Na
Mg
K
r SFBDUJWJUZEFDSFBTJOH
Hence sodium is more reactive than magnesium.
(2) Phosphorus and chlorine are non-metals. They combine to form a covalent compound.
(3) Electron diagram of the compound formed between silicon and chlorine:
$M
$M
4J
$M
$M
(Only electrons in the outermost shells are shown.)
60 C (1) A Group IV element seldom forms an ion.
(3) X and Y are non-metals. They combine to form a covalent compound.
61 B Bromine and chlorine belong to the same group as their atoms have the same number of outermost
shell electrons.
62 C Neon is a noble gas. A neon atom has 8 outermost shell electrons. A special stability is obtained when
this happens. Hence neon will NOT form a compound with nitrogen.
63 B Phosphorus and chlorine form a covalent compound as they are both non-metals.
85
64 D Hydrogen and chlorine are non-metals. They combine to form a covalent compound.
Electron diagram of the compound formed between hydrogen and chlorine:
$M
)
(Only electrons in the outermost shells are shown.)
65 C Calcium carbonate is an ionic compound.
Each carbonate ion (CO32–) is a group of four atoms held together by covalent bonding.
m
0
$
0
0
FMFDUSPOGSPNBO
FYUFSOBMTPVSDF
FMFDUSPOGSPNBO
FYUFSOBMTPVSDF
Unit 9 Relating the properties of substances to structures and
bonding
Fill in the blanks
1
giant ionic
2
giant covalent
3
simple molecular
4
giant metallic
5
water; non-aqueous
6
mobile
7
Allotropes
8
oxygen; covalent; oxygen
9
three; covalent; Van der Waals’
10 water; non-aqueous
11 simple molecular; covalent bond; van der Waals’ forces
12 mobile electrons
86
True or false
13 F
Allotropes are two (or more) forms of the same element in which the atoms or molecules are arranged
in different ways.
Quartz is a form of silicon dioxide while graphite is a form of carbon.
Hence quartz and graphite are NOT allotropes.
14 T Both silicon and diamond have a giant covalent structure.
Structure of silicon:
LFZ
TJMJDPOBUPN
15 F
Silicon carbide has a giant covalent structure.
Structure of silicon carbide:
LFZ
DBSCPOBUPN
TJMJDPOBUPN
16 T Electron diagram of silane (SiH4):
)
)
4J
)
)
(Only electrons in the outermost shells are shown.)
87
17 F
Carbon disulphide (CS2) has a simple molecular structure.
Electron diagram of carbon disulphide:
4
$
4
(Only electrons in the outermost shells are shown.)
18 T The melting point of sugar is low and it is a non-conductor of electricity. It can be deduced that sugar
has a simple molecular structure.
19 T
20 F
Ionic compounds do not conduct electricity in solid state. In solid state, ions in the compound are held
together by strong ionic bonds. They are not free to move.
21 T Sodium chloride is hard due to the strong ionic bonds between oppositely charged ions. Relative
motion of the ions is restricted.
22 F
The melting point of graphite (3 730 °C) is higher than that of diamond (3 500 °C).
23 F
In quartz, each silicon atom is joined to four oxygen atoms by covalent bonds, while each oxygen
atom is joined to two silicon atoms by covalent bonds.
Structure of quartz:
LFZ
PYZHFOBUPN
TJMJDPOBUPN
24 T In graphite, the carbon atoms are arranged in flat parallel layers.
There are weak van der Waals’ forces between the adjacent layers in graphite. The layers can easily slide
over each other. Hence graphite has a slippery feel.
Structure of graphite:
WBOEFS8BBMTGPSDFT
LFZ
DBSCPOBUPN
88
25 T Diamond has a giant covalent structure consisting of a network of covalent bonds. Relative motion of
the atoms is restricted. This makes diamond very hard.
Diamond is harder than graphite. In graphite, the layers of carbon atoms are held by weak van der Waals’
forces. The layers can slide over each other easily.
26 T Graphite is a good conductor of electricity. Zinc-carbon dry cells use graphite as electrodes.
27 F
Dry ice consists of separate carbon dioxide (CO2) molecules. In each molecule, strong covalent bonds
hold the carbon and oxygen atoms together. The carbon dioxide molecules are packed close to one
another in a regular pattern. Weak van der Waals’ forces hold the molecules together.
Structure of dry ice:
LFZ
DBSCPOEJPYJEFNPMFDVMF
28 F
In iodine, strong covalent bonds hold the atoms in each molecule together.
Much weaker van der Waals’ forces hold the separate molecules together.
Structure of iodine:
LFZ
JPEJOFNPMFDVMF
29 F
Iodine is slightly soluble in water. The weak attractive forces between iodine and water molecules are
not strong enough to overcome the attractive forces between the water molecules.
30 F
Carbon dioxide has a simple molecular structure while silicon dioxide has a giant covalent structure.
Hence carbon dioxide and silicon dioxide have different physical properties.
31 T The melting point of hydrogen chloride is lower than that of potassium chloride.
Hydrogen chloride has a simple molecular structure. Little heat is needed to separate the molecules.
Potassium chloride has a giant ionic structure. A lot of heat is needed to overcome the strong ionic bonds
between the ions during melting.
89
32 T Silver conducts electricity due to the movement of mobile electrons. When silver is connected to a
battery, mobile electrons in the metal flow towards the positive terminal of the battery. At the same
time, electrons flow into the other end of the metal from the negative terminal of the battery.
CBUUFSZ
DPOEVDUJOHXJSF
TJMWFSXJSF
FMFDUSPOTGMPXJOUPUIF
TJMWFSXJSFGSPNUIFOFHBUJWF
UFSNJOBMPGUIFCBUUFSZ
FMFDUSPOTGMPXUPXBSET
UIFQPTJUJWFUFSNJOBMPG
UIFCBUUFSZ
Multiple choice questions
33 B
34 A Structure of sodium chloride:
LFZ
DIMPSJEFJPO
TPEJVNJPO
35 C
Option
Substance
Electrical conductivity
A
argon
B
potassium
C
potassium fluoride
an electrolyte; conducts electricity in molten state or aqueous solution,
and decomposed by electricity during conduction
D
tetrachloromethane
a non-conductor
a non-conductor
a conductor; not chemically changed during conduction
36 D Zinc chloride is an ionic compound. It does not conduct electricity in solid state. In solid state, ions in
the compound are held together by strong ionic bonds. They are not free to move.
90
37 C To melt an ionic compound, a lot of heat is needed to overcome the strong attractive forces (ionic
bonds) between the ions. Therefore ionic compounds have high melting points.
38 C Magnesium oxide is an ionic compound. The ions in it are held together by ionic bonds.
39 B
Option
Chloride
Structure of chloride
A
HCl
simple molecular structure
B
KCl
giant ionic structure
C
SCl2
simple molecular structure
D
PCl3
simple molecular structure
To melt KCl, a lot of heat is needed to overcome the strong ionic bonds between the ions.
The attractive forces between the molecules in HCl, SCl2 and PCl3 are weak. Little heat is needed to
separate the molecules.
Hence KCl has the highest melting point.
40 A Diamond has a giant covalent structure consisting of a network of covalent bonds. Relative motion of
the atoms is restricted. Diamond is the hardest substance known.
Structure of diamond:
LFZ
DBSCPOBUPN
41 C In silicon dioxide, each silicon atom is joined to four oxygen atoms by covalent bonds, while each
oxygen atom is joined to two silicon atoms by covalent bonds. Hence silicon dioxide has a giant
covalent structure.
Structure of silicon dioxide:
LFZ
PYZHFOBUPN
TJMJDPOBUPN
42 B Quartz has a giant covalent structure. To melt quartz, a lot of heat is needed to overcome the strong
covalent bonds between the atoms. Therefore quartz has a high melting point.
91
43 C Option A — Calcium (Ca) tarnishes in moist air. This is because calcium reacts with oxygen in the air
to form an oxide layer on the surface.
Option B — Sodium (Na) is stored in paraffin oil as it can react with oxygen in the air.
Option C — Silicon dioxide (SiO2) has a giant covalent structure. It is the most stable in moist air.
Option D — Sulphur dioxide (SO2) can react with moisture in air to form an acid.
44 D Graphite has a layered structure. Weak van der Waals’ forces exist between the layers.
The layers can easily slide over each other. Hence graphite has a slippery feel and can be used as a
lubricant.
Structure of graphite:
WBOEFS8BBMTGPSDFT
LFZ
DBSCPOBUPN
45 D Structure of dry ice:
LFZ
DBSCPOEJPYJEFNPMFDVMF
46 C
92
Option
Substance
Structure
A
diamond
giant covalent structure
B
mercury
giant metallic structure
C
nitrogen
simple molecular structure
D
quartz
giant covalent structure
47 C
Option
Substance
Structure
A
calcium oxide
giant ionic structure
B
graphite
giant covalent structure
C
iodine
simple molecular structure
D
sodium
giant metallic structure
∴ iodine consists of separate molecules.
48 D
Option
Oxide
Structure
A
MgO
giant ionic structure
B
Al2O3
giant ionic structure
C
SiO2
giant covalent structure
D
Cl2O
simple molecular structure
Electron diagram of Cl2O:
$M
0
$M
(Only electrons in the outermost shells are shown.)
49 C In solid carbon dioxide, van der Waals’ forces hold the molecules together.
50 C The attractive forces between bromine molecules are weak. Little heat is needed to separate the
molecules. Hence bromine has a low melting point. It exists as a liquid at room temperature and
pressure.
51 A Option A — For a substance with a simple molecular structure, the attractive forces between the
molecules are weak. Little heat is needed to separate the molecules. Hence the substance
has low melting and boiling points.
Option B — A substance with a simple molecular structure is generally quite soluble in non-aqueous
solvents.
Option C — The substance melts at –10 °C and boils at 58 °C. Hence it is a liquid at room
temperature and pressure.
Option D — Substances with simple molecular structures do not conduct electricity.
52 C Options A, B and D — Common salt, sugar and sodium nitrate are soluble in water.
Option C — Sulphur has a simple molecular structure. It is insoluble in water but soluble in nonaqueous solvents.
93
53 C
Option
Atomic number of
element
Name of element
Compound formed between the element and
chlorine
A
10
neon
no compound formed
B
11
sodium
an ionic compound formed
C
16
sulphur
a covalent compound formed
D
20
calcium
an ionic compound formed
Element X with an atomic number 16 (sulphur) reacts with chlorine to form a covalent compound with
a simple molecular structure.
Electron diagram of the compound formed between sulphur and chlorine:
$M
4
$M
(Only electrons in the outermost shells are shown.)
54 B Element X is phosphorus.
Option A — X (phosphorus) is a solid at room temperature and pressure.
Option B — To obtain stable electronic arrangements, one atom of X (phosphorus) bonds with three
hydrogen atoms.
Electron diagram of the compound formed:
)
1
)
)
(Only electrons in the outermost shells are shown.)
Option C — There are 15 electrons in an atom of X.
Option D — White phosphorus is composed of P4 molecules. The molecules are held together by van
der Waals’ forces.
55 A
94
Element
Atomic number
Name of element
X
8
oxygen
Y
12
magnesium
Z
17
chlorine
Option A — X (oxygen) and Z (chlorine) are non-metals. They combine to form a covalent compound.
Electron diagram of the compound formed:
$M
0
$M
(Only electrons in the outermost shells are shown.)
Option B — Y (magnesium) has a giant metallic structure.
Option C — Z (chlorine) exists as a gas at room conditions.
Option D — X (oxygen) supports combustion, but it is NOT flammable.
56 B
Element
Atomic number
Name of element
X
7
nitrogen
Y
14
silicon
Z
20
calcium
Option A — X (nitrogen) is a gas at room temperature and pressure. Its melting point is quite low.
Option B — Y (silicon) has a giant covalent structure similar to that of diamond.
Option C — Z (calcium) gives a brick-red flame in flame test.
Option D — X (nitrogen) and Z (calcium) combine to form an ionic compound with a giant ionic
structure.
Electron diagram of the compound formed:
$B
m
/
$B
m
/
$B
(Only electrons in the outermost shells are shown.)
57 A
Substance
Structure
Melting point (°C)
CO2
simple molecular structure
–56 (5.2 atm); sublimes (1 atm)
SiO2
giant covalent structure
1 610
Na2O
giant ionic structure
1 275
95
58 A Option A — W is probably a metal and hence it should NOT be a brittle solid. Metals are malleable
and ductile.
Option B — X has high melting and boiling points, and it is a non-conductor of electricity. It probably
has a giant covalent structure.
Option C — Y is a liquid at room temperature and pressure, and it is a good conductor of electricity.
It is probably mercury. Hence it is probably a good conductor of heat as well.
Option D — Z has low melting and boiling points, and it is a non-conductor of electricity. It is probably
a non-metal with a simple molecular structure.
59 B Y is a metal that melts at –39 °C and boils at 357 °C. It is a liquid at room temperature and pressure.
Therefore Y is likely to be mercury.
60 B Both W and X have low melting points, do not conduct electricity and are slightly soluble / insoluble in
water. Hence they probably have simple molecular structures.
W melts at –7 °C while X melts at 46 °C. Hence only X exists as a simple molecular solid at room
temperature.
61 B X (a metal) and Y (a non-metal) combine to form an ionic compound.
Option A — The ionic compound formed between X and Y has a giant ionic structure.
Option B — To melt or boil an ionic compound, a lot of heat is needed to overcome the strong
attractive forces (ionic bonds) between the ions. Therefore ionic compounds have high
melting and boiling points. The compound formed between X and Y should be a solid at
room temperature and pressure.
Option C — Electron diagram of the compound formed between X and Y:
m
9
:
(Only electrons in the outermost shells are shown.)
The chemical formula of the compound is XY.
Option D — The molten compound contains mobile ions and hence it can conduct electricity.
62 C
Element
Electronic arrangement of atom
Name of element
X
2,8,5
phosphorus
Y
2,7
fluorine
X and Y are non-metals. They combine to form a covalent compound Z.
96
Electron diagram of the compound Z:
1
'
'
'
(Only electrons in the outermost shells are shown.)
Option A — The bonds in Z are formed by electron sharing.
Option B — Z has a simple molecular structure.
Option C — Z has a low melting point.
Option D — The chemical formula of Z is XY3.
63 C
Chloride of
Melting point (°C)
Structure of the
chloride
Bonding in the
chloride
Metal / non-metal
X
–82
simple molecular
structure
covalent bonding
X is a non-metal
Y
808
giant ionic structure
ionic bonding
Y is a metal
Option B — The chloride of X is a covalent compound with a simple molecular structure.
Option C — The chloride of Y melts at 808 °C. Hence it is a solid at room temperature and pressure.
Option D — The chloride of Y conducts electricity in molten state.
64 C Magnesium is a good conductor of electricity due to the presence of mobile electrons.
LFZ
NBHOFTJVNJPO
NPCJMF
FMFDUSPOT
FMFDUSPO
65 D The ions in copper are packed closely and the metallic bonds holding them together are very strong.
To melt a piece of copper, a lot of heat is needed to overcome the strong attractive forces.
97
66 A Silver conducts electricity due to the movement of mobile electrons. When silver is connected to a
battery, mobile electrons in the metal flow towards the positive terminal of the battery. At the same
time, electrons flow into the other end of the metal from the negative terminal of the battery.
CBUUFSZ
DPOEVDUJOHXJSF
TJMWFSXJSF
FMFDUSPOTGMPXJOUPUIF
TJMWFSXJSFGSPNUIFOFHBUJWF
UFSNJOBMPGUIFCBUUFSZ
FMFDUSPOTGMPXUPXBSET
UIFQPTJUJWFUFSNJOBMPG
UIFCBUUFSZ
67 D A metal conducts electricity in both solid and molten states. It is also insoluble in water.
68 B
Option
Substance
Melting point (°C)
Boiling point (°C)
State at –50 °C and
1 atm pressure
A
bromine
–7
59
solid
B
chlorine
–101
–35
liquid
C
oxygen
–218
–183
gas
D
hydrogen bromide
–88
–67
gas
Only chlorine exists as a liquid at –50 °C and 1 atm pressure.
69 C All the four substances have simple molecular structures. In these substances, strong covalent bonds
hold the atoms in each molecule together. Much weaker van der Waals’ forces hold the separate
molecules together.
e.g. van der Waals’ forces exist between oxygen molecules:
XFBLWBOEFS8BBMTGPSDFT
CFUXFFOPYZHFONPMFDVMFT
PYZHFONPMFDVMF
98
70 D
Option
Substance
Structure
(1)
oxygen
simple molecular structure
(2)
potassium oxide
giant ionic structure
(3)
silicon dioxide
giant covalent structure
Option
Substance
Can conduct electricity?
Particles responsible for the
conduction
(1)
graphite
yes
mobile delocalized electrons
(2)
molten zinc chloride
yes
mobile ions
(3)
magnesium sulphate solution
yes
mobile ions
71 D
72 C (1) The chemical formula of silicon carbide is SiC.
(2) Silicon carbide has a giant covalent structure.
To melt it, a lot of heat is needed to overcome the strong covalent bonds between the atoms.
Hence it has a high melting point.
(3) Silicon carbide has a giant covalent structure.
It is insoluble in water because the atoms are held together by strong covalent bonds and it is very
difficult to separate the atoms.
73 A (2) The structure of germanium is similar to that of silicon. It has a giant covalent structure.
To melt it, a lot of heat is needed to overcome the strong covalent bonds between the atoms.
Hence germanium has a high melting point (937 °C).
(3) Electron diagram of the fluoride of germanium:
'
'
(F
'
'
(Only electrons in the outermost shells are shown.)
The fluoride has a simple molecular structure.
99
74 A This form of solid carbon is composed of C60 molecules. It probably has a simple molecular structure.
(1) Substances with simple molecular structures are usually insoluble in water.
(2) Diamond is the hardest substance known. The attractive forces between C60 molecules are weak.
Hence this form of solid carbon is softer than diamond.
(3) The melting point of graphite is higher than that of this form of solid carbon.
To melt graphite, a lot of heat is needed to overcome the strong covalent bonds between the
atoms.
The attractive forces between the C60 molecules are weak. Little heat is needed to separate the
molecules.
75 B (1) Selenium and hydrogen are non-metals. They combine to form a covalent compound.
(2) Electron diagram of chloride of selenium:
$M
4F
$M
(Only electrons in the outermost shells are shown.)
The chloride of selenium has a simple molecular structure.
(3) Selenium (a non-metal) and a Group I element (a metal) combine to form an ionic compound.
A selenium atom has 6 outermost shell electrons. In the above reaction, a selenium atom gains
2 electrons to obtain a stable electronic arrangement. An ion carrying two negative charges (Se2–)
forms.
76 B Structure of graphite:
WBOEFS8BBMTGPSDFT
LFZ
DBSCPOBUPN
(1) Graphite is slippery because weak van der Waals’ forces exist between the carbon layers in it.
Hence the layers can easily slide over each other.
(2) Graphite has a layered structure. Within each layer, each carbon atom is covalently bonded to three
other atoms.
To melt graphite, the covalent bonds between the atoms must be overcome. Graphite has a high
melting point; that means a lot of heat is needed to overcome the covalent bonds. The fact that
‘graphite has a high melting point’ is an evidence to support that covalent bonds are strong.
(3) Graphite can conduct electricity due to the presence of delocalized electrons.
100
77 A (1) X exists as diatomic molecules. It has a simple molecular structure.
(2) Weak van der Waals’ forces hold the molecules together. Hence X has a very low melting point.
(3) Within each molecule, a strong covalent bond holds the two atoms together.
78 A (3) Sodium chloride melts at 808 °C. It is still a solid at 750 °C and thus cannot conduct electricity at
750 °C.
79 A (2) SiH4 has a simple molecular structure as deduced from its low melting point.
(3) H2S is a gas at room temperature and pressure.
80 A
Structure of substance
Giant covalent
Giant ionic
Giant metallic
Melting point
high
high
high
Electrical conductivity at room temperature
non-conductor
(except graphite)
non-conductor
conductor
Property of substance
∴ X may have a giant covalent or giant ionic stucture.
81 B
Element
Atomic number
Name of element
Electronic arrangement of atom
X
19
potassium
2,8,8,1
Y
16
sulphur
2,8,6
(1) When X and Y combine, two atoms of X lose 2 electrons and the electrons are accepted by one
atom of Y. An ionic compound with a chemical formula X2Y forms.
Electron diagram of the compound:
m
9
:
9
(Only electrons in the outermost shells are shown.)
(2) The ionic compound formed between X and Y is insoluble in non-aqueous solvents.
(3) The aqueous solution of the compound formed between X and Y contains mobile ions. Hence it
can conduct electricity.
82 B
Element
Atomic number
Name of element
Electronic arrangement of atom
X
8
oxygen
2,6
Y
9
fluorine
2,7
101
To obtain stable electronic arrangements, one atom of X forms a single bond with each of two atoms of Y.
Electron diagram of the compound formed:
:
9
:
(Only electrons in the outermost shells are shown.)
(1) A covalent compound with a simple molecular structure is formed.
(2) The chemical formula of the compound is XY2, i.e. OF2.
(3) The compound is slightly soluble in water.
83 D (1) An oxide ion carries 2 negative charges. The simplest ratio of ion of X to O2– in X2O3 is 2 : 3 and
the net charge of the compound is zero. Hence the ion of X carries 3 positive charges.
Use the following steps to work out the chemical formula of the chloride of X:
Step
Chloride of X
1 Write down the symbols of ions in the compound.
X
Cl
2 Write down the number of charges of each ion on the
top of each symbol.
3
X
1
Cl
3
X
= X1
3 Cross multiply the numbers and the symbols.
4 Combine the symbols.
1
Cl
= Cl3
XCl3
(2) X (a metal) and oxygen (a non-metal) combine to form an ionic oxide. Hence X2O3 has a giant
ionic structure.
(3) X2O3 (an ionic oxide) conducts electricity in molten state due to the presence of mobile ions.
84 A
Element
Atomic number
Name of element
Electronic arrangement
of atom
Relative atomic mass
P
9
fluorine
2,7
19.0
Q
17
chlorine
2,8,7
35.5
(1) P and Q are non-metals. They combine to form a covalent compound by electron sharing.
Electron diagram of compound X:
2
1
(Only electrons in the outermost shells are shown.)
102
(2) Relative molecular mass of X = relative atomic mass of P + relative atomic mass of Q
= 19.0 + 35.5
= 54.5
(3) X has a simple molecular structure. It is a gas at room temperature and pressure.
85 A
Element
Atomic number
Name of element
Electronic arrangement
of atom
Relative atomic mass
P
7
nitrogen
2,5
14.0
Q
12
magnesium
2,8,2
24.3
(1) P (a non-metal) and Q (a metal) combine to form an ionic compound by electron transfer.
Electron diagram of the compound X:
2
m
1
2
m
1
2
(Only electrons in the outermost shells are shown.)
(2) X has a giant ionic structure. It should have a high melting point and hence it should be a solid at
room temperature and pressure.
(3) Formula mass of X = 2 x relative atomic mass of P + 3 x relative atomic mass of Q
= 2 x 14.0 + 3 x 24.3
= 100.9
86 D Not all ionic compounds are soluble in water. For example, calcium carbonate is insoluble in water.
Attractive forces exist between ions in an ionic compound and water molecules. However, the attractive
forces between water molecules are weaker than ionic bonds.
87 A To melt potassium chloride, a lot of heat is needed to overcome the strong ionic bonds between the
ions. Hence potassium chloride has a high melting point.
103
88 D Iodine is very soluble in non-aqueous solvents. The attractive forces between molecules of non-aqueous
solvents are similar to those between iodine molecules. Hence molecules of iodine and non-aqueous
solvents can mix together easily.
JPEJOFNPMFDVMFT
BUUSBDUJWFGPSDFTCFUXFFO
JPEJOFNPMFDVMFTBOE
NPMFDVMFTPGOPOBRVFPVT
TPMWFOUTJNJMBSUPUIPTF
CFUXFFONPMFDVMFTPG
OPOBRVFPVTTPMWFOU
NPMFDVMFTPG
OPOBRVFPVT
TPMWFOU
89 C Many covalent compounds consist of separate molecules held together by weak van der Waals’ forces.
The molecules separate easily when heated. Therefore these covalent compounds exist as gases, liquids
or solids with low melting points.
90 C Graphite can conduct electricity but diamond cannot.
Diamond and graphite are different forms of carbon.
91 C Calcium carbonate is an ionic compound but it is insoluble in water.
Ammonia is a covalent compound but it dissolves well in water.
92 D Sugar is a non-conductor of electricity. It does not form ions when dissolved in water.
93 B When iodine sublimes, it absorbs heat so as to overcome the van der Waals’ forces between iodine
molecules.
94 C Carbon dioxide has a simple molecular structure while silicon dioxide has a giant covalent structure.
Therefore carbon dioxide and silicon dioxide have very different physical properties.
95 A Metals are good conductors of electricity due to the movement of mobile electrons in a metal. When
heating one end of a piece of metal, the mobile electrons get more energy. They move faster and
collide with neighbouring electrons. This helps to transfer the heat.
Fm
m
F
Fm
Fm
Fm
104
Fm
Part B
Topic-based exercise
Multiple choice questions
1
B Option A — Atoms are NOT indivisible. They contain subatomic particles.
Option C — Isotopes of an element are NOT identical.
Option D — Atoms of different elements may have the same mass.
2
B
Substance
X
Y
Z
Melting point (°C)
–146
–210
–108
Boiling point (°C)
–80
–105
–45
State at –70 °C
gas
gas
liquid
∴ only Z exists in the liquid state at –70 °C.
3
B Atomic number of
39
19
K = 19 = number of protons in an atom
= number of electrons in an atom
Number of neutrons = mass number – atomic number
= 39 – 19
= 20
4
D
Option
∴
5
35
17
Species
A
9
B
23
Number of electrons
Number of neutrons
4
5
10
12
4Be
11Na
+
C
27
13Al
13
14
D
35
18
18
–
17Cl
Cl– contains equal numbers of electrons and neutrons.
B An atom of M loses 2 electrons to form a M2+ ion with 27 electrons.
∴ number of electrons in an atom of M =
=
=
=
6
29
number of protons in an atom of M
atomic number of M
atomic number of Cu
84 x 45.0 + 86 x 55.0
100
= 85.1
C Relative atomic mass of Kr =
105
7
D Let the relative abundance of
126
I and
127
I be (100 – y)% and y% respectively.
126 x (100 – y) + 127 x y
100
y = 90.0
126.9 =
8
D
Halogen
Atomic number
F
9
Cl
17
Difference in atomic number
Electronic arrangement of atom
2,7
8
2,8,7
18
Br
35
2,8,18,7
The difference in atomic number between the first three successive halogens is either 8 or 18. This
is because the second electron shell can hold 8 electrons while the third electron shell can hold 18
electrons.
Hence the atomic number of a halogen A is x, then the atomic number of another halogen B could
be x + 18.
9
A
Option
Atom
4
2
20
10Y
2,8
12
2,4
24
12Y
2,8,2
23
11X
2,8,1
35
17Y
2,8,7
40
2,8,8
40
2,8,8,2
2X
A
6X
B
C
18X
D
20Y
4
2
X and
20
10
Electronic arrangement of atom
Y are atoms of noble gases. They have similar chemical properties.
10 A The reactivity of Group I elements increases down the group. Hence the reactivity of Rb is greater than
that of Na.
The reactivity of Group VII elements decreases down the group. Hence the reactivity of Cl is greater
than that of Br.
Therefore Rb of Group I and Cl of Group VII would react with each other most vigorously.
106
11 D A helium atom has two outermost shell electrons. Atoms of other noble gases have eight outermost
shell electrons.
12 B
Atom
W
X
Y
Z
Atomic number
7
17
8
18
Number of neutrons
7
18
8
20
Name of atom
nitrogen
chlorine
oxygen
argon
Period to which it belongs
Period 2
Period 3
Period 2
Period 3
Option A — W and Y belong to the same period, Period 2.
Option B — X and Z are atoms of different elements belonging to different groups. Hence they have
different chemical properties.
Option C — W (nitrogen) is a gas at room temperature and pressure.
Option D — X (chlorine) is a non-metal.
13 D Number of neutrons = mass number – atomic number
= 79 – 34
= 45
14 A Atomic number of Ce = 58 = number of electrons in an atom
Number of electrons in the Ce3+ ion = 58 – 3 = 55
15 D An atom of a Group I element X forms an ion X+ by losing 1 electron.
e.g.
Q
O
Q
O
TPEJVNBUPN
TPEJVNJPO
m
F
Options A, B and C — Losing an electron does not affect the mass number, nuclear charge and
atomic number of the atom of element X.
Option D — An atom of X has one more occupied electron shell than the X+ ion.
16 A Option A — An atom of X gains 1 electron to form the ion X–. Hence the electronic arrangement of
an atom of X is 2,8,7. X is chlorine, a halogen.
Option C — X is a Group VII element.
Option D — An atom of X has 3 occupied electron shells. Hence X is a Period 3 element.
107
17 A
Option
Compound
Electron diagram
/B
A
m
Na2O
0
/B
-J
m
B
Li3N
/
-J
-J
m
C
CaO
0
$B
m
D
SrCl2
$M
4S
m
$M
∴ the Na+ ion and O2– ion in Na2O have the same electronic arrangement.
18 B Solid copper(II) chloride does not conduct electricity. The ions in solid copper(II) chloride are held
together by strong ionic bonds. They are not free to move.
When water is added to dissolve the solid copper(II) chloride, the ions become mobile. Hence the
solution can conduct electricity.
108
19 C Chlorine exists as diatomic molecules. It is a gas at room temperature and pressure. The molecules are
arranged in a random way.
3–
20 B The atom of X gains 3 electrons in order to obtain a stable electronic arrangement. A stable X
formed.
ion is
For atoms of non-metals in Group V, VI and VII, they gain ‘8 – group number’ electrons in order to
obtain stable electronic arrangements. Hence X probably belongs to Group V of the periodic table.
X and chlorine are non-metals. They combine to form a covalent compound. To obtain stable electronic
arrangements, one atom of X forms a single bond with each of three chlorine atoms.
Electron diagram of the compound:
$M
9
$M
$M
(Only electrons in the outermost shells are shown.)
The chemical formula of the compound is XCl3.
21 C E forms a sulphate with the chemical formula E2(SO4)3. The sulphate ion (SO42–) carries 2 negative
charges and the net charge of the compound must be zero. The simplest ratio of ion of E to sulphate
ion in the sulphate is 2 : 3. It can be deduced that the ion of E carries 3 positive charges.
Z forms a molecular hydride. Hence Z is a non-metal. The chemical formula of the hydride is H2Z. That
means to obtain stable electronic arrangements, one atom of Z forms single covalent bonds with two
hydrogen atoms. An atom of Z contributes 2 electrons for sharing as it needs 2 electrons to obtain a
stable electronic arrangement. Therefore it probably has 6 outermost shell electrons.
An atom of Z forms a stable anion Z2– by gaining 2 electrons. Use the following steps to work out the
chemical formula of the compound formed between E and Z:
Step
Compound formed between E and Z
1 Write down the symbols of ions in the compound.
E
Z
2 Write down the number of charges of each ion on the top
of each symbol.
3
E
2
Z
3
E
= E2
2
Z
= Z3
3 Cross multiply the numbers and the symbols.
4 Combine the symbols.
E2Z3
109
22 A In the compound, the atom of Y contributes 1 electron for sharing with an atom of X and 3 electrons
for sharing with an atom of Z. The atom of X contributes 4 electrons for sharing as it needs 4
electrons to obtain a stable electronic arrangement. It can be deduced that an atom of Y has 4
outermost shell electrons.
The electron diagram of hydrogen cyanide is shown in the question.
)
$
/
(Only electrons in the outermost shells are shown.)
23 C An atom of X gains 3 electrons to form the anion X3– with an electronic arrangement 2,8. Hence the
electronic arrangement of an atom of X is 2,5. X is nitrogen.
Option A — X (nitrogen) is in Group V of the periodic table.
Option B — X (nitrogen) is a gas at room temperature and pressure.
Option C — X (nitrogen) exists as diatomic molecules.
Option D — Neon is a noble gas. A neon atom has 8 outermost shell electrons. A special stability is
obtained when this happens. Hence neon will NOT form a compound with nitrogen.
24 A A colour moved towards the right (i.e the negative electrode) because positive copper(II) ions, which
are blue in colour, move towards the negative electrode.
25 C The aqueous solution of compound YX is colourless. Hence Y2+(aq) and X2–(aq) ions are colourless.
The green colour of the aqueous solution of compound WX is due to the W2+(aq) ions. Hence W2+(aq)
ion is green in colour.
The purple colour of the aqueous solution of compound YZ is due to the Z2–(aq) ions. Hence Z2–(aq)
ion is purple in colour.
26 C Zinc chloride is an ionic compound. It probably has a high melting point, does not conduct electricity
in solid state and is soluble in water.
27 B
Option
Solid
Melting point
Electrical conductivity in solid state
A
Iodine
low
non-conducting
B
Potassium
low
conducting
C
Potassium fluoride
high
non-conducting
D
Silicon dioxide
high
non-conducting
28 C Option A — Solid carbon dioxide has a simple molecular structure.
Option B — Carbon dioxide is denser than air. It is NOT used to fill weather balloons.
Option C — Carbon dioxide dissolves in water to give carbonic acid.
Option D — Carbon dioxide has a simple molecular structure while silicon dioxide has a giant covalent
structure. Hence they have different physical properties.
110
29 C Silicon has a giant covalent structure. To melt silicon, a lot of heat is needed to overcome the strong
covalent bonds between the atoms. Hence silicon has a high melting point.
Structure of silicon:
LFZ
TJMJDPOBUPN
30 C Dry ice has a simple molecular structure. Weak van der Waals’ forces hold the molecules together.
31 D
Option
Substance
A
magnesium fluoride
Bonding type
barium fluoride
ionic bonding
potassium fluoride
carbon monoxide
B
sulphur dioxide
covalent bonding
methane
chromium
C
magnesium
metallic bonding
nickel
boron trichloride
D
silicon tetrachloride
sodium chloride
32 A
covalent bonding
covalent bonding
ionic bonding
Option
Chloride
State at room temperature and pressure
A
HCl
gas
B
KCl
solid
C
AlCl3
solid
D
CCl4
liquid
∴ HCl has the lowest boiling point.
111
33 D The element (melting point above 3 000 °C) probably has a giant structure.
It forms a gaseous oxide. It can be deduced that the oxide is a covalent compound with a simple
molecular structure. Hence element X is probably a non-metal.
Thus X should have a covalent network structure.
34 D Options A, B and C — Barium chloride, caesium chloride and calcium chloride are ionic compounds.
They are likely to be insoluble in tetrachloromethane, a non-aqueous solvent.
Option D — Phosphorus trichloride is a covalent compound with a simple molecular structure. It is
likely to be soluble in tetrachloromethane, a non-aqueous solvent.
35 C Substances X and Y are probably ionic compounds as they are non-conductors of electricity when in
solid state but conducts electricity when in molten state.
Substance Z has very high melting and boiling points. It conducts electricity when in solid state. It is
likely to be graphite.
36 C Options A and B — W melts at 71 °C while X melts at 98 °C. They are solids at room temperature
and pressure.
Option C — Y melts at –130 °C and boils at 36 °C. It is a liquid at room temperature and pressure. It
is a poor conductor of electricity. It probably has a simple molecular structure.
Option D — Z melts at –138 °C and boils at –0.5 °C. It is a gas at room temperature and pressure.
37 D
Element
Atomic number
Name of element
Electronic arrangement of atom
X
8
oxygen
2,6
Y
9
fluorine
2,7
X and Y are non-metals. They combine to form a covalent compound Z.
Electron diagram of Z:
:
9
:
(Only electrons in the outermost shells are shown.)
Option A — The chemical formula of Z is XY2, i.e. OF2.
Option B — Z is formed by electron sharing.
Option C — Van der Waals’ forces exist between molecules of Z.
Option D — Z is a gas at room temperature and pressure.
112
38 D
Element
Atomic number
Name of element
Electronic arrangement of atom
X
14
silicon
2,8,4
Y
17
chlorine
2,8,7
Option A — X (silicon) has a giant covalent structure.
Structure of silicon:
LFZ
TJMJDPOBUPN
Option B — Y (chlorine) is a gas at room temperature and pressure.
Option C — X and Y are non-metals. They combine to form a covalent compound with a simple
molecular structure.
Electron diagram of the compound:
$M
$M
4J
$M
$M
(Only electrons in the outermost shells are shown.)
Option D — Weak van der Waals’ forces exist between molecules of the compound formed between X
and Y. Hence the compound has a low melting point.
39 B Options A and B — From its low boiling point, it can be deduced that the oxide of X is a covalent
compound with a simple molecular structure. Therefore X is probably a non-metal.
Option C — The oxide of Y boils at 2 230 °C. It is a solid at room temperature and pressure.
Option D — From its high boiling point, it can be deduced that the oxide of Y has a giant structure.
Van der Waals’ forces do not exist in it.
113
40 C (1) Number of neutrons in
68
31
Ga = mass number – atomic mass
= 68 – 31
= 37
(2) The atomic number of the gallium isotope
68
31
Ga is 31.
(3) Isotopes of the same element have the same electronic arrangement and hence the same chemical
properties.
41 D
Species
Electronic arrangement of species
+
2
3–
2,8
Li
N
Ar
2,8,8
∴ N3– and Ar have an octet structure in their outermost shells.
42 D (1) The reactivity of Group II elements increases down the group, i.e. increases with relative atomic
mass.
3FMBUJWFBUPNJDNBTT
CFSZMMJVN
NBHOFTJVN
DBMDJVN
SFBDUJWJUZ
JODSFBTJOH
EPXOUIFHSPVQ
(2) The melting point of Group VII elements increases down the group, i.e. increases with atomic
number.
3FMBUJWFBUPNJDNBTT
NFMUJOHQPJOU
GMVPSJOF
DIMPSJOF
CSPNJOF
JPEJOF
JODSFBTJOH
EPXOUIFHSPVQ
114
43 A (1) Across the third period of the periodic table, the atomic size of the elements shows a gradual
decrease.
1FSJPE
"UPNJDSBEJVTQN
/B
.H
"M
4J
4
1
$M "S
"UPNJDOVNCFS
(2) Across the third period of the periodic table, the elements change from metals through metalloids
to non-metals, i.e. the metallic character of the elements decreases.
Group
I
II
III
IV
V
VI
VII
Period 3 element
Sodium
Magnesium
Aluminium
Silicon
Phosphorus
Sulphur
Chlorine
metalloid
metals
non-metals
Type of element
metallic character decreasing
(3) Across the third period of the periodic table, the melting point of elements rises to Group IV and
then falls to low values.
.FMUJOHQPJOUž$
m
&MFNFOU
(SPVQOVNCFS
44 A (1)
Group
Period 3 element
Electronic
arrangement of atom
I
Sodium
2,8,1
/B
*
.H
**
II
"M
***
4J
*7
III
Magnesium Aluminium
2,8,2
2,8,3
1
7
4
7*
$M
7**
IV
V
VI
VII
Silicon
Phosphorus
Sulphur
Chlorine
2,8,4
2,8,5
2,8,6
2,8,7
∴ the number of outermost shell electrons in atoms of Period 3 elements increases from sodium
to chlorine.
115
(2) Across the third period of the periodic table, the metals on the left tend to lose electrons while the
non-metals on the right tend to gain electrons. The ability of elements to gain electrons increases.
(3) Across the third period of the periodic table, the metals on the left (e.g. sodium and magnesium)
form ionic chlorides while the non-metals on the right (e.g. phosphorus and sulphur) form covalent
chlorides. Chlorides of the elements change from ionic to covalent.
Group
I
Period 3 element
II
Sodium
III
Magnesium Aluminium
IV
V
VI
VII
Silicon
Phosphorus
Sulphur
Chlorine
form ionic chlorides
form covalent chlorides
45 D (3) Potassium reacts with water to give an alkaline solution.
potassium + water
46 B (1)
potassium hydroxide + hydrogen
Element
Relative atomic mass
Electronic arrangement of atom
Beryllium
9.0
2,2
Magnesium
24.3
2,8,2
Calcium
40.0
2,8,8,2
A beryllium atom has 2 occupied electron shells, a magnesium atom has 3 and a calcium atom has 4.
Hence the atomic size of the elements increases down the group, i.e. increases with atomic mass.
(2)
Element
Relative atomic mass
Melting point (°C)
Beryllium
9.0
1 280
Magnesium
24.3
650
Calcium
40.0
838
Group II elements do not show a gradual change in melting point with increasing relative atomic
mass.
(3) The reactivity of Group II elements increases down the group, i.e. increases with relative atomic
mass.
3FMBUJWFBUPNJDNBTT
CFSZMMJVN
NBHOFTJVN
DBMDJVN
SFBDUJWJUZ
JODSFBTJOH
EPXOUIFHSPVQ
116
47 B Caesium (a metal) and bromine (a non-metal) react to form an ionic compound.
Electron diagram of the compound:
m
$T
#S
(Only electrons in the outermost shells are shown.)
(2) The compound is soluble in water but insoluble in non-aqueous solvents.
48 B (1) Strontium is a reactive metal. It will burn in air.
(2) Strontium and calcium are Group II elements. Calcium reacts with water to liberate hydrogen.
As the reactivity of Group II elements increases down the group, strontium is more reactive than
calcium and should react with water to liberate hydrogen as well.
(3) Calcium and magnesium form many white compounds. Hence many strontium compounds should
also be white in colour.
49 D (1) The structure of germanium is similar to that of silicon. Both have a giant covalent structure.
(2) A germanium atom has 4 occupied electron shells while a silicon atom has 3. The atomic size of
germanium is larger than that of silicon.
Group IV element
Atomic number
Electronic arrangement of atom
Silicon
14
2,8,4
Germanium
32
2,8,18,4
(3) Electron diagram of fluoride of germanium:
'
'
(F
'
'
(Only electrons in the outermost shells are shown.)
50 D
117
51 D (3) Iodine changes from the solid state to the gaseous state directly when heated.
FWBQPSBUJOH
EJTI
JPEJOF
WBQPVS
JPEJOFTPMJE
IFBU
52 D (1)
Element
Colour
Chlorine
greenish yellow
Bromine
reddish brown
Iodine
black
(2) The reactivity of halogens decreases down the group.
DIMPSJOF
SFBDUJWJUZ
EFDSFBTJOHEPXO
CSPNJOF
UIFHSPVQ
JPEJOF
(3) The halogens exist as diatomic molecules. Van der Waals’ forces exist between the molecules.
53 D Electron diagram of hydrogen peroxide:
)
0
0
)
(1) The number of bonding electrons contributed by each hydrogen atom in the molecule is 1.
(2) The number of bonding electrons contributed by each oxygen atom in the molecule is 2.
(3) The total number of electrons in the molecule
= 2 x number of electrons in one oxygen atom + 2 x number of electron in one hydrogen atom
= 2 x 8 + 2 x 1
= 18
118
54 A Electron diagram of calcium carbonate:
m
0
$B
$
0
0
(Only electrons in the outermost shells are shown.)
Both ionic bond and covalent bond exist in calcium carbonate.
55 B (1) Copper has a giant metallic structure.
(2) Nitrogen dioxide has a simple molecular structure.
(3) Silicon dioxide has a giant covalent structure.
56 D (1) Neon exists as monoatomic molecules. Van der Waals’ forces exist between its molecules.
57 A
Option
Substance
Bonding type
(1)
copper
mercury
tungsten
metallic bonding
(2)
dry ice
nitrogen dioxide
water
covalent bonding
(3)
copper(II) chloride
hydrogen chloride
zinc chloride
ionic bonding
covalent bonding
ionic bonding
58 C (1) Graphite has a layered structure. Within each layer, each carbon atom uses three outermost shell
electrons in forming covalent bonds with three other atoms. The remaining electron is delocalized
between the layers of carbon atoms.
(2) Electron diagram of a methane molecule:
)
)
$
)
)
(Only electrons in the outermost shells are shown.)
The outermost shell electrons of atoms in a methane molecule are shared to form covalent bonds.
There are NO delocalized electrons.
119
(3) Sodium consists of tightly packed positive ions surrounded by a sea of delocalized electrons.
LFZ
NFUBMJPO
FMFDUSPO
NPCJMF
FMFDUSPOT
59 B To melt potassium chloride, the ionic bonds between the ions in potassium chloride must be overcome.
Potassium chloride has a high melting point; that means a lot of heat is needed to overcome the ionic
bonds. The fact that ‘potassium chloride has a high melting point’ is an evidence to support that ionic
bonds are strong.
60 C
Element
Atomic number
Name of element Electronic arrangement
Relative atomic mass
P
7
nitrogen
2,5
14.0
Q
9
fluorine
2,7
19.0
(1) P and Q are non-metals. They combine to form a covalent compound by electron sharing.
Electron diagram of compound X:
2
1
2
2
(Only electrons in the outermost shells are shown.)
(2) Relative molecular mass of X = relative atomic mass of P + 3 x relative atomic mass of Q
= 14.0 + 3 x 19.0
= 71.0
(3) Van der Waals’ forces exist between the molecules of X.
61 A
62 B Argon is used to fill electric light bulbs because it does not react with the metal filament in light bulbs.
63 B Nitrogen is used to fill the packets of potato chips because it is unreactive and can provide an inert
atmosphere.
64 C When carbon dioxide dissolves in water, carbonic acid is formed. The solution can conduct electricity.
120
65 A Metals are good conductors of electricity due to the movement of mobile electrons in a metal.
When a metal is connected to a battery, mobile electrons in the metal flow towards the positive
terminal of the battery. At the same time, electrons flow into the other end of the metal from the
negative terminal of the battery.
CBUUFSZ
DPOEVDUJOHXJSF
BNFUBM
FMFDUSPOTGMPXJOUPUIF
NFUBMGSPNUIFOFHBUJWF
UFSNJOBMPGUIFCBUUFSZ
FMFDUSPOTGMPXUPXBSET
UIFQPTJUJWFUFSNJOBMPG
UIFCBUUFSZ
66 C Bromine is a non-metal and does not conduct electricity.
Bromine is a liquid at room conditions and contains mobile molecules.
67 D Some covalent substances are soluble in water, e.g. carbon dioxide, sulphur dioxide.
Some covalent substances have giant covalent structures, e.g. silicon dioxide.
68 A The boiling point of Group VII elements increases down the group.
GMVPSJOF
DIMPSJOF
CPJMJOHQPJOU
CSPNJOF
EPXOUIFHSPVQ
JODSFBTJOH
JPEJOF
Hence the boiling point of bromine is higher than that of chlorine.
The boiling points of bromine and chlorine are related to the strength of the van der Waals’ forces
between their molecules.
121
69 B The melting point of hydrogen chloride is low because weak van der Waals’ forces exist between its
molecules. Little heat is needed to separate the molecules.
70 C An ammonia molecule contains 3 hydrogen atoms and 1 nitrogen atom.
However, the mass of hydrogen is NOT three times that of nitrogen as the relative atomic masses of
hydrogen and nitrogen are different.
Relative molecular mass of ammonia
= 3 x relative atomic mass of hydrogen + relative atomic mass of nitrogen
= 3 x 1.0 + 14.0
= 17.0
3 x 1.0
x 100%
17.0
= 17.6%
Percentage by mass of hydrogen in ammonia =
Percentage by mass of nitrogen in ammonia =
14.0
x 100%
17.0
= 82.4%
Short questions
71
122
Element
Symbol
Metal / Metalloid / Non-metal
Argon
Ar
non-metal
Carbon
C
non-metal
Calcium
Ca
metal
Fluorine
F
non-metal
Germanium
Ge
metalloid
Lithium
Li
metal
Magnesium
Mg
metal
Neon
Ne
non-metal
Nitrogen
N
non-metal
Potassium
K
metal
Phosphorus
P
non-metal
Silicon
Si
metalloid
(0.5 x 24)
72
Atomic
number
Atom
Mass
number
Symbol
Number of
protons neutrons electrons
Oxygen
8
16
16
8
Sodium
11
23
23
11
Aluminium
13
27
27
13
Sulphur
16
32
32
16
S
16
16
16
Chlorine
17
35
35
17
Cl
17
18
17
Potassium
19
39
39
19
K
19
20
19
Calcium
20
40
40
20
Ca
20
20
20
Iron
26
56
56
26
26
30
26
O
8
8
8
Na
11
12
11
13
14
13
Al
Fe
(0.5 x 40)
73 a) chlorine
(1)
b) copper
(1)
c) phosphorus
(1)
d) nitrogen
(1)
e) boron
(1)
f) nickel
(1)
74 a)
Number of
Species
Atomic
number
Mass
number
protons
neutrons
electrons
Electronic
arrangement
i)
Beryllium atom
4
9
4
5
4
2,2
ii)
Neon atom
10
20
10
10
10
2,8
iii) Silicon atom
14
28
14
14
14
2,8,4
iv) Phosphorus atom
15
31
15
16
15
2,8,5
v)
19
39
19
20
18
2,8,8
vi) Nitride ion
7
14
7
7
10
2,8
vii) Magnesium ion
12
24
12
12
10
2,8
viii) Fluoride ion
9
19
9
10
10
2,8
ix) Sodium ion
11
23
11
12
10
2,8
x)
11
23
11
12
11
2,8,1
Potassium ion
Sodium atom
(0.5 x 40)
b) Species (vi) & (viii) / nitride ion and fluoride ion
(0.5 x 2)
c) Group IV;
it has 4 electrons in its outermost shell.
d) Species (ix) is the cation of species (x) / (ix) and (x) are the ion and atom of the same element.
(1)
(1)
(1)
123
75 a)
b)
Name
Chemical formula
Aluminium hydroxide
Al(OH)3
Ammonium dichromate
(NH4)2Cr2O7
Calcium phosphate
Ca3(PO4)2
Copper(II) chloride
CuCl2
Iron(III) oxide
Fe2O3
Magnesium hydroxide
Mg(OH)2
Potassium carbonate
K2CO3
Sodium sulphite
Na2SO3
(1 x 8)
Chemical formula
Name
KHCO3
potassium hydrogencarbonate
Fe2(SO4)3
iron(III) sulphate
Cu(OH)2
copper(II) hydroxide
Mg3N2
magnesium nitride
Zn(NO3)2
zinc nitrate
NaS
sodium sulphide
Al2O3
aluminium oxide
AgCl
silver chloride
76
Cation
(1 x 8)
Anion
Compound
Name
Formula
Name
Formula
Name
Formula
Colour of
aqueous
solution
ammonium
NH4+
carbonate
CO32–
ammonium carbonate
(NH4)2CO3
colourless
copper(II) nitrate
Cu(NO3)2
blue
copper(II)
2+
Cu
nitrate
NO3
–
iron(II)
Fe2+
sulphate
SO42–
iron(II) sulphate
FeSO4
pale green
potassium
K+
permanganate
MnO4–
potassium permanganate
KMnO4
purple
nickel(II)
Ni2+
chloride
Cl–
nickel(II) chloride
NiCl2
green
iodide
–
aluminium iodide
All3
colourless
aluminium
chromium(III)
sodium
zinc
3+
Al
3+
Cr
Na+
2+
Zn
I
–
chloride
Cl
chromium(III) chloride
CrCl3
green
dichromate
Cr2O72–
sodium dichromate
Na2Cr2O7
orange
bromide
Br
zinc bromide
ZnBr2
colourless
–
(0.5 x 42)
124
77 a) i)
m
$M
$M
$B
$B
m
$M
$M
(1)
ii)
.H
m
/
/
.H
.H
.H
m
/
.H
/
.H
(1)
iii)
/B
m
/B
4
4
/B
/B
(1)
125
b) i)
1
$M
$M
$M
(1)
ii)
0
)
)
)
0
)
(1)
iii)
$M
'
$M
'
(1)
78
Substance
Chemical formula
Relative atomic
mass(es)
Formula mass /
relative molecular
mass
Oxygen
O2
O = 16.0
32.0
Carbon dioxide
CO2
C = 12.0
O = 16.0
44.0
Potassium nitrate
KNO3
N = 14.0
O = 16.0
K = 39.1
101.1
Calcium hydroxide
Ca(OH)2
H = 1.0
O = 16.0
Ca = 40.1
74.1
Fe2(SO4)3
O = 16.0
S = 32.1
Fe = 55.8
399.9
Iron(III) sulphate
(1 x 5)
79 a) Dative covalent bond
(1)
b) The phosphorus atom
(1)
supplies both bonding electrons to the hydrogen ion.
80
Giant
ionic structure
copper(II) sulphate,
magnesium fluoride,
sodium sulphide
126
Giant
covalent structure
Simple
molecular structure
diamond,
quartz
carbon dioxide,
chlorine,
nitrogen
(1)
Gaint
metallic structure
copper,
sodium
(0.5 x 10)
81
Forces of attraction
covalent bond
Between carbon atoms in diamond
Between carbon dioxide molecules in dry ice
van der Waals’ forces
Between particles in calcium nitride
ionic bond
Between particles in magnesium
metallic bond
Between carbon and oxygen atoms in a
carbon dioxide molecule
covalent bond
82 a)
Conductor, with decomposition
at the electrodes
Non-conductor
molten sulphur
(1)
liquid hydrogen chloride
(1)
molten potassium chloride
(1)
(1 x 5)
Conductor, but without
decomposition
molten potassium
b) molten potassium chloride — mobile potassium ions and chloride ions
molten potassium — mobile electrons
(1)
(1)
(1)
Structured questions
83 a) The relative atomic mass of an element is the weighted average relative isotopic mass of all the naturally
occurring isotopes of that element
(1)
on the
b)
12
C = 12.00 scale.
(1)
107 x 55 + 109 x 45
100
= 107.9
(1)
(1)
c) Isotopes of silver have the same chemical properties.
(1)
Hence it is impossible to separate the isotopes of silver by chemical means.
(1)
84 a) Isotopes are different atoms of an element which have the same number of protons
(1)
but a different number of neutrons.
b) Let the relative abundance of
28.09 =
28
Si and
(1)
29
Si be y% and (96.9 – y)% respectively.
28 x y + 29 x (96.9 – y) + 30 x 3.1
100
y = 94.1
c) Making semi-conductors
(1)
(1)
(1)
127
d)
LFZ
DBSCPOBUPN
TJMJDPOBUPN
(1 mark for the correct arrangement of atoms; 1 mark for the correct labelling of silicon and carbon
atoms)
(2)
e)
$M
$M
4J
$M
$M
(1)
f) Silicon carbide has a giant covalent structure while silicon tetrachloride has a simple molecular
structure.
(1)
To melt silicon carbide, a lot of heat is needed to overcome the strong covalent bonds between the
atoms.
(1)
Weak van der Waals’ forces exist between molecules of silicon tetrachloride. Little heat is needed to
separate the molecules.
(1)
85 a) • Both have 5 protons and 5 electrons.
11
(1)
• Both have the same number of protons and electrons.
(1)
•
10
(1)
B has 5 neutrons while
B has 6 neutrons.
OR
•
10
B and
11
B have different number of neutrons.
b) The weighted average relative isotopic mass of all the naturally occurring isotopes of that element
on the
128
12
C = 12.00 scale.
(1)
(1)
(1)
c)
d)
10 x 19.7 + 11 x 80.3
100
= 10.8
11
BF3 would give steamy fumes because the chemical propertries of isotopes are the same.
(1)
(1)
(1)
e) Boron is a semi-conductor.
(1)
f) i) The bond pair electrons are provided by the F– ion.
(1)
ii) Dative covalent bond
86 a) Its atomic number
b) Atomic size / metallic character of elements
c) Sodium and chlorine
d) i) Sodium and potassium react with cold water vigorously.
Hydrogen is evolved. / An alkaline solution is formed.
ii) Chlorine and fluorine react with metal to form salts.
e) Potassium and fluorine
(1)
(1)
(1)
(0.5, 0.5)
(1)
(1)
(1)
(0.5, 0.5)
f) Silicon has the highest melting point.
(1)
Silicon has a giant covalent structure.
(1)
To melt silicon, a lot of heat is needed to overcome the strong covalent bonds between the atoms. (1)
g) x = 18
y = 8
h) i) Rubidium is more reactive.
ii) It should be stored in paraffin oil.
87 a) Boron / silicon
b) Nitrogen and oxygen
(0.5)
(0.5)
(1)
(1)
(1)
(0.5, 0.5)
c) An argon atom has 8 electrons in its outermost shell. A special stability is obtained when this happens.
(1)
An argon atom has very little tendency to share electrons with other argon atoms. Therefore argon is
monoatomic.
(1)
d) Alkali metals
(1)
e) i) They have the same number of outermost shell electrons.
(1)
ii) They have different number of occupied electron shells.
(1)
129
f) The reactivity of Group II elements increases down the group.
(1)
g) There is a gradual increase in the melting / boiling point of the elements.
(1)
There is a gradual change in the intensity of the colour of the elements.
(1)
h) Aluminium has a giant metallic structure. It consists of tightly packed positive ions surrounded by a sea
of delocalized electrons.
(1)
The attractive forces between the electrons and positive ions hold the particles of aluminium together.
(1)
Fm
Fm
Fm
Fm
Fm
Fm
Fm
Fm
Fm
Fm
Fm
m
F
Fm
Fm
88 a) Increasing atomic number
(1)
(1)
b) They have the same number of occupied electron shells.
(1)
c) Lithium / Li
(1)
d) Atoms of Group 0 elements have stable electronic structure.
(1)
e) i) Potassium / K
(1)
ii) It should be stored in paraffin oil.
(1)
iii) Flammable / corrosive
(1)
f) Silicon / Si
(1)
g) Making bleach / hydrochloric acid / organic solvents
(1)
h) Lower than argon
(1)
89 a) f, h
(0.5, 0.5)
b) Any four of the following:
• Melting / boiling point usually high (1) • Shiny appearance (1) • Good conductor of electricity (1)
• Good conductor of heat (1)
• Ductile / malleable (1)
• Lose electrons / form positive ions (1)
130
c) 3
(1)
d) d
(1)
e) e
(1)
f) i) Alkali metals
(1)
ii) Any two of the following:
• Wear safety glasses. (1)
• Use forceps. (1)
• Use a safety screen. (1)
g) Noble gases
(1)
h) d and h
(0.5, 0.5)
i)
G
.H
m
m
C
/
G
PS
m
.H
m
C
/
G
.H
(1)
90 a) Nichrome / graphite
(1)
b) A reddish brown gas evolves.
(1)
c) A white shiny solid deposits on the electrode.
(1)
d) Inside a fume cupboard
(1)
The reddish brown gas (bromine) evolved is toxic.
(1)
e) The light bulb gradually goes out.
(1)
As the temperature drops, movement of ions in molten lead(II) bromide slows down. Therefore a smaller
current flows through the external circuit.
(1)
When the molten lead(II) bromide becomes solid, there are no mobile ions. Hence no current flows through
the external circuit.
(1)
91 a) For the conduction of electricity
(1)
b) i) Purple
ii)
(1)
DPMPVSFEQBUDI
QPTJUJWFFMFDUSPEF
OFHBUJWFFMFDUSPEF
m
EDQPXFSTVQQMZ
c) The coloured patch would move towards the new position of the positive electrode
because the negative permanganate ions would be attracted towards the positive electrode.
(2)
(1)
(1)
131
92 a) Any three of the following:
• Lithium fizzes / produces a gas. (1) • The universal indicator turns blue / purple (alkaline colour). (1)
• The water level in the test tube goes down (or gas fills the test tube). (1)
• Lithium moves around on the surface of water. (1) • Lithium dissolves. (1)
b) lithium + water
lithium hydroxide + hydrogen
(1)
c) Caesium is more reactive.
(1)
d) i) The melting point decreases as the atomic number increases
(1)
and the rate of decrease slows down.
ii) 26 °C
e) i) C
(1)
(1)
(1)
ii) The reactivity of Group I elements increases as we move down the group (i.e. potassium is more
reactive than lithium).
(1)
The reactivity of Group VII elements decreases as we move down the group (i.e. chlorine is more
reactive than iodine).
(1)
93 a) i) They have the same number of outermost shell electrons.
ii) They have different number of occupied electron shells.
b) i) Calcium sinks in water. / Calcium dissolves.
ii) Calcium is covered by a layer of calcium oxide.
Reaction between calcium and water starts only when the oxide layer dissolves.
c) x = 18
y = 2
(1)
(1)
(1)
(1)
(0.5)
(0.5)
d) Strontium is more reactive than calcium.
(1)
e) i) Isotopes are different atoms of an element which have the same number of protons
(1)
but a different number of neutrons.
ii) The chemical properties of strontium are similar to those of calcium.
Thus strontium can replace some of the calcium required.
132
(1)
(1)
(1)
(1)
94 a)
m
$M
.H
m
$M
(1)
b) i)
$M
4
$M
(1)
ii) Liquid
c) Sodium chloride and magnesium chloride have giant ionic structures.
(1)
(1)
To melt them, a lot of heat is needed to overcome the strong ionic bonds between the ions. Hence
sodium chloride and magnesium chloride have high melting points.
(1)
Phosphorus trichloride and sulphur dichloride have simple molecular structures.
(1)
The attractive forces between the molecules are weak. Little heat is needed to separate the molecules.
Hence phosphorus trichloride and sulphur dichloride have low melting points.
(1)
d) Magnesium chloride conducts electricity in molten state or aqueous solution but not in solid state.
(1)
In solid state, the ions in magnesium chloride are held together by strong ionic bonds. They are not free
to move.
(1)
The ions become mobile in molten state or aqueous solution.
e) Sodium chloride is soluble in water.
(1)
(1)
Strong attractive forces exist between ions in sodium chloride and water molecules.
(1)
These forces cause the ions to move away from the solid and go into the water.
(1)
95 a) fluorine
(1)
chlorine
(1)
calcium
(1)
b) Alkaline earth metals
(1)
133
c)
9
:
'
PS
$M
(1)
d)
m
m
9
'
;
PS
m
$B
m
9
'
(1)
e) The melting point of Q is higher than that of P.
(1)
To melt Q, a lot of heat is needed to overcome the strong ionic bonds between the ions.
(1)
The attractive forces between the molecules of P are weak. Little heat is needed to separate the
molecules.
(1)
f)
9
/
9
'
/
'
PS
9
'
(1)
g)
;
$B
m
m
/
/
;
PS
m
$B
m
/
/
;
$B
(1)
96 a)
134
Number of
Species
Atomic
number
Mass
number
protons
electrons
neutrons
A
8
16
8
8
8
B
8
18
8
8
10
C
8
16
8
10
8
D
9
19
9
9
10
E
12
24
12
10
12
F
12
24
12
12
12
(0.5 x 12)
b) They are isotopes.
(1)
c) C is an anion of A.
(1)
d) E is a cation of F.
(1)
e)
%
%
(1)
f) i)
%
"
%
(1)
ii)
'
m
"
(1)
iii) The melting point of Y is higher than that of X.
To melt Y, a lot of heat is needed to overcome the strong ionic bonds between the ions.
(1)
(1)
The attractive forces between the molecules of X are weak. Little heat is needed to separate the
molecules.
(1)
97 a) Ionic bond
(1)
b) Covalent bond
(1)
c) i) Dative covalent bond
(1)
ii) The nitrogen atom supplies both bonding electrons to the hydrogen ion.
d) Ammonium chloride is soluble in water.
(1)
(1)
Strong attractive forces exist between ions in ammonium chloride and water molecules.
(1)
These forces cause the ions to move away from the solid and go into the water.
(1)
e) Dissolve the sample in water.
(1)
Then add excess dilute nitric acid, followed by silver nitrate solution.
(1)
A white precipitate forms.
(1)
135
98 a) An atom of argon has 8 electrons in its outermost shell.
(1)
A special stability is obtained when this happens. Hence argon seldom forms compounds with other
elements.
(1)
b) Atoms of magnesium and calcium have the same number of outermost shell electrons.
(1)
c) A sulphur atom has an electronic arrangement 2,8,6. It obtains a stable electronic arrangement (2,8,8)
by gaining two electrons.
(1)
d) The ions in copper are packed closely and the metallic bonds holding them together are very strong. (1)
To melt a piece of copper, a lot of heat is needed to overcome the strong attractive forces. Hence copper
has a high melting point.
(1)
e) Quartz has a giant covalent structure.
(1)
To melt quartz, a lot of heat is needed to overcome the strong covalent bonds between the atoms. Hence
quartz has a high melting point.
(1)
Carbon dioxide has a simple molecular structure.
(1)
The attractive forces between the molecules are weak. Little heat is needed to separate the molecules.
Hence carbon dioxide has a low boiling point.
(1)
99 a) sodium chloride
(1)
argon
(1)
iodine
(1)
b) Van der Waals’ forces
(1)
c) i) X is hard
(1)
due to the strong ionic bonds between the ions. Relative motion of the ions is restricted.
ii) X does not conduct electricity in solid state but it does in molten state.
(1)
(1)
In solid state, the ions in X are held together by strong ionic bonds. They are not free to move. (1)
The ions become mobile in molten state.
d) i) Z is slightly soluble in water.
(1)
(1)
The weak attractive forces between molecules of Z and water are not strong enough to overcome the
attractive forces between water molecules.
(1)
ii) Z is very soluble in non-aqueous solvents.
(1)
The attractive forces between molecules of non-aqueous solvents are similar to those between molecules
of Z. Hence molecules of Z and non-aqueous solvents mix together easily.
(1)
iii) Z does not conduct electricity
becsuse it does not contain mobile electrons or ions.
136
(1)
(1)
100 a) Group I
(1)
b) Store X in paraffin oil.
(1)
c) i) XBr
(1)
ii) The compound is not volatile.
(1)
Its ions are held together by strong ionic bonds.
(1)
A lot of heat is needed to overcome the strong ionic bonds.
(1)
iii) The compound conducts electricity in molten state or aqueous solution but not in solid state.
(1)
In solid state, the ions in the compound are held together by strong ionic bonds. They are not free
to move.
(1)
The ions become mobile in molten state or aqueous solution.
101 a) Allotropes are two (or more) forms of the same element
in which the atoms or molecules are arranged in different ways.
b) Diamond is insoluble in water.
(1)
(1)
(1)
(1)
This is because the atoms are held together by strong covalent bonds. It is very difficult to separate the
atoms.
(1)
c) As a stone cutter
(1)
d) In graphite, the layers of carbon atoms are held by weak van der Waals’ forces. The layers can easily
slide over each other. Hence graphite is quite soft.
(1)
In diamond, each carbon atom is bonded to other carbon atoms by strong covalent bonds. Relative motion
of the atoms is restricted. Hence diamond is very hard.
(1)
e) Graphite has a layered structure. Weak van der Waals’ forces exist between the layers.
The layers can easily slide over each other.
(1)
(1)
Hence graphite has a slippery feel.
f) Graphite is a good conductor of electricity.
(1)
Graphite has a layered structure. Within each layer, each carbon atom uses three outermost shell electrons
in forming covalent bonds with three other atoms.
(1)
The remaining electron is delocalized between the layers of carbon atoms.
(1)
Graphite is a good conductor of electricity due to the presence of delocalized electrons.
137
102 a) Giant covalent structure
(1)
b) The attractive forces between carbon dioxide molecules are weak. Little heat is needed to separate the
molecules. Hence carbon dioxide has a low boiling point.
(1)
Diamond consists of a network of covalent bonds. A lot of heat is needed to overcome the strong covalent
bonds between the atoms. Hence diamond has a high melting point.
(1)
c) i)
WBOEFS8BBMTGPSDFT
LFZ
DBSCPOBUPN
(1 mark for the hexagonal arrangement of atoms; 1 mark for labelling the van der Waals’ forces
between the layers of atoms)
(2)
ii) When graphite is pressed onto a peice of paper, the layers of atoms slide over each other
and flake off easily onto the paper.
(1)
(1)
103 a) Giant covalent structure
(1)
b)
LFZ
PYZHFOBUPN
TJMJDPOBUPN
(1 mark for the correct arrangement of atoms; 1 mark for the correct labelling of silicon and oxygen
atoms)
(2)
c)
$M
$M
4J
$M
$M
(1)
138
d) To melt silicon dioxide, a lot of heat is needed to overcome the strong covalent bonds between the
atoms.
(1)
The attractive forces between silicon tetrachloride molecules are weak. Little heat is needed to separate
the molecules.
(1)
e)
70 x 24.4 + 72 x 32.4 + 74 x 43.2
100
= 72.4
(1)
(1)
104 a) i)
m
m
m
m
(1)
ii) 6
(1)
b) i) The particles in solid sodium metal are held together by a ‘sea’ of mobile electrons.
(1)
ii) The particles in solid sodium chloride are held together by ionic bonds.
(1)
iii) The ionic bonding in sodium chloride is stronger / requires more heat to break than the metallic
bonding in sodium.
(1)
c) Solid sodium conducts electricity but solid sodium chloride does not.
(1)
Solid sodium contains mobile electrons
(1)
but the ions in solid sodium chloride are not free to move.
(1)
d) Ions in sodium are packed in layers.
(1)
As the metal is struck by a hammer, the ion layers slide through the ‘sea’ of electrons to new positions.
The metal does not break because the ions are still bound together by the ‘sea’ of electrons.
(1)
As a result, sodium is malleable.
105
Element
p
q
r
s
t
u
v
w
Boiling point (°C)
2 480
3 930
4 830
–196
–183
–190
–246
890
Answer:
Be
B
C
N
O
F
Ne
Na
a) Element r has the highest boiling point and
(1)
a sudden drop in boiling point occurs from r to s.
(1)
Hence r is carbon while s is nitrogen.
(1)
b) The attractive forces between the molecules of t are weak.
Little heat is needed to separate the molecules.
(1)
(1)
139
c) q
(1)
d) Element w is stored in paraffin oil
(1)
as it is very reactive.
(1)
e) In advertising signs
(1)
f)
'
0
PS
'
V
U
V
(1)
g)
/B
X
m
0
m
PS
U
/B
X
(1)
h) The melting point of Y is higher than that of X.
To melt Y, a lot of heat is needed to overcome the strong ionic bonds between the ions.
(1)
(1)
The attractive forces between the molecules of X are weak. Little heat is needed to separate the
molecules.
(1)
106 a) The outermost shell electrons of each magnesium atom are free to move randomly in magnesium.
Thus, magnesium consists of positively charged ions surrounded by a ‘sea’ of electrons.
(1)
(1)
Magnesium is a good conductor of electricity due to the movement of mobile electrons in the metal.
(1)
b) Each aluminium atom has three outermost shell electrons while a magnesium atom has two.
There are more delocalized electrons in aluminium.
(1)
(1)
So, the electrical conductivity of aluminium is higher.
c) Ions in a metal are packed in layers.
(1)
As the metal is struck by a hammer, the ion layers slide through the ‘sea’ of electrons to new positions.
The metal does not break because the ions are still bound together by the ‘sea’ of electrons.
(1)
As a result, the metal is ductile.
d) Molten sulphur cannot conduct electricity
because it does not contain mobile electrons or ions.
140
(1)
(1)
107 a) i) Group 0
ii) Not yet discovered at that time
(1)
(1)
b) Exists as diatomic molecules / exists as a gas at room conditions / any other general property of nonmetals
(1)
c) Any two of the following:
• Many elements in the groups have very dissimilar properties, e.g. K and Cu.
(2)
• Two elements were put in one place, e.g. Ce and La.
(2)
• Metals and non-metals were mixed up, e.g. Cl and Co in the same group.
(2)
d) Any two of the following:
• Elements with similar properties were grouped together.
(1)
• Gaps left for elements to be added when discovered.
(1)
• A new group created / iron, cobalt and nickel put in a group.
(1)
• Metals and non-metals were separated.
(1)
e) In order of increasing atomic number
108 a) Allotropes are two (or more) forms of the same element
in which the atoms or molecules are arranged in different ways.
b) Buckminsterfullerene has a simple molecular structure.
Buckminsterfullerene is soluble in non-aqueous solvents.
(1)
(1)
(1)
(1)
(1)
It can be deduced that the attractive forces between molecules of buckminsterfullerene are similar to
those between molecules of non-aqueous solvents.
(1)
Hence it can be concluded that buckminsterfullerene has a simple molecular structure.
c) The melting point of diamond is higher than that of buckminsterfullerene.
(1)
Diamond has a giant covalent structure. The carbon atoms are held together by strong covalent bonds.
(1)
There are weak van der Waals’ forces between the buckminsterfullerene molecules.
(1)
More heat is needed to break the strong covalent bonds between atoms in diamond. Hence diamond
has a higher melting point.
d) The carbon atoms are held together by strong covalent bonds.
(1)
e) i) In graphite, the carbon atoms are arranged in flat parallel layers.
(1)
Within each layer, each carbon atom uses three electrons in forming covalent bonds with three other
carbon atoms.
(1)
The remaining outermost shell electron of each carbon atom is delocalized between the layers of
carbon atoms.
(1)
141
ii) Graphite has a layered structure. Weak van der Waals’ forces exist between the layers.
The layers of atoms can slide over each other easily.
(1)
(1)
f) Adding a non-aqueous solvent to each solid separately, buckminsterfullerene is soluble while graphite is
insoluble.
(1)
Buckminsterfullerene has a simple molecular structure and is soluble in non-aqueous solvents.
(1)
Graphite has a giant covalent structure. It is insoluble in most solvents.
(1)
109 a) A mixture consists of two or more pure substances
(1)
which have not been chemically joined together.
(1)
b)
m
0
/
0
0
(1)
c) The pressure from the expanding gases propels the shell into the air
(2)
d) i) Potassium sulphide
(1)
ii) K2S
(1)
e) i) Alkaline earth metals
(1)
ii) Sr(NO3)2
(1)
f) Sodium nitrate / any sodium compound
(1)
110 Sodium, magnesium and aluminium are metals. The strength of the metallic bond depends on the number
of delocalized electrons in the metal structure.
(1)
Sodium has one outermost shell electron per atom, magnesium has two while aluminium has three. The
strength of metallic bond and hence the melting point increase from sodium to aluminium.
(1)
Silicon has a giant covalent structure. Each silicon atom is covalently bonded to four other silicon atoms.(1)
To melt silicon, a lot of heat is needed to overcome the strong covalent bonds between the atoms. Hence
it has a very high melting point.
(1)
Phosphorus, sulphur and chlorine exist as simple molecules. The molecules are attracted to one another by
weak van der Waals’ forces.
(1)
Little heat is needed to separate the molecules. Hence they have low melting points.
(1)
(3 marks for organization and presentation)
142
111 Atoms of Group VI elements have six outermost shell electrons. They can obtain the electronic arrangements
of atoms of noble gases by gaining or sharing electrons.
(1)
Oxygen is a Group VI element. Take the combination of oxygen and sodium as an example. An oxygen atom
has an electronic arrangement 2,6. It tends to gain two electrons to obtain the electronic arrangement of a
stable neon atom (2,8).
(1)
Sodium is a Group I element. A sodium atom has an electronic arrangement 2,8,1. It tends to lose one
electron to obtain the electronic arrangement of a stable neon atom (2,8).
(1)
When sodium and oxygen react, two sodium atoms would combine with one oxygen atom.
(1)
Take the combination of oxygen and carbon as another example. An oxygen atom has an electronic arrangement
2,6 while that of a carbon atom is 2,4.
Both atoms require electrons to obtain the electronic arrangement of a stable neon atom (2,8). They achieve
that by sharing outermost shell electrons.
(1)
One carbon atom forms a double bond with each of the two oxygen atoms.
(1)
(3 marks for organization and presentation)
112 The melting point of potassium chloride is higher than that of silicon tetrachloride.
(1)
Potassium chloride has a giant ionic structure.
To melt potassium chloride, a lot of heat is needed to overcome the strong ionic bonds between the ions.
Hence potassium chloride has a high melting point.
(1)
Silicon tetrachloride has a simple molecular structure.
The attractive forces between the molecules are weak. Little heat is needed to separate the molecules. Hence
silicon tetrachloride has a low melting point.
(1)
Potassium chloride conducts electricity in molten state or aqueous solution while silicon tetrachloride does
not conduct electricity.
(1)
In the solid state, the ions in potassium chloride are held together by strong ionic bonds and are not free
to move.
The ions become mobile when potassium chloride is in molten state or aqueous solution. Hence potassium
chloride can conduct electricity under these conditions.
(1)
Silicon tetrachloride does not conduct electricity because it does not contain mobile electrons or ions.
(1)
(3 marks for organization and presentation)
143