Topic Microscopic World I / Microscopic World (Combined Science)home.sbc.edu.hk/~chem/Supp Ex...

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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 Molten sulphur contains mobile molecules only. There are no mobile ions or electrons. It does not conduct electricity.

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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 A hydrogen atom (11H) 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.

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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

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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


47 D Particle Atomic number Mass number

Number of neutrons= mass number – atomic number

316

2S 16 32 16

211

3Na 11 23 12

212

4Mg 12 24 12

2148Si 14 28 14

315

1P 15 31 16

∴ 3151P contains the same number of neutrons as 3

162S.

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.

50 B Relative atomic mass of lithium = 6 x 7.4 + 7 x 92.6

100

= 6.93

51 B Relative atomic mass of X = 85 x 72.1 + 87 x 27.9

100 = 85.6

52 C Relative atomic mass of X = 189 x 25 + 190 x 30 + 192 x 45

100 = 190.7

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53 D Let the relative abundance of 79X and 81X be (100 – y)% and y% respectively.

Relative atomic mass of X = 79 x (100 – y) + 81 x y

100 = 79.9

∴ 7 900 – 79y + 81y = 7 990

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 69X and aX are 65.0% and 35.0% respectively.

Relative atomic mass of X = 69 x 65 + a x 35

100 = 69.7

∴ 4 485 + 35a = 6 970

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 elementMetal / 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.

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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 6270Co atom = mass number – atomic number

= 60 – 27 = 33

(2) The atomic number of 6270Co is 27. Hence a 6

270Co 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 3162S atom

contains 16 protons.

72 C 5244X and 5

264Y represent different elements. They are NOT isotopes.

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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

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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.

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.

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36 T

37 F 35Cl and 37Cl have the same electronic arrangement and hence the same chemical properties.

38 F A helium atom has a duplet structure in its outermost shell.

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.


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 arrangementof atom

Number of occupiedelectron 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.

Li Be

Na Mg

K Ca

B C N O F

Al Si P S Cl

H

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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 Option Element Group number

A Boron III

B Bromine VII

C Chlorine VII

D Silicon IV

49 A Option Element Group

Aargon 0

neon 0

Bcarbon IV

chlorine VII

Ccalcium II

potassium I

Dmagnesium II

sodium I

Argon and neon belong to the same group.

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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

A7 2,5 V

13 2,8,3 III

B9 2,7 VII

17 2,8,7 VII

C11 2,8,1 I

18 2,8,8 0

D14 2,8,4 IV

20 2,8,8,2 II

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 Difference in atomic number Electronic arrangement of atom

B 5 8

18

2,3

Al 13 2,8,3

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.

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54 A Option Element

Atomic number


Period to which the element belongs

Aargon 18 2,8,8 3

aluminium 13 2,8,3 3

Bberyllium 4 2,2 2

silicon 14 2,8,4 3

Cchlorine 17 2,8,7 3

nitrogen 7 2,5 2

Dphosphorus 15 2,8,5 3

oxygen 8 2,6 2

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.

Li Be

Na Mg

K Ca

B C N O F

Al Si P S Cl

H

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.

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57 C Option A — The reactivity of Group I elements increases down the group.

Hence potassium is more reactive than sodium.

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.

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.

Hence fluorine is more reactive than chlorine.

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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.

K

Na Mg

Be

Hence Z (potassium) is the most reactive metal.

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.

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.

67 A Noble gas Atomic number Difference in atomic number Electronic arrangement of atom

He 2 8

8

18

2

Ne 10 2,8

Ar 18 2,8,8

Kr 36 2,8,18,8

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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.

Li Be

Na Mg

K Ca

B C N O F

Al Si P S Cl

H

The reactivity of Group I elements increases down the group. Hence element g is the most reactive 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 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.

DThe reactivity of Group I elements increases down the group. Hence rubidium is more reactive than potassium.

72 C Option Description

AStrontium 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.

CStrontium 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.

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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.

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

ALi+ 2 2

H 1 1

BO2– 2,8 10

Cl– 2,8,8 18

CNa+ 2,8 10

S2– 2,8,8 18

DNe 2,8 10

F– 2,8 10

Ne and F– have the same number of electrons.

76 C

SpeciesNumber of

protons neutrons electrons

5266Fe 26 30 26

Fe3+ 26 30 23

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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 X is a Group II element, an alkaline earth metal.

D 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 IV V VI VII

Element Lithium Beryllium BoronCarbon

(graphite)Carbon

(diamond)Nitrogen Oxygen Fluorine

Atomic radius (pm) (1 pm = 10–12 m)

152 112 83 77 75 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 IV V VI VII

Element Lithium Beryllium BoronCarbon

(graphite)Carbon

(diamond)Nitrogen Oxygen Fluorine

Melting point (°C) 180 1 280 2 030 3 730 3 500 –210 –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

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(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) 4180X 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.

Hence the chemical reactivity is in the order lithium < sodium < potassium.

(3) The melting point of Group I elements decreases down the group.

Hence the melting point is in the order lithium > sodium > potassium.

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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.

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.

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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.

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.

Hence strontium is more reactive than calcium.

93 C (1) The reactivity of halogens decreases down the group.

Hence the reactivity of halogens is in the order chlorine > bromine > iodine.

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(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.


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.

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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.

(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.

(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.

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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.

Elements in the second period of the periodic table have 2 occupied electron shells in their atoms.

52

105 B Across the third period of the periodic table, the atomic size of the elements decreases gradually.

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

Type of element

metals metalloid non-metals

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.

53

108 C The reactivity of halogens decreases down the group.

The number of occupied electron shells in atoms of halogens increases down the group. Hence the atomic size also increases down the group.

Halogen Electronic arrangement of atom Atomic radius (pm) (1 pm = 10–12 m)

Fluorine 2,7 72

Chlorine 2,8,7 99

Bromine 2,8,18,7 114

Iodine 2,8,18,18,7 133

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.

54

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.

55

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.

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.

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.

56

Electron diagram of the compound formed:

Ca

(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.


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 Metallic bond is a type of bond in which positive metal ions are held together by a ‘sea’ of mobile electrons.

57


21 A Alcohol is made up of carbon, hydrogen and oxygen. Compounds made up of non-metals are non-conductors.

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.

58

28 C Electron transfer during the reaction between potassium and oxygen:

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:


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.

59

Electron transfer during the reaction between X and Y:


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:


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.

60

33 B Option Compound Chemical formula Electron diagram

A aluminium oxideAl2O3

The chemical formula is not in the form of XY.

—

B magnesium oxide MgO

C lithium fluoride LiF

D sodium chloride NaCl

∴ 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.



The chemical formula of the compound formed is Mg3N2, i.e. Mg3X2.

61

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.



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.

62

The reactivity of Group VII elements decreases down the group. Hence fluorine is more reactive than chlorine.

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.

63

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.

49 D A calcium ion (Ca2+) 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.

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 1 M Cl

3 Cross multiply the numbers and the symbols. 3 1 M Cl = M1 = Cl3

4 Combine the symbols. MCl3

64

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.

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.

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.

Hence chlorine is more reactive than astatine.

65

Option C — Astatine is in Period 6 of the periodic table.

Option D — A calcium ion (Ca2+) 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.

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.

Option D — The cation of X (X+) carries 1 positive charge while the oxide ion (O2–) 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.

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.



(2) An atom of X has 4 occupied electron shells. Hence X is in Period 4 of the periodic table.

66

57 A (1) The reactivity of Group I elements increases down the group.

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.

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.


Hence barium is more reactive than calcium.

67

(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.

60 B (1) A strontium ion (Sr2+) carries 2 positive charges while a carbonate ion (CO32–) carries 2 negative

charges. The simplest ratio of Sr2+ to CO32– 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.



The chemical formula of the compound is Mg3N2, i.e. Mg3X2.

68

62 D Compound Chemical formula Electron diagram

Lithium oxide Li2O

Potassium sulphide K2S

Sodium oxide Na2O

∴ 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).

67 C A helium atom has 2 outermost shell electrons.

69

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



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

70

True or false

11 T

12 F Nitrogen exists as diatomic molecules.

≡


13 T Carbon and silicon are non-metals. They combine to form a covalent compound (silicon carbide, SiC).

Structure of silicon carbide:

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:

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

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16 T In a SiCl4 molecule, one Si atom forms a single bond with each of four Cl atoms.

Electron diagram of SiCl4:


17 T Covalent bonding exists in both iodine and oxygen.


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:


72

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.


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.


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.



73

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.


Option C — In a hydrogen chloride molecule, one hydrogen atom forms a single bond with one chlorine atom.


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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:


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.

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).

75

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).

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

Electron diagram (only electrons in the outermost shells are shown)

Number of pairs of bond pair electrons

C2H4 6

CO2 4

H2S 2

PCl3 3

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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

NH3 1

SCl2 2

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:


77

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.


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)

magnesium and fluorine

lithium and oxygen

chlorine and fluorine

chlorine and oxygen

78

39 B Option Substance Bonding type

A

aluminium

metallic bondingmercury

sodium

B

calcium chloride ionic bonding

hydrogen chloride covalent bonding

silver chloride ionic bonding

Ccarbon dioxide

nitrogenoxygen

covalent bonding

Diodine

methanesulphur dioxide

covalent bonding

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:


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:

StepCompound formed between X

and calcium

1 Write down the symbols of ions in the compound. Ca X


2 3 Ca X

3 Cross multiply the numbers and the symbols. 2 3 Ca X = Ca3 = X2

4 Combine the symbols. Ca3X2

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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.

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.


Element lithium beryllium boron carbon nitrogen oxygen fluorine

forms an ionic oxide

forms a covalent

oxide

43 D In a piece of metal, metal ions are packed tightly together in a regular pattern to form a giant structure.

80

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 theoutermost shells are shown)

With singlebonds only?

(1) CH4 yes

(2) CO2 no

(3) H2O yes

81

49 D Option Molecule

Electron diagram (only electrons in theoutermost shells are shown)

With multiple bond(s)?

(1) CO2 yes

(2) N2 yes

(3) HCN 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.


(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:


(3) Both chlorine and bromine can react with sodium sulphite solution.

e.g. aqueous bromine + sodium sulphite sodium sulphate + hydrogen bromide

82

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:


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.

83

(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.

(3) Each atom of X in the compound has 2 lone pairs of electrons.

54 B The electron diagram of ammonium chloride is shown in the question.


(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.


(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).

84

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:


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:


85

59 D (1) Group II elements are less reactive than Group I elements.

K

Na Mg

Be

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:


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.

86

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:


65 C Calcium carbonate is an ionic compound.

Each carbonate ion (CO32–) is a group of four atoms held together by covalent bonding.

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

87

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:

15 F Silicon carbide has a giant covalent structure.

Structure of silicon carbide:

16 T Electron diagram of silane (SiH4):


88

17 F Carbon disulphide (CS2) has a simple molecular structure.

Electron diagram of carbon disulphide:


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:

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:

89

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:

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:

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.

90

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.


33 B

34 A Structure of sodium chloride:

35 C Option Substance Electrical conductivity

A argon a non-conductor

B potassium a conductor; not chemically changed during conduction

C potassium fluoridean electrolyte; conducts electricity in molten state or aqueous solution, and decomposed by electricity during conduction

D tetrachloromethane a non-conductor

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.

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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:

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:

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.

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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:

45 D Structure of dry ice:

46 C Option Substance Structure

A diamond giant covalent structure

B mercury giant metallic structure

C nitrogen simple molecular structure

D quartz giant covalent structure

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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:


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 non-aqueous solvents.

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53 C Option

Atomic number of element

Name of elementCompound 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:


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.



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 Element Atomic number Name of element

X 8 oxygen

Y 12 magnesium

Z 17 chlorine

95

Option A — X (oxygen) and Z (chlorine) are non-metals. They combine to form a covalent compound.



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.



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

96

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:


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.

97

Electron diagram of the compound Z:


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 –82simple molecular

structurecovalent 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.

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.

98

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.

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 and1 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:

99

70 D Option Substance Structure

(1) oxygen simple molecular structure

(2) potassium oxide giant ionic structure

(3) silicon dioxide giant covalent structure

71 D 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

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:


The fluoride has a simple molecular structure.

100

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:


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:

(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.

101

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

Property of substanceGiant covalent Giant ionic Giant metallic

Melting point high high high

Electrical conductivity at room temperaturenon-conductor

(except graphite)non-conductor conductor

∴ 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:


(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

102

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:


(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


3 1 X Cl

3 Cross multiply the numbers and the symbols. 3 1 X Cl = X1 = Cl3

4 Combine the symbols. 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.



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:


103

(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.



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) 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.

104

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.

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.

105

Part B Topic-based exercise


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 3199K = 19 = number of protons in an atom

= number of electrons in an atom


4 D Option Species Number of electrons Number of neutrons

A94Be 4 5

B211

3Na+10 12

C2137Al 13 14

D317

5Cl– 18 18

∴ 3175Cl– contains equal numbers of electrons and neutrons.

5 B An atom of M loses 2 electrons to form a M2+ ion with 27 electrons.

∴ number of electrons in an atom of M = 29 = number of protons in an atom of M = atomic number of M = atomic number of Cu

6 C Relative atomic mass of Kr = 84 x 45.0 + 86 x 55.0

100

= 85.1

106

7 D Let the relative abundance of 126I and 127I be (100 – y)% and y% respectively.

126.9 = 126 x (100 – y) + 127 x y

100

y = 90.0

8 D Halogen Atomic number Difference in atomic number Electronic arrangement of atom

F 9 8

18

2,7

Cl 17 2,8,7

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 Electronic arrangement of atom

A

42X 2

210

0Y 2,8

B

162X 2,4

212

4Y 2,8,2

C

211

3X 2,8,1

317

5Y 2,8,7

D

418

0X 2,8,8

420

0Y 2,8,8,2

42X and 2

100Y 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.

107

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.

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.

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17 A Option Compound Electron diagram

A Na2O

B Li3N

C CaO

D SrCl2

∴ 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.

109

19 C Chlorine exists as diatomic molecules. It is a gas at room temperature and pressure. The molecules are arranged in a random way.

20 B The atom of X gains 3 electrons in order to obtain a stable electronic arrangement. A stable X3– ion is formed.

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.



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


3 2 E Z

3 Cross multiply the numbers and the symbols. 3 2 E Z = E2 = Z3

4 Combine the symbols. E2Z3

110

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.


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.

111

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.


30 C Dry ice has a simple molecular structure. Weak van der Waals’ forces hold the molecules together.

31 D Option Substance Bonding type

A

barium fluoride

magnesium fluoride

potassium fluoride

ionic bonding

B

carbon monoxide

sulphur dioxide

methane

covalent bonding

C

chromium

magnesium

nickel

metallic bonding

D

boron trichloride

silicon tetrachloride

sodium chloride

covalent bonding

covalent bonding

ionic bonding

32 A 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.

112

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:


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.

113

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.


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.



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.

114

40 C (1) Number of neutrons in 6318Ga = mass number – atomic mass

= 68 – 31 = 37

(2) The atomic number of the gallium isotope 6318Ga 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

Li+ 2

N3– 2,8

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.

(2) The melting point of Group VII elements increases down the group, i.e. increases with atomic number.

115

43 A (1) Across the third period of the periodic table, the atomic size of the elements shows a gradual decrease.

(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.



Type of element

metals metalloid non-metals

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.

44 A (1) Group I II III IV V VI VII



2,8,1 2,8,2 2,8,3 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.

116

(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.



form ionic chlorides form covalent chlorides

45 D (3) Potassium reacts with water to give an alkaline solution.

potassium + water potassium hydroxide + hydrogen

46 B (1) 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.

117

47 B Caesium (a metal) and bromine (a non-metal) react to form an ionic compound. Electron diagram of the compound:


(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:


50 D

118

51 D (3) Iodine changes from the solid state to the gaseous state directly when heated.

52 D (1) Element Colour

Chlorine greenish yellow

Bromine reddish brown

Iodine black

(2) The reactivity of halogens decreases down the group.

(3) The halogens exist as diatomic molecules. Van der Waals’ forces exist between the molecules.

53 D Electron diagram of hydrogen peroxide:

(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

119

54 A Electron diagram of calcium carbonate:


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:


The outermost shell electrons of atoms in a methane molecule are shared to form covalent bonds. There are NO delocalized electrons.

120

(3) Sodium consists of tightly packed positive ions surrounded by a sea of delocalized electrons.

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) 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.

121

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.

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.

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.

122

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

Percentage by mass of hydrogen in ammonia = 3 x 1.0

17.0 x 100%

= 17.6%

Percentage by mass of nitrogen in ammonia = 14.017.0

x 100%

= 82.4%

Short questions

71 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)

123

72 Atom

Atomic number

Massnumber

SymbolNumber of

protons neutrons electrons

Oxygen 8 16 816O 8 8 8

Sodium 11 23 1231Na 11 12 11

Aluminium 13 27 1273Al 13 14 13

Sulphur 16 32 132

6S 16 16 16

Chlorine 17 35 1357Cl 17 18 17

Potassium 19 39 139

9K 19 20 19

Calcium 20 40 240

0Ca 20 20 20

Iron 26 56 256

6Fe 26 30 26 (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) Species

Atomicnumber

Massnumber

Number of Electronicarrangementprotons neutrons electrons

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) Potassium ion 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) Sodium atom 11 23 11 12 11 2,8,1

(0.5 x 40)

b) Species (vi) & (viii) / nitride ion and fluoride ion (0.5 x 2)

c) Group IV; (1)

it has 4 electrons in its outermost shell. (1)

d) Species (ix) is the cation of species (x) / (ix) and (x) are the ion and atom of the same element. (1)

124

75 a) 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)

b) 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 (1 x 8)

76 Cation Anion Compound

Name Formula Name Formula Name FormulaColour of aqueous solution

ammonium NH4+ carbonate CO3

2– ammonium carbonate (NH4)2CO3 colourless

copper(II) Cu2+ nitrate NO3– copper(II) nitrate Cu(NO3)2 blue

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

aluminium Al3+ iodide I– aluminium iodide All3 colourless

chromium(III) Cr3+ chloride Cl– chromium(III) chloride CrCl3 green

sodium Na+ dichromate Cr2O72– sodium dichromate Na2Cr2O7 orange

zinc Zn2+ bromide Br– zinc bromide ZnBr2 colourless

(0.5 x 42)

125

77 a) i)

(1)

ii)

(1)

iii)

(1)

126

b) i)

(1)

ii)

(1)

iii)

(1)

78

Substance Chemical formulaRelative 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

Iron(III) sulphate Fe2(SO4)3

O = 16.0 S = 32.1 Fe = 55.8

399.9(1 x 5)

79 a) Dative covalent bond (1)

b) The phosphorus atom (1)

supplies both bonding electrons to the hydrogen ion. (1)

80 Giant ionic structure

Giant covalent structure

Simple molecular structure

Gaint metallic structure

copper(II) sulphate,magnesium fluoride,sodium sulphide

diamond,quartz

carbon dioxide,chlorine,nitrogen

copper,sodium

(0.5 x 10)

127

81 Forces of attraction

Between carbon atoms in diamond covalent bond

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(1 x 5)

82 a) Non-conductor

Conductor, with decomposition at the electrodes

Conductor, but without decomposition

molten sulphur (1)

liquid hydrogen chloride (1)

molten potassium chloride (1) molten potassium (1)

b) molten potassium chloride — mobile potassium ions and chloride ions (1)

molten potassium — mobile electrons (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 12C = 12.00 scale. (1)

b) 107 x 55 + 109 x 45

100 (1)

= 107.9 (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. (1)

b) Let the relative abundance of 28Si and 29Si be y% and (96.9 – y)% respectively.

28.09 = 28 x y + 29 x (96.9 – y) + 30 x 3.1100

(1)

y = 94.1 (1)

c) Making semi-conductors (1)

128

d)

(1 mark for the correct arrangement of atoms; 1 mark for the correct labelling of silicon and carbon atoms) (2)

e)

(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. (1)

• 10B has 5 neutrons while 11B has 6 neutrons. (1)

OR

• Both have the same number of protons and electrons. (1)

• 10B and 11B have different number of neutrons. (1)

b) The weighted average relative isotopic mass of all the naturally occurring isotopes of that element (1)

on the 12C = 12.00 scale. (1)

129

c) 10 x 19.7 + 11 x 80.3

100 (1)

= 10.8 (1)

d) 11BF3 would give steamy fumes because the chemical propertries of isotopes are the same. (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 (1)

86 a) Its atomic number (1)

b) Atomic size / metallic character of elements (1)

c) Sodium and chlorine (0.5, 0.5)

d) i) Sodium and potassium react with cold water vigorously. (1)

Hydrogen is evolved. / An alkaline solution is formed. (1)

ii) Chlorine and fluorine react with metal to form salts. (1)

e) Potassium and fluorine (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 (0.5)

y = 8 (0.5)

h) i) Rubidium is more reactive. (1)

ii) It should be stored in paraffin oil. (1)

87 a) Boron / silicon (1)

b) Nitrogen and oxygen (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)

130

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)

(1)

88 a) Increasing atomic number (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)

c) 3 (1)

d) d (1)

e) e (1)

131

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)

(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 (1)

ii)

(2)

c) The coloured patch would move towards the new position of the positive electrode (1)

because the negative permanganate ions would be attracted towards the positive electrode. (1)

132

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. (1)

ii) 26 °C (1)

e) i) C (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. (1)

ii) They have different number of occupied electron shells. (1)

b) i) Calcium sinks in water. / Calcium dissolves. (1)

ii) Calcium is covered by a layer of calcium oxide. (1)

Reaction between calcium and water starts only when the oxide layer dissolves. (1)

c) x = 18 (0.5)

y = 2 (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. (1)

ii) The chemical properties of strontium are similar to those of calcium. (1)

Thus strontium can replace some of the calcium required. (1)

133

94 a)

(1)

b) i)

(1)

ii) Liquid (1)

c) Sodium chloride and magnesium chloride have giant ionic structures. (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. (1)

e) Sodium chloride is soluble in water. (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)

134

c)

(1)

d)

(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)

(1)

g)

(1)

96 a) Species

Atomicnumber

Massnumber

Number of

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)

135

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)

(1)

iii) The melting point of Y is higher than that of X. (1)

To melt Y, a lot of heat is needed to overcome the strong ionic bonds between the ions. (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. (1)

d) Ammonium chloride is soluble in water. (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)

136

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. (1)

ii) X does not conduct electricity in solid state but it does in molten state. (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. (1)

d) i) Z is slightly soluble in water. (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 (1)

becsuse it does not contain mobile electrons or ions. (1)

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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. (1)

101 a) Allotropes are two (or more) forms of the same element (1)

in which the atoms or molecules are arranged in different ways. (1)

b) Diamond is insoluble in water. (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. (1)

The layers can easily slide over each other. (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.

138

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)

(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 (1)

and flake off easily onto the paper. (1)

103 a) Giant covalent structure (1)

b)

(1 mark for the correct arrangement of atoms; 1 mark for the correct labelling of silicon and oxygen atoms) (2)

c)

(1)

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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 (1)

= 72.4 (1)

104 a) i)

(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. (1)

Little heat is needed to separate the molecules. (1)

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c) q (1)

d) Element w is stored in paraffin oil (1)

as it is very reactive. (1)

e) In advertising signs (1)

f)

(1)

g)

(1)

h) The melting point of Y is higher than that of X. (1)

To melt Y, a lot of heat is needed to overcome the strong ionic bonds between the ions. (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. (1)

Thus, magnesium consists of positively charged ions surrounded by a ‘sea’ of electrons. (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. (1)

There are more delocalized electrons in aluminium. (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 (1)

because it does not contain mobile electrons or ions. (1)

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107 a) i) Group 0 (1)

ii) Not yet discovered at that time (1)

b) Exists as diatomic molecules / exists as a gas at room conditions / any other general property of non-metals (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 (1)

108 a) Allotropes are two (or more) forms of the same element (1)

in which the atoms or molecules are arranged in different ways. (1)

b) Buckminsterfullerene has a simple molecular structure. (1)

Buckminsterfullerene is soluble in non-aqueous solvents. (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)

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ii) Graphite has a layered structure. Weak van der Waals’ forces exist between the layers. (1)

The layers of atoms can slide over each other easily. (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)

(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)

143

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)


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)


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