3
3.1
3.2
3.3
3.4
3.5
Ionic (Electrovalent) Bonding
Covalent Bonding and Co-ordinate Bonding
Bond Properties and Intermolecular Forces
Metallic Bonding
Bonding and Physical Properties
3.1



Ionic (Electrovalent) Bonding
An ionic bond is an attraction between oppositely charged ions, which are formed by the transfer of
electrons from one atom to another.
A high difference of electronegativity between the two atoms is necessary for an ionic bond to form.
Ionic bonding is not possible between similar atoms.
In sodium chloride, each sodium atom transfers an electron to a chlorine atom, forming 𝑁𝑎+ and
𝐶𝑙 − ions. These two ions attract each other by electrostatic forces to form a stable compound.
Note that:
- Both species end up with the electronic configuration of the nearest noble gas.
- On the ‘dot-and-cross’ diagram, only the last shell of electrons needs to be shown.
3.2
Covalent Bonding and Co-ordinate Bonding
Covalent Bonding


A covalent bond is a pair of electrons shared between two atoms, usually of similar electronegativity.
In a normal covalent bond, each atom provides one of the electrons in the bond. The atoms achieve
the stable noble gas structure by sharing electrons.
A covalent bond holds the two atoms together because the pair of electrons is attracted to both
nuclei. The following are examples of dot-and-cross diagrams for some common molecules.
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

Most covalent bonds require large amount of energy to break due to the strong electrostatic forces
of attraction between shared pairs of electrons and the nuclei.
Electron pairs involved in covalent bonding are called bond pairs; those not involved are called lone
pairs. Examples:
Molecule
Dot-and-Cross Diagram
Electron Pairs
𝐶𝑂2
There are 4 bond pairs and 0 lone
pairs around the C atom
𝑆𝑂2
There are 4 bond pairs and 1 lone
pair around S
𝐻2 𝑂
There are 2 bond pairs and 2 lone
pairs around the O atom
𝐶𝐻2 = 𝐶𝐻2
There are 6 bond pairs and 0 lone
pairs around both C atoms
Table 3.1 Some dot-and-cross diagrams
Co-ordinate (Dative) Bonding




A dative covalent bond is a pair of electrons shared between two atoms, one of which provides both
electrons to the bond.
The two conditions for dative bonding to occur are:
1. The donor atom must have lone pairs in its outer shell; and
2. The acceptor atom must be short of the “octet”.
A dative covalent bond is represented by an arrow pointing away from the atom donating the lone
pair to the atom accepting it.
Example:
The formation of a dative bond in the ammonium ion:
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
The following are examples of molecules that have dative bonds:
Molecule or
Ion
𝑪𝑶
𝑵𝑶𝟐
𝑵𝟐 𝑶𝟒
𝑶𝟑 (ozone)
Bonding
Outline
Molecule or
Ion
𝑪𝑵−
𝑯𝑵𝑶𝟑
𝑨𝒍𝟐 𝑪𝒍𝟔
𝑨𝒍𝑪𝒍𝟒−
Bonding
Outline
Table 3.2 Some molecules with dative bonding



Although dative bonds are drawn with an arrow, there is no difference between a dative covalent
bond and an ordinary covalent bond.
Dative bonding also occurs in hydrated complex ions such as the hexa-aqua-aluminium ion,
𝐴𝑙(𝐻2 𝑂)3+
6 :
The 6 H2O molecules are bonded datively to the central 𝐴𝑙 3+ cation.
Shapes of Molecules

Molecular shapes and the angles between bonds can be predicted by the VSEPR (Valence Shell
Electron Pair Repulsion) theory which states that:
1.
The shape of a molecule or ion is determined by both the bond pairs and the lone pairs around
the central atom;
2.
The electron pairs spread out as far apart as possible so as to minimize repulsion between
them.
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3.
Lone pairs of electrons have a greater repulsive power than bond pairs, so their presence will
affect the bond angles as they push the bond pairs away. The order of repulsive power is:
lone pair - lone pair repulsion > lone pair - bond pair repulsion > bond pair - bond pair repulsion
Shape
Linear
2 bond directions, 0
lone pairs
Trigonal Planar
Examples of Molecules
CO
BeCl2
BH2+
NO+
2
CNO−
BF3
CO2−
3
3 bond directions, 0
lone pairs
C2 H4
N2 O4
Tetrahedral
4 bond directions, 0
lone pairs
NH4+
OH −
NO−
3
SO3
SO2−
4
CH4
C2 H6
H2 SO4
*tetrahedral around S,
bent around O atoms.
BrFO3
Octahedral
6 bond directions, 0
lone pairs
SF6
Bent/Non-Linear
2 bond directions, at
least 1 lone pair
PCl−
6
SO2
NO−
2
H2 O
NO2
O3
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Trigonal
Pyramidal
NH3
SO2−
3
N2 H4
H3 O+
C𝑙FO2
3 bond directions, 1
lone pair
AsBr3
Trigonal Bipyramidal
5 bond directions, 0
lone pairs
PF5
SbF5
Square Planar
4 bond directions, 2
lone pairs
XeF4
C𝑙F4−
Table 3.3 Shapes of some molecules
Sigma Bonds

Sigma bond (σ-bond) is formed when atoms overlap directly along the inter-nuclear axis (i.e. headon overlap). A σ-bond forms from three possibilities:
+
The overlapping of
𝑠 −orbitals, e.g. in 𝐻2
The overlapping of an
𝑠 −orbital and a 𝑝 −orbital,
e.g. in 𝐻𝐵𝑟, 𝐻𝐶𝑙
+
+
The head-on overlapping of
𝑝 −orbitals, e.g. in 𝐶𝑙2 , 𝐼2
Fig.3.1 Formation of sigma bonds


All single bonds between any two atoms are σ-bonds.
It is not possible to form more than one σ-bond between any two atoms, since doing so would force
too many electrons into a small space and generate repulsion.
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Pi-Bonds

A π-bond is one formed from overlap of atomic orbitals above and below the inter-nuclear axis (i.e.
sideways overlap of p-orbitals).
Fig.3.2 Formation of pi bonds




All double bonds consist of one σ-bond and one π-bond.
All triple bonds consist of a σ-bond and two π-bonds. If the first π -bond results from overlap above
and below the inter-nuclear axis, the second results from overlap behind and in front of the internuclear axis.
Note that π -bonds can only be formed by overlap of p-orbitals, since s-orbitals do not have the
correct geometry.
A pi-bond has two regions, one above and the
other below the plane of the molecule.
Fig.3.3 Sigma and pi bonds in alkenes
Comparing Sigma and Pi-bonds
Sigma (σ-bond)
Formed by the axial (along the axis) overlap of
orbitals
Formed by overlap of 𝒔 − 𝒔, 𝒔 − 𝒑 or 𝒑 − 𝒑
orbitals (head-on)
Bond is stronger because the overlapping can
take place to a greater extent
There can be a free rotation of atoms around a
sigma bond
The molecular orbital consists of a single charged
cloud and is symmetrical about the inter-nuclear
axis
May exist between the two atoms either alone or
along with a pi-bond
Pi (π-bond)
Formed by the sideway overlap of orbitals
Formed by overlap of 𝑝 −orbitals (side-ways) only
Bond is weaker because the overlapping is only to a
smaller extent
Rotation about a pi-bond is not possible as it
involves breaking the pi-bond
The molecular orbital consist of two regions, one
above and one below the plane of atoms
Only exists along with a sigma bond
Table 3.4 Comparing sigma and pi-bonds
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3.3
Bond Properties and Intermolecular Forces
Electronegativity




Electronegativity is the relative ability of an atom to attract electrons in a covalent bond.
Electronegativity increases across a period as the nuclear charge on the atoms increases but the
shielding stays the same, so that the nucleus is more able to attract to electrons.
Electronegativity decreases down a group as the number of shells increases, so shielding increases
and the nucleus is less able to attract to electrons.
Fluorine is the most electronegative element. Noble gases cannot be ascribed an electronegativity
value since they do not form bonds.
Polar Bonds



Consider a covalent bond between two atoms 𝐴 and 𝐵. If both atoms have a similar electronegativity,
their atoms attract to the bonding electrons with equal strength and the bonding electrons will
remain mid-way between the two. The bond between 𝐴 and 𝐵 is said to be non-polar. 𝐻2 or 𝐶𝑙2 are
examples of molecules with non-polar bonds.
If element 𝐵 is slightly more electronegative than element 𝐴, 𝐵 will attract the electron pair closer to
itself. This makes 𝐵 partially negative, and 𝐴 becomes slightly positive. The bond is said to be polar.
Examples of polar bonds are 𝐻 − 𝐶𝑙 and 𝑂 − 𝐻 bonds.
The greater the difference in electronegativity, the greater the polarity of the bond, e.g., the 𝐻 − 𝐹
bond is more polar than the 𝐻 − 𝐵𝑟 bond because fluorine is more electronegative than bromine.
Non-polar bond, e.g. in 𝐻2
H
x
o
Polar bond, e.g. the 𝑂 − 𝐻 in 𝐻2 𝑂
- x
H
O o
+
H
Polar Molecules



Non-polar molecules are those in which there are no polar bonds or in which the dipoles resulting
from the polar bonds all cancel each other out.
Polar molecules are those in which there are polar bonds and in which the dipoles resulting from the
polar bonds do not cancel out.
The molecules 𝐵𝐹3 and 𝐶𝑂2 are non-polar because the dipoles cancel out:
F
F
B

C
O
F
O
The following molecules are polar because the dipoles do not cancel out:
O

- S
O

+

N
H
H
H
+
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EXPERIMENT
Aim: To determine whether a liquid
substance is polar
Method:



Example 1:
𝐶𝐶𝑙3 𝐻 is polar. The hydrogen at the top of the molecule is less
electronegative than carbon and so is slightly positive. This means
that the molecule has a slightly positive "top" and a slightly negative
"bottom", and so is overall a polar molecule.
Place the liquid in a burette;
Allow a narrow stream to run out;
Place a charged rod next to the
flow;
Observation:
𝐶𝐻𝐶𝑙3 is polar
Example 2:
𝐶𝐶𝑙4 , is non-polar. Each bond is polar but the molecule as a whole is
non-polar. The whole of the outside of the molecule is somewhat
negative, but there is no overall separation of charge from top to
bottom or from left to right.


Polar substance will be attracted.
Non-polar substance will not be
attracted.
𝐶𝐶𝑙4 is non-polar
Intermolecular Forces


a)



Intermolecular forces are attractions between one molecule and a neighbouring molecule. There are
three main types of intermolecular forces:
a) Van der Waals’ forces;
b) Permanent dipole – permanent dipole attractions;
c) Hydrogen bonds.
It is these forces which must be overcome when a simple molecular substance such as water is being
melted or boiled.
Van der Waal's Forces (Induced or Temporary Dipole Attractions)
The electrons in a molecule are in a state of constant motion. At any given time the distribution of
electrons will not be exactly symmetrical - there is likely to be a slight surplus of electrons on one end
of the atom. The other end will be temporarily short of electrons and so becomes slightly positively
charged.
This temporary separation of positive and negative charges is called a temporary or instantaneous
dipole. It lasts for a very short time as the electrons are constantly moving. If another molecule
approaches, its electrons will be attracted by the slightly positive end of the already polarized molecule.
This sets up an induced dipole in the approaching molecule.
This random movement of electrons can occur over huge numbers of molecules such that a whole
lattice of molecules could be held together by these induced forces, see Fig.3.3:
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Fig.3.3 Induced dipoles (van der Waals’ forces)

This temporary, instantaneous attraction between the dipoles is called a Van der Waal's force. The
strength of the Van der Waal's forces depends on two factors:
1. Molecular size (the number of electrons in the molecule)
With a greater number of electrons in a molecule, there is a greater magnitude of the temporary
dipoles. Thus the boiling points of the noble gases increase with increasing number of electrons down
the group:
substance
number of electrons
boiling point/°C
𝑯𝒆
2
−269
𝑵𝒆
10
−250
𝑨𝒓
18
−186
𝑲𝒓
𝑿𝒆
36
54
−152
−108
Table 3.5 Boiling points of the noble gases
The boiling points of the straight-chain alkanes increase with increasing number of electrons per
molecule:
𝑪𝑯𝟒
𝑪𝟐 𝑯𝟔
𝑪𝟑 𝑯𝟖
𝑪𝟒 𝑯𝟏𝟎
substance
number of electrons
10
18
26
34
boiling point/°C
−164
−88
−42
0
Table 3.6 Boiling points of some alkanes
2. Molecular structure (surface area of contact between molecules)
Long thin molecules mean there is a greater surface area for contact, making the attractions more
effective. Therefore, branched molecules have lower melting and boiling points than straight-chain
molecules of the same number of electrons. Consider butane and 2-methylpropane, each of which has
34 electrons:
Boiling point:
−0.5°C
Boiling point:
−11.7°C
b) Permanent Dipole Attractions



Polar molecules have permanent dipoles.
permanent
dipole
Permanent dipoles attract the molecules more strongly than van der
Waal’s forces. Molecules which have permanent dipoles thus have
boiling points higher than molecules which only have temporary dipoles.
For
example,
ethane,
𝐶2 𝐻6
(boiling
point
−88°C)
and
fluoromethane,
𝐶𝐻3 𝐹 (boiling point −78°C) have same numbers of electron but fluoromethane has a higher boiling
point due to the large permanent dipole it has due to the high electronegativity of fluorine.
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c)



Hydrogen Bonds
Hydrogen bonds form between a hydrogen atom in one molecule, and a highly electronegative atom
in another molecule. The highly electronegative atom could be oxygen, nitrogen, or fluorine.
Hydrogen bonding is just a stronger form of permanent dipole attractions. Examples of substances
containing hydrogen bonds are 𝐻𝐹 , 𝐻2 𝑂, 𝑁𝐻3 , alcohols, carboxylic acids, primary amines, etc.
Conditions for hydrogen bonding to exist:
1. The hydrogen must be attached directly to one of the most electronegative elements (𝐹, 𝑂, and
𝑁). This gives the 𝐻 atom a high partial positive charge.
2. Each of the atoms to which the hydrogen is attached must have at least one "active" lone pair.
Ethanoic acid (CH3CO2H, Mr of 60.0) has an
apparent Mr of 120. Because of hydrogen
bonding, the molecules pair up to form
dimers of Mr = 120.
Effects of Hydrogen Bonds
1. Unusually high boiling points
Substances containing hydrogen bonds have much higher boiling points than would be predicted from
Van der Waal's forces alone.
Boling Points of Hydrides of Group V, VI and VII
150
Boiling Point /˚C
100
50
H2O
HF
H2Te
0
-50
-100
NH3
H2Se
H2S
HBr
HCl
PH3
HI
SbH3
AsH3
Table 3.7 Boiling points of hydrides of Groups V, VI and VII

In each case the hydride of 𝑁, 𝑂 or 𝐹 shows a boiling point which is abnormally high. They have very
strong intermolecular hydrogen bonds between their molecules despite the fact that their Van der
Waal's forces are weaker than in the other hydrides.
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2. The low density of ice



In ice, each water molecule is hydrogen-bonded to 4 others in a tetrahedral formation, giving ice a
“diamond-like” structure.
The result is that ice has a very open hexagonal structure with large spaces within the crystal. This
accounts for its large volume, and hence its low density.
Ice floats on water because it is less dense. Also, this explains why water freezes from the top surface.
Fig.3.4 Hydrogen bonding in ice


In ice, each water molecule forms two hydrogen bonds with other neighbouring water molecules
such that each oxygen atom is surrounded by four hydrogen bonded atoms. There are a total of 4
bond pairs and no lone pairs around 𝑂, hence shape is tetrahedral about 𝑂 and bond angle is 109.5°.
Water expands when it freezes – most substances actually contract during freezing. When ice melts,
the structure collapses slightly and the molecules come closer together.
3. High surface tension of water

Because hydrogen bonds in water are fairly strong, they give it a skinny surface. The ‘skin’ is strong
enough such that small insects can crawl on the water surface without sinking!
4. The helical nature of DNA


Molecules of DNA contain 𝑁 − 𝐻 bonds and so hydrogen bonding is possible. The long chains also
contain 𝐶 = 𝑂 bonds and the 𝐻 atoms can form a hydrogen bond with this electronegative 𝑂 atom.
This results in the molecule spiralling, as the 𝐶 = 𝑂 bonds and the 𝑁 − 𝐻 bonds approach each other.
This is an example of an intramolecular hydrogen bond, where the attraction is between a hydrogen
atom and an electronegative atom on the same molecule. With intermolecular hydrogen bonding, the
attraction is between a hydrogen atom and an electronegative atom on a neighbouring molecule.
3.4



Metallic Bonding
Metallic bonding is the electrostatic attraction between metal cations and a sea of delocalised electrons.
The cations are arranged to form a lattice, with the electrons free to move between them.
In metals, the outermost shell electrons delocalise to form a “cloud” or “sea” of electrons. The
electrostatic forces between the electron cloud and the metal ions form the metallic bonds
The delocalised electrons in a metal are free to move throughout the lattice, so that metals can conduct
electricity in solid or molten state.
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
Strength of metallic bonding depends on:
1. The number of electrons in outer shell: - for metallic atoms with same
number of shells, the one with more electrons in outer shell has higher melting
point. The more the electrons delocalised per atom, the stronger the
electrostatic forces that form.
The melting points increase from Na to Al as the strength of metallic bonding
increases:
𝑁𝑎
𝑀𝑔
𝐴𝑙
metal
electrons in outer shell
1
2
3
Table 3.8
melting point /oC
98
649
660
2. The radius of the metal atom: - a larger atom/ion makes weaker electrostatic attractions with its
delocalised electron cloud, thus lowering the melting point.
𝐿𝑖
𝑁𝑎
𝐾
metal
atomic radius /nm
0.152
0.186
0.227
Table 3.9
melting point /oC
181
98
63
3.5 Bonding and Physical Properties
substance
nature of bonding

Ionic
e.g. 𝑵𝒂𝟐 𝑶, 𝑨𝒍𝑭𝟑

Metallic
e.g. 𝑴𝒈, 𝒁𝒏, 𝑨𝒈

Giant Molecular
e.g. 𝑪 (graphite or
diamond), 𝑺𝒊𝑶𝟐


Simple Molecular
e.g. 𝑯𝟐 𝑶, 𝑰𝟐 , 𝑺𝟖 ,
𝑶𝟐

Electrostatic
attraction
between oppositely charged
ions.
Electrostatic
attraction
between
cations
and
delocalised electrons.
Lattice of atoms linked by
covalent bonds in three
dimensions.
Covalent bonds are pairs of
electrons shared between
two atoms.
Discrete molecules. Atoms
in molecule linked by
covalent bonds. (σ or π)
Weak intermolecular forces
between molecules.
physical properties










High melting and boiling points. They are
hard, strong and brittle.
Good conductors in molten state; poor
conductors in solid state.
High melting and boiling points. They are
strong and malleable.
Good electrical conductors in solid and
liquid states.
Very high melting and boiling points.
Poor conductors in solid and liquid states
(graphite conducts along the layers).
They are hard, strong and brittle.
Low melting and boiling points.
Poor electrical conductors in solid and liquid
states.
Soft and weak
Table 3.9 Summary properties of different substances
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Examination Practice 3
1. At room temperature, both sodium metal and sodium chloride are crystalline solids which contain ions.
(a) Copy the diagrams for sodium metal and sodium chloride below, and mark the charge for each ion.
Sodium metal
Sodium chloride
(2)
(b) (i) Explain how the ions are held together in solid sodium metal.
(ii) Explain how the ions are held together in solid sodium chloride.
(iii) The melting point of sodium chloride is much higher than that of sodium metal. What can be
deduced from this information?
(3)
(c) Compare the electrical conductivity of solid sodium metal with that of solid sodium chloride.
Explain your answer.
(3)
(d) Explain why sodium metal is malleable (can be hammered into shape).
(1)
(Total 9 marks)
2. The equation below shows the reaction between boron trifluoride and a fluoride ion.
𝑩𝑭𝟑 + 𝑭−
𝑩𝑭−
𝟒
(i) Draw diagrams to show the shape of the BF3 molecule and the shape of the 𝐵𝐹4− ion. In each case,
3.
4.
5.
6.
name the shape. Account for the shape of the 𝐵𝐹4− ion and state the bond angle present.
(ii) In terms of the electrons involved, explain how the bond between the 𝐵𝐹3 molecule and the 𝐹 −
ion is formed. Name the type of bond formed in this reaction.
(Total 9 marks)
Draw the shape of a molecule of 𝐵𝑒𝐶𝑙2 and the shape of a molecule of 𝐶𝑙2𝑂. Show any lone pairs of
electrons on the central atom. Name the shape of each molecule.
(4)
(Total 4 marks)
Ammonia, 𝑁𝐻3, reacts with sodium to form sodium amide, 𝑁𝑎𝑁𝐻2, and hydrogen.
(a) Draw the shape of an ammonia molecule and that of an amide ion, 𝑁𝐻2− .
In each case show any lone pairs of electrons.
(b) State the bond angle found in an ammonia molecule.
(c) Explain why the bond angle in an amide ion is smaller than that in an ammonia molecule.
(5)
(Total 5 marks)
(a) Describe the bonding that is present in metals.
(3)
(b) Explain how the bonding and structure lead to the typical metallic properties of electrical
conductivity and malleability.
(4)
(c) Suggest a reason why aluminium is a better conductor of electricity than magnesium.
(2)
(Total 9 marks)
The table below shows the electronegativity values of some elements.
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Fluorine
Chlorine
Bromine
Iodine
Carbon
Hydrogen
4.0
3.0
2.8
2.5
2.5
2.1
Electronegativity
(a) Define the term electronegativity.
(2)
(b) The table below shows the boiling points of fluorine, fluoromethane (𝐶𝐻3𝐹) and hydrogen fluoride.
F–F
H–F
F
C
H
Boiling point/K
85
H
H
194
293
(i) Name the strongest type of intermolecular force present in liquid 𝐹2, liquid 𝐶𝐻3𝐹 and liquid
𝐻𝐹.
(ii) Explain how the strongest type of intermolecular force in liquid 𝐻𝐹 arises.
(6)
(c) The table below shows the boiling points of some other hydrogen halides.
Boiling point / K
𝐻𝐶𝑙
𝐻𝐵𝑟
𝐻𝐼
188
206
238
(i) Explain the trend in the boiling points of the hydrogen halides from 𝐻𝐶𝑙 to 𝐻𝐼.
(ii) Give one reason why the boiling point of 𝐻𝐹 is higher than that of all the other hydrogen
halides.
(3)
(Total 11 marks)
7. (a) Methanol has the structure shown below:
H
H
C
O
H
H
Explain why the 𝑂– 𝐻 bond in a methanol molecule is polar.
(2)
(b) The boiling point of methanol is +65 °C; the boiling point of oxygen is –183 °C. Methanol and
oxygen each have an 𝑀𝑟 value of 32. Explain, in terms of the intermolecular forces present in each
case, why the boiling point of methanol is much higher than that of oxygen.
(3)
(Total 5 marks)
8. (a) The diagram below shows the melting points of some of the elements in Period 3.
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2000
1600
Melting
point/K
1200
800
400
0
Na Mg
9.
10.
11.
12.
Al
Si
P
S
Cl
Ar
(i) On the diagram, use crosses to mark the approximate positions of the melting points for the
elements silicon, chlorine and argon. Complete the diagram by joining the crosses.
(ii) By referring to its structure and bonding, explain your choice of position for the melting point
of silicon.
(iii) Explain why the melting point of sulphur, 𝑆8, is higher than that of phosphorus, 𝑃4.
(8)
(b) State and explain the trend in melting point of the Group II elements 𝐶𝑎– 𝐵𝑎.
(3)
(Total 11 marks)
State and explain the trend in the melting points of the Period 3 metals 𝑁𝑎, 𝑀𝑔 and 𝐴𝑙.
(3)
(Total 3 marks)
(a) (i) Describe the bonding in a metal.
(ii) Explain why magnesium has a higher melting point than sodium.
(4)
(b) Why do diamond and graphite both have high melting points?
(3)
(c) Why is graphite a good conductor of electricity?
(1)
(d) Why is graphite soft?
(2)
(Total 10 marks)
Sodium sulphide, 𝑁𝑎2𝑆, is a high melting point solid which conducts electricity when molten. Carbon
disulphide, 𝐶𝑆2, is a liquid which does not conduct electricity.
(a) Deduce the type of bonding present in 𝑁𝑎2𝑆 and that present in 𝐶𝑆2.
(b) By reference to all the atoms involved explain, in terms of electrons, how 𝑁𝑎2𝑆 is formed from its
atoms.
(c) Draw a diagram, including all the outer electrons, to represent the bonding present in 𝐶𝑆2.
(6)
(Total 6 narks)
(a) The diagram below represents a part of the structure of sodium chloride. The ionic charge is shown
on the centre of only one of the ions.
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+
(i) Copy the diagram, and mark the charges on the four negative ions.
(ii) What change occurs to the motion of the ions in sodium chloride when it is heated from room
temperature to a temperature below its melting point?
(2)
(b) Sodium chloride can be formed by reacting sodium with chlorine.
(ii) A chloride ion has one more electron than a chlorine atom. In the formation of sodium chloride,
from where does this electron come?
(ii) What property of the atoms joined by a covalent bond causes the bond to be polar?
(2)
(Total 4 marks)
13. Phosphorus and nitrogen are in Group V of the Periodic Table and both elements form hydrides.
Phosphine, 𝑃𝐻3, reacts to form phosphonium ions, 𝑃𝐻4+ , in a similar way to that by which ammonia,
𝑁𝐻3, forms ammonium ions, 𝑁𝐻4+
(a) Give the name of the type of bond formed when phosphine reacts with an 𝐻 + ion. Explain
how this bond is formed.
(3)
(b) Draw the shapes, including any lone pairs of electrons, of a phosphine molecule and of a
phosphonium ion.
Give the name of the shape of the phosphine molecule and state the bond angle found in the
phosphonium ion.
(4)
(Total 7 marks)
14. (a) State the meaning of the term electronegativity.
(2)
(b) State and explain the trend in electronegativity values across Period 3 from sodium to chlorine.
(3)
(Total 5 marks)
15. (a) The shape of the molecule 𝐵𝐶𝑙3 and that of the unstable molecule 𝐶𝐶𝑙2 are shown below.
Cl
B
Cl
(i)
(ii)
C
Cl
Cl
Cl
Why is each bond angle exactly 120° in 𝐵𝐶𝑙3?
Predict the bond angle in 𝐶𝐶𝑙2 and explain why this angle is different from that in 𝐵𝐶𝑙3.
(5)
(b) Give the name which describes the shape of molecules having bond angles of 109° 8'.
Give an example of one such molecule.
(2)
(c) The shape of the 𝑋𝑒𝐹4 molecule is shown below.
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F
F
Xe
F
F
(i) State the bond angle in 𝑋𝑒𝐹4
(ii) Suggest why the lone pairs of electrons are opposite each other in this molecule.
(iii) Name the shape of this molecule, given that the shape describes the positions of the 𝑋𝑒 and 𝐹
atoms only.
(4)
(d) Draw a sketch of the 𝑁𝐹3 molecule. Indicate in your sketch any lone pairs of electrons on nitrogen.
(2)
(Total 13 marks)
16. (a) Describe the motion of the particles in solid iodine and in iodine vapour.
(3)
(b) Explain why solid iodine vaporises when warmed gently.
(2)
(c) Silver and sodium chloride melt at similar temperatures. Give two physical properties of silver which
are different from those of sodium chloride and, in each case, give one reason why the property of
silver is different from that of sodium chloride.
(4)
2−
(d) Draw the shapes of 𝐵𝑒𝐶𝑙2, 𝑁𝐶𝑙3 and 𝐵𝑒𝐶𝑙4 . In each case, show any lone-pair electrons on the
central atom and state the value of the bond angle.
(6)
(Total 15 marks)
17. Silicon dioxide has a macromolecular structure. Draw a diagram to show the arrangement of atoms
around a silicon atom in silicon dioxide. Give the name of the shape of this arrangement of atoms and
state the bond angle.
(3)
(Total 3 marks)
18. (a) When considering electron pair repulsions in molecules, why does a lone pair of electrons repel
more strongly than a bonding pair?
(1)
(b) The diagram below shows a hydrogen peroxide molecule.
H
O
O
H
(i) Copy the diagram, and draw the lone pairs, in appropriate positions, on the oxygen atoms.
(ii) Indicate, on the diagram, the magnitude of one of the bond angles.
(iii) Name the strongest type of intermolecular force which exists between molecules of hydrogen
peroxide in the pure liquid.
(4)
(c) Draw a diagram to illustrate the shape of a molecule of 𝑆𝐹4 and predict the bond angle(s).
(4)
(d) Name two types of intermolecular force which exist between molecules in liquid 𝑆𝐹4.
(2)
(Total 11 marks)
19. (a) Name the type of force that holds the particles together in an ionic crystal.
(1)
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(b) What is a covalent bond?
(1)
(c) State how a co-ordinate bond is formed.
(2)
(d) Describe the bonding in a metal.
(2)
(e) A molecule of hydrogen chloride has a dipole and molecules of hydrogen chloride attract each
other by permanent dipole-dipole forces. Molecules of chlorine are non-polar.
(i) What is a permanent dipole?
(ii) Explain why a molecule of hydrogen chloride is polar.
(iii) Name the type of force which exists between molecules of chlorine.
(5)
(f) Show, by means of a diagram, how two molecules of hydrogen fluoride are attracted to each other
by hydrogen bonding; include all lone-pair electrons and partial charges in your diagram.
(3)
(g) Why is there no hydrogen bonding between molecules of hydrogen bromide?
(1)
(Total 15 marks)
20. The table below gives the boiling points, 𝑇𝑏 , of some hydrogen halides.
Hydrogen halide
𝑇𝑏 /𝐾
HF
293
HCl
188
HBr
206
HI
238
(a) By referring to the types of intermolecular force involved, explain why energy must be supplied in
order to boil liquid hydrogen chloride.
(3)
(b) Explain why the boiling point of hydrogen bromide lies between those of hydrogen chloride and
hydrogen iodide.
(2)
(c) Explain why the boiling point of hydrogen fluoride is higher than that of hydrogen chloride.
(2)
(d) Draw a sketch to illustrate how two molecules of hydrogen fluoride interact in liquid hydrogen
fluoride.
(2)
(Total 9 marks)
21. Sulphur will combine separately with carbon, hydrogen and sodium to form carbon disulphide (𝐶𝑆2),
hydrogen sulphide (𝐻2𝑆) and sodium sulphide (𝑁𝑎2𝑆) respectively. The bonding in these compounds
is similar to that in 𝐶𝑂2, 𝐻2𝑂 and 𝑁𝑎2𝑂.
(a) Complete the table in Figure 2 by classifying the compounds as either ionic or covalent.
Figure 2
Carbon disulphide
Hydrogen sulphide
Sodium sulphide
Melting point/K
162
187
1450
Ionic or covalent
(3)
(b) One of the compounds in Figure 2 shows high electrical conductivity under appropriate conditions.
Identify the compound, by name or formula, and state one condition under which it shows high
electrical conductivity.
(2)
(Total 5 marks)
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22. (a) State what is meant by the term polar bond.
(1)
(b) Sulphuric acid is a liquid that can be represented by the formula drawn below.
O
O
H
O
H
S
O
Given that the electronegativity values for hydrogen, sulphur and oxygen are 2.1, 2.5 and 3.5
respectively, clearly indicate the polarity of each bond present in the formula given.
(2)
(c) Suggest the strongest type of intermolecular force present in pure sulphuric acid.
Briefly explain how this type of intermolecular force arises.
(2)
(Total 5 marks)
23. (a) Sketch the shapes of each of the following molecules, showing any lone pairs of electrons. In each
case, state the bond angle(s) present in the molecule and name the shape.
Molecule
Sketch of shape
Bond angle(s)
Name of shape
𝐵𝐹3
𝑁𝐹3
𝐶𝑙𝐹3
(9)
(b) State the types of intermolecular force which exist, in the liquid state, between pairs of
𝐵𝐹3 molecules and between pairs of 𝑁𝐹3 molecules.
(3)
(c) Name the type of bond which you would expect to be formed between a molecule of 𝐵𝐹3 and a
molecule of 𝑁𝐹3. Explain how this bond is able to form.
(3)
(Total 15 marks)
24. Sketch a diagram to show the shape of a molecule of 𝑁𝐻3 and indicate on your
diagram how this molecule is attracted to another 𝑁𝐻3 molecule in liquid ammonia.
(3)
(Total 3 marks)
25. (a) Define the term electronegativity.
(2)
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(b) State and explain the trend in electronegativity of the elements across Period 3 from sodium to
chlorine.
(3)
(c) State the bond type in sodium oxide and the bond type in sulphur dioxide. In each case, explain
the link between the bond type and the electronegativity of the elements involved.
(4)
(Total 9 marks)
26. (a) Co-ordinate bonding can be described as dative covalency. In this context, what is the meaning of
each of the terms covalency and dative?
(2)
(b) Write an equation for a reaction in which a co-ordinate bond is formed.
(2)
(c) Why is sodium chloride ionic rather than covalent?
(2)
(d) Why is aluminium chloride covalent rather than ionic?
(2)
(e) Why is molten sodium chloride a good conductor of electricity?
(1)
(f) Explain, in terms of covalent bonding, why the element iodine exists as simple molecules whereas
the element carbon does not.
(3)
(Total 12 marks)
27. (a) Describe the nature and strength of the bonding in solid calcium oxide.
(3)
(b) Use the kinetic theory to describe the changes that take place as calcium oxide is heated from 25°C
to a temperature above its melting point.
(3)
(c) State two properties of calcium oxide that depend on its bonding.
(2)
(Total 8 marks)
28. What is a covalent bond?
(1)
(Total 1 mark)
29. (a) Figure 1 shows some data concerned with halogens.
Element
Fluorine
Chlorine
Bromine
Iodine
Electronegativity
4.0
3.0
2.8
2.5
Figure 1
Boiling point of hydride / 𝑲
293
188
206
238
(i) Define the term electronegativity.
(ii) Explain the trend in boiling points from hydrogen chloride to hydrogen iodide.
(iii) Explain why hydrogen fluoride does not fit this trend.
(2 + 2 + 2)
(b) The oxygen atoms in the sulphate ion surround the sulphur in a regular tetrahedral shape.
(i) Write the formula of the ion.
(ii) State the O–S–O. bond angle.
(1 + 1)
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(Total 8 marks)
30. (a) State the type of bonding in a crystal of potassium bromide.
(1)
(b) Sketch a diagram to show the shape of a 𝐵𝑟𝐹3 molecule. Show on your sketch any lone pairs of
electrons in the outermost shell of bromine and name the shape.
(3)
(Total 4 marks)
31. (a) (i) State one feature which molecules must have in order for hydrogen bonding to occur between
them.
(ii) Give the name of the type of intermolecular bonding present in hydrogen sulphide, 𝐻2𝑆, and
explain why hydrogen bonding does not occur.
(iii) Account for the much lower boiling point of hydrogen sulphide (–61 °C) compared with that
of water (100 °C).
(1 + 2 + 2)
(b) Protein molecules are composed of sequences of amino acid molecules that have joined together,
with the elimination of water, to form long chains. Part of a protein chain is represented by the
displayed formula given below.
R
O
C
C
H
R
O
N
C
C
H
H
R
O
N
C
C
H
H
R
N
C
H
H
Explain the formation of hydrogen bonding between protein molecules.
(4)
(Total 9 marks)
32. (a) Describe the bonding found in metals.
(3)
(b) Use data from table above and your knowledge of the bonding in these metals to explain why the
melting point of magnesium is higher than that of sodium.
(3)
(c) State and explain the similarities and differences in electrical conductivity of sodium, graphite and
diamond.
(4)
(Total 10 marks)
33. The table below contains electronegativity values for the Period 3 elements, except chlorine.
Element
Na
Mg
Al
Si
P
S
Cl
Ar
Electronegativity
0.9
1.2
1.5
1.8
2.1
2.5
(a) How can electronegativity values be used to predict whether a given chloride is likely to be ionic
or covalent?
(2)
(b) State the type of bonding in sodium oxide.
(1)
(Total 3 marks)
34. The diagram below shows how a water molecule interacts with a hydrogen fluoride molecule.
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H
O
+
–
H
F
H
(a) What is the value of the bond angle in a single molecule of water?
(1)
(b) Explain your answer to part (a) by using the concept of electron pair repulsion.
35.
36.
37.
38.
39.
(4)
(c) Name the type of interaction between a water molecule and a hydrogen fluoride molecule shown
in the diagram above.
(1)
(d) Explain the origin of the δ+ charge shown on the hydrogen atom in the diagram.
(2)
(e) When water interacts with hydrogen fluoride, the value of the bond angle in water changes slightly.
Predict how the angle is different from that in a single molecule of water and explain your answer.
(2)
(Total 10 marks)
(a) State which one of the elements neon, sodium, magnesium, aluminium and silicon has the lowest
melting point and explain your answer in terms of the structure and bonding present in that
element.
(b) State which one of the elements neon, sodium, magnesium, aluminium and silicon has the highest
melting point and explain your answer in terms of the structure and bonding present in that
element.
(3 + 3)
(Total 6 marks)
Diamond and graphite are both forms of carbon.
Diamond is able to scratch almost all other substances, whereas graphite may be used as a lubricant.
Diamond and graphite both have high melting points.
Explain each of these properties of diamond and graphite in terms of structure and bonding. Give one
other difference in the properties of diamond and graphite.
(Total 9 marks)
Iodine and diamond are both crystalline solids at room temperature. Identify one similarity in the
bonding,
and
one
difference
in
the
structures,
of
these
two
solids.
Explain why these two solids have very different melting points.
(Total 6 marks)
Phosphorus exists in several different forms, two of which are white phosphorus and red phosphorus.
White phosphorus consists of 𝑃4 molecules, and melts at 44°C. Red phosphorus is macromolecular,
and has a melting point above 550°C.
Explain what is meant by the term macromolecular. By considering the structure and bonding present
in these two forms of phosphorus, explain why their melting points are so different.
(Total 5 marks)
(a) Predict the shapes of the 𝑆𝐹6 molecule and the 𝐴𝑙𝐶𝑙4− ion. Draw diagrams of these species to show
their three-dimensional shapes. Name the shapes and suggest values for the bond angles. Explain your
reasoning.
(8)
(b) Perfume is a mixture of fragrant compounds dissolved in a volatile solvent.
When applied to the skin the solvent evaporates, causing the skin to cool for a short time. After a
while, the fragrance may be detected some distance away. Explain these observations.
(4)
(Total 12 marks)
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40. (a) Iodine and graphite crystals both contain covalent bonds and yet the physical properties of their
crystals are very different.
For iodine and graphite, state and explain the differences in their melting points and in their
electrical conductivities.
(9)
(b) Draw the shape of the 𝐵𝑒𝐶𝑙2 molecule and explain why it has this shape.
(2)
(Total 11marks)
41. (a) The table below gives the melting point for each of the Period 3 elements 𝑁𝑎 – 𝐴𝑟.
42.
43.
44.
45.
46.
Element
Na
Mg
Al
Si
P
S
Cl
Ar
Melting point / K
371
923
933
1680
317
392
172
84
In terms of structure and bonding, explain why silicon has a high melting point, and why the melting
point of sulphur is higher than that of phosphorus.
(7)
(b) Draw a diagram to show the structure of sodium chloride. Explain, in terms of bonding, why sodium
chloride has a high melting point.
(4)
(Total 11 marks)
Explain the meaning of the term periodicity as applied to the properties of rows of elements in the
Periodic Table. Describe and explain the trends in atomic radius, in electronegativity and in conductivity
for the elements sodium to argon.
(13)
(Total 13 marks)
(a) Describe the structure of, and bonding in, three different types of crystal. Illustrate your answer
with a specific example of each type of crystal and sketch labelled diagrams of the structures. In
each case, explain how the ability to conduct electricity is influenced by the type of bonding.
(18)
(b) Explain how the concept of bonding and lone (non-bonding) pairs of electrons can be used to
predict the shape of, and bond angles in, a molecule of sulphur tetrafluoride, 𝑆𝐹4. Illustrate your
answer with a sketch of the structure.
(8)
(Total 26 marks)
Sketch a graph to show how the melting points of the elements vary across Period 3 from sodium to
argon. Account for the shape of the graph in terms of the structure of, and the bonding in, the elements.
(Total 21 marks)
(a) With the aid of diagrams, describe the structure of, and bonding in, crystals of sodium chloride,
graphite and magnesium. In each case, explain how the melting point and the ability to conduct
electricity of these substances can be understood by a consideration of the structure and bonding
involved.
(23)
(b) Explain how the electron-pair repulsion theory can be used to predict the shapes of the
molecules 𝐻2𝑂 and 𝑃𝐹5. Illustrate your answer with diagrams of the molecules on which the bond
angles are shown.
(7)
(Total 30 marks)
In this question, one mark is available for the quality of spelling, punctuation and grammar.
Many physical properties can be explained in terms of bonding and structure. The table below shows
the structures and some properties of sodium chloride and graphite in the solid state.
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substance
structure
sodium chloride
graphite
electrical
conductivity of solid
melting and boiling
point
solubility in water
poor
good
high
high
good
insoluble
Explain these properties in terms of bonding and structure.
[7]
Quality of Written Communication [1]
[Total 8 marks]
47. Sodium reacts with excess oxygen to form sodium peroxide, 𝑁𝑎2𝑂2.
𝑁𝑎2𝑂2 is used in laundry bleaches. When added to water a reaction takes place forming an alkaline
solution and hydrogen peroxide, 𝐻2𝑂2.
(i) Construct a balanced equation for the formation of sodium peroxide from sodium.
(ii) Construct a balanced equation for the reaction of sodium peroxide with water.
(iii) Draw a ‘dot-and-cross’ diagram for a molecule of 𝐻2𝑂2. Show outer electrons only.
[1 + 1 + 2]
[Total 4 marks]
48. In water treatment plants, care must be taken as chlorine can react with nitrogen compounds to form
the highly explosive compound, nitrogen trichloride, 𝑁𝐶𝑙3. Molecules of 𝑁𝐶𝑙3 have a bond angle of
107°.
(i) Name the shape of an 𝑁𝐶𝑙3 molecule.
(ii) Explain why a molecule of 𝑁𝐶𝑙3 has this shape and a bond angle of 107°.
[1 + 3]
[Total 4 marks]
49. Chemists have developed models for bonding and structure which are used to explain different
properties. Ammonia, 𝑁𝐻3, is a covalent compound.
(i) Explain what is meant by a covalent bond.
(ii) Draw a ‘dot-and-cross’ diagram to show the bonding in 𝑁𝐻3 . Show outer electrons only.
(iii) Name the shape of the ammonia molecule. Explain, using your ‘dot-and-cross’ diagram, why
ammonia has this shape and has a bond angle of 107°.
[1 + 1 + 3]
[Total 5 marks]
50. The shape of a water molecule is different from the shape of a carbon dioxide molecule.
(i) Draw the shapes of these molecules and state the bond angles.
(ii) Explain why a water molecule has a different shape from a carbon dioxide molecule.
[4 + 2]
[Total 6 marks]
51. Ammonia reacts with hydrogen chloride, 𝐻𝐶𝑙, to form ammonium chloride, 𝑁𝐻4𝐶𝑙.
𝑁𝐻4𝐶𝑙 is an ionic compound containing 𝑁𝐻4+ and 𝐶𝑙 − ions.
(i) Write the electron configuration of the 𝐶𝑙 − ion.
(ii) Draw a ‘dot-and-cross’ diagram to show the bonding in 𝑁𝐻4+ . Show outer electrons only.
(iii) State the shape of, and bond angle in, an 𝑁𝐻4+ ion.
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(iv) A student investigated the conductivity of ammonium chloride. She noticed that when the
ammonium chloride was solid it did not conduct electricity. However, when ammonium chloride
was dissolved in water, the resulting solution did conduct electricity.
Explain these observations.
[1 + 1 + 2 +2]
[Total 6 marks]
52. Compounds with covalent bonding often have polar bonds. Polarity can be explained in terms of
electronegativity.
(i) Explain the term electronegativity.
(ii) Use a suitable example to show how the presence of a polar bond can be explained in terms of
electronegativity. You may find it useful to draw a diagram in your answer.
[2 + 2]
[Total 4 marks]
53. Liquid ammonia, 𝑁𝐻3, and water, 𝐻2𝑂, both show hydrogen bonding.
(i) Draw a labelled diagram to show hydrogen bonding between two molecules of liquid ammonia.
(ii) Water has several anomalous properties as a result of its hydrogen bonding. Describe and explain
one anomalous property of water which results from hydrogen bonding.
[3 + 2]
[Total 5 marks]
54. Some polar molecules are able to form hydrogen bonds. Draw a diagram to show an example of
hydrogen bonding.
[Total 2 marks]
55. At room temperature, 𝑿 is a liquid which does not conduct electricity. What does this information
suggest about the bonding and structure in 𝑿?
[Total 2 marks]
56. Sulphuric acid was added to aqueous barium hydroxide until the solution was just neutralised, forming
the insoluble salt, 𝐵𝑎𝑆𝑂4, and water.
𝑩𝒂(𝑶𝑯)𝟐(𝒂𝒒) + 𝑯𝟐𝑺𝑶𝟒(𝒂𝒒)
𝑩𝒂𝑺𝑶𝟒(𝒔) + 𝟐𝑯𝟐𝑶(𝒍)
The electrical conductivity of the solution steadily decreased as the sulphuric acid was added.
Explain why the electrical conductivity decreased.
[Total 2 marks]
57. The metal magnesium reacts with the non-metal chlorine to form a compound magnesium chloride,
𝑀𝑔𝐶𝑙2, which has ionic bonding.
(i) State what is meant by an ionic bond.
(ii) ‘Dot-and-cross’ diagrams are used to model which electrons are present in the ion.
Draw a ‘dot-and-cross’ diagram, including outer electron shells only, to show the ions present in
magnesium chloride, 𝑀𝑔𝐶𝑙2.
(iii) A student finds that solid magnesium chloride and pure water do not conduct electricity. The
student dissolved the magnesium chloride in the water and the resulting solution does conduct
electricity. Explain these observations.
[1 + 2 + 3]
[Total 6 marks]
58. In this question, one mark is available for the quality of use and organisation of scientific terms.
Nitrogen and oxygen are elements in Period 2 of the Periodic Table. The hydrogen compounds of
oxygen and nitrogen, 𝐻2𝑂 and 𝑁𝐻3, both form hydrogen bonds.
(i) Draw a diagram containing two 𝐻2𝑂 molecules to show what is meant by hydrogen bonding. On
your diagram, show any lone pairs present and relevant dipoles.
(ii) State and explain two anomalous properties of water resulting from hydrogen bonding.
[3 + 4]
[Total 7 marks]
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59. State and explain two anomalous properties of 𝐻2𝑂 that are a result of its intermolecular forces.
[4]
Quality of Written Communication [1]
[Total 5 marks]
60. The figure below shows the boiling points of four hydrides of Group 6 elements.
100 H2O
boiling point / °C
0
H2Te
H2Se
H2S
–100
(i) Explain, with the aid of a diagram, the intermolecular forces in 𝐻2𝑂 that lead to the relatively high
boiling point of 𝐻2𝑂.
[3]
(ii) Suggest why 𝐻2𝑆 has a much lower boiling point than 𝐻2𝑂.
[1]
[Total 4 marks]
61. Water forms hydrogen bonds which influences its properties.
Explain, with a diagram, what is meant by hydrogen bonding and explain two anomalous properties of
water resulting from hydrogen bonding.
[Total 6 marks]
62. Ammonia, 𝑁𝐻3, and methane, 𝐶𝐻4, are the hydrides of elements which are next to one another in the
Periodic Table.
(a) In the boxes below, draw the ‘dot-and-cross’ diagram of a molecule of each of these compounds.
Show outer electrons only.
State the shape of each molecule.
NH3
CH4
shape
(b)
shape
[3]
Ammonia is polar whereas methane is non-polar. The physical properties of the two
compounds are different.
(i) Explain, using ammonia as the example, the meaning of the term bond polarity.
(ii) Explain why the ammonia molecule is polar.
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(iii) State one physical property of ammonia which is caused by its polarity.
[4]
(c) When ammonia gas is mixed with hydrogen chloride, white, solid ammonium chloride is formed.
State each type of bond that is present in one formula unit of ammonium chloride and how many
of each type are present. You may draw diagrams.
[3]
[Total 10 marks]
[© UCLES 2013 9701/23/O/N/13]
62. Valence Shell Electron Pair Repulsion theory (VSEPR) is a model of electron-pair repulsion (including
lone pairs) that can be used to deduce the shapes of, and bond angles in, simple molecules.
(a) Complete the table below by using simple hydrogen-containing compounds. One example has
been included.
number of bond
number of
shape of
formula of a molecule
pairs
lone pairs
molecule
with this shape
NH3
3
0
trigonal planar
4
3
2
0
1
2
(b) Tellurium, 𝑇𝑒, proton number 52, is used in photovoltaic cells.
When fluorine gas is passed over tellurium at 150 °C, the colourless gas 𝑇𝑒𝐹6 is formed.
(i) Draw a ‘dot-and-cross’ diagram of the 𝑇𝑒𝐹6 molecule, showing outer electrons only.
(ii) What will be the shape of the 𝑇𝑒𝐹6 molecule?
(ii) What is the 𝐹– 𝑇𝑒– 𝐹 bond angle in 𝑇𝑒𝐹6 ?
[3]
[Total 6 marks]
[© UCLES 2013 9701/21/O/N/13]
63. (a) In 1962 Barlett made an orange yellow solid, 𝑋𝑒 + [𝑃𝑡𝐹6 ]− by reacting the noble gas xenon with
platinum (VI) fluoride.
(i) State the valence electron pair repulsion theory for predicting shapes of molecules.
(ii) Draw the shape of [𝑃𝑡𝐹6 ]− .
(iii) Suggest a reason why it has been found difficult to form the corresponding compound of
helium.
[4]
(b) Explain the difference between the bond angles in phosphine, 𝑃𝐻3 (97°) and ammonia, 𝑁𝐻3 (107°).
[2]
(c) Describe, by means of diagrams, the shapes of the following chlorine containing compounds in
terms of σ and π bonds:
(i) 𝐶𝐻3 𝐶𝑙
(ii) 𝐶2 𝐻2 𝐶𝑙
(iii) 𝐶6 𝐻5 𝐶𝑙
[6]
[Total 12 marks]
[© ZIMSEC 9189/1 N2010]
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