Chemical Bonds
Ionic
(Metal & NonMetal)
Covalent
(2 or > nonmetals)
Metallic
Intermolecular
forces
Hydrogen bond
Van Der Waals
forces
Substances
Giant Structure
- High melting & boiling
points
Ionic compounds
– forms giant ionic
structures
Graphite
Giant covalent structure
- Atoms are held by
strong covalent bonds
-No van der waals forces
Diamond
Silicon dioxide
Metals
Simple molecular
structure
- Low melting & boiling
points
Covalent
- Molecules are held by
weak van der waals
forces
 Nature
of the bond
 Physical properties
- Melting point / boiling point
- Electrical conductivity
- Solubility
- Hardness
 electrostatic
attraction between oppositely
charged ions
 The charged ions are formed when atoms
lose or gain electrons.
Sodium atom Na 2,8,1
Chlorine atom Cl
2,8,7
Sodium ion
Na+
2,8
Chloride ion
Cl2,8,8
Na x
 [Na]+ + e-
Cl + e-  [ Cl ]Na
x
+ Cl  [Na]+[ Cl x]-
Sodium metal reacts with chlorine gas in a violent exothermic
reaction to produce NaCl.
What ions will be formed by these elements?
Na , Mg , Al , P , S and Cl
Draw the dot and cross diagram for magnesium
oxide
 The
transition elements form more than one
stable positive ion.
Name of Transition Element
Positive Charge
Silver
Ag+
Iron
Fe2+ , Fe3+
Copper
Cu+ , Cu2+
Manganese
Mn2+ , Mn3+ , Mn4+
Chromium
Cr2+ , Cr3+ (not stable in air)
-
-
Melting point / boiling point
Electrical conductivity
Solubility
Hardness
Volatility
Students’ posters
-
Hardness
Hard but brittle because a slight
displacement of a layer of ions will cause
repulsion between the ions and the crystal
will break.
Volatility
(how readily a substance evaporates)
- Low volatility because of the strong
electrostatics forces between ions.
-
 Ionic
bonding typically occurs between metal
and non-metal. E.g. Barium fluoride, BaF2
 The reactivity of metals and non-metals can
be assessed using electronegativity (ability of
an atom in a covalent bond to attract shared
pairs of electrons to itself).
 Metals
generally have low electronegativity
values, while non-metals have relatively
high electronegativity values.
 Fluorine, which has the greatest attraction
for electrons in bond-forming situations
(highest E value).
If the difference in E values is > 1.8 => ionic bond
 If the difference in E values is 0, non-polar
covalent bond
 If the difference is 0 – 1.8, polar covalent bond

Polar covalent bonds are covalent bonds with
ionic character.
Ionic bond
Non polar
covalent bond
Polar
covalent bond
Electrons are not shared.
E.g. Na+ Cl- , electron is transferred.
Electrons are equally shared.
E.g.Cl-Cl
Electrons are not equally shared.
E.g.  +  -
H  Cl
Atoms have different
electronegativity values
 Use
the table above to predict the type of
bonding beween
Fluorine , F2
Hydrogen iodide, HI and
Lithium fluoride, LiF
Name of ion
Formula Example of compound
Ammonium
NH4+
NH4Cl, ammonium chloride
Hydroxonium
H3O+
H3O+Cl-, hydrochloric acid
Sulfate
SO42-
MgSO4 , magnesium sulfate
Hydrogencarbonate HCO3-
KHCO3 , potassium hydrogencarbonate
Nitrate
NO3-
AgNO3 , silver nitrate
Phosphate
PO43-
K3PO4 , potassium phosphate
Hydroxide
OH-
NaOH , sodium hydroxide
Carbonate
CO32-
Na2CO3 , sodium hydroxide

In an ionic compound,
constituent ions are held in fixed
positions in an orderly
arrangement by strong ionic
bonds.
A crystal of NaCl consists of a giant
lattice of Na+ and Cl- ions.
All Na+ ions in the crystal attract
all the Cl- ions, and vice versa.
This strong attraction between
oppositey
charged ions holds the crystal
together.
It is difficult to break apart the
lattice.
Page 88
 forms
when atoms share electrons
 A covalent bond is the electrostatic attraction
between the shared pair of electrons and the
nuclei of the atoms making the bond.
E.g. Formation of methane molecule, CH4
Lewis Structures of methane
Bond pair _____
Lone pair
..
Page 90




The bonding pair of electrons spends most of its time
between the two atomic nuclei, thereby screening the
positive charges from one another and enabling the nuclei
to come closer together than if the bonding electrons were
absent. Negative charge on the electron pair attracts both
nuclei and holds them together in a covalent bond.
When two atoms are chemically bonded, the two atoms
close together have less energy and therefore are more
stable than when separated.
Energy is given off by the atoms to form a bond, and
energy must be supplied (absorbed) to break the bond.
A covalent bond is the result of electrostatic attraction
between the nuclei of the 2 atoms and the pair of shared
electrons.
Bond pair _____
Lone pair
..
Draw Lewis structures of the following molecules:
Chlorine, Cl2
Hydrogen chloride, HCl
Oxygen, O2
Nitrogen, N2
Nitrogen Trifluoride, NF3
Carbon dioxide, CO2
Water, H2O
Hydrogen peroxide, H2O2
Ethene, C2H4
Follow structures on page 91
Bond pair _____
Lone pair
..
Draw Lewis structures of the following molecules:
BH3 , NH3 , PH3
In which of the above obey the octet rule?
 Not
all covalent bond results in a noble gas
configuration.
In BF3, B has only 6 electrons in its outer shell.
In SF6, S has 12 electrons in its outer shell
(expanded octet)
Page 92
 Bond
Strength
Triple bonds > Double bonds > Single bonds
(1)
(2)
The attraction between the 2 nuclei for 3
electron pairs in a triple bond is > that for
2 electron pairs in a double bond which is >
than that for 1 electron pair in a single
bond.
Triple bonds are shorter due to greater
attraction between the bonding electrons
and the nuclei with more electrons in the
bond.
Page 93
 In
some molecules and polyatomic ions, both
electrons to be shared come from the same
atom. The covalent formed is called the
coordinate or dative bond.
 Carbon monoxide (CO) can be viewed as
containing one coordinate bond and two
"normal" covalent bonds between the C atom
and the O atom.
How do you draw the Lewis structure?
Page 94
A
thick white smoke of solid ammonium
chloride is formed in the reaction below:
 Ammonium
ions, NH4+, are formed by the
transfer of a hydrogen ion from the hydrogen
chloride to the lone pair of electrons on the
ammonia molecule.
Page 94, Lewis acid-base reaction
When the ammonium ion, NH4+, is formed, the
fourth hydrogen is attached by a dative covalent
bond, because only the hydrogen's nucleus is
transferred from the chlorine to the nitrogen.
The hydrogen's electron is left behind on the
chlorine to form a negative chloride ion.
 Once the ammonium ion has been formed it is
impossible to tell any difference between the
dative covalent and the ordinary covalent bonds.

 Something
similar happens. A hydrogen ion
(H+) is transferred from the chlorine to one
of the lone pairs on the oxygen atom.
 The
H3O+ ion is variously called the
hydroxonium ion.
Other examples:
 The reaction between ammonia and boron
trifluoride, BF3
In BF3, there are only 6 electrons in the outer shell of boron.
There is space for the B to accept a pair of electrons.
Rules





Calculate the total no. of valence electrons for all
atoms in the molecule or ion.
Divide by 2 to get the no. of electron pairs.
Each electron pair is represented by a line.
Arrange the lines (electron pairs) so that all the
atoms are joined together by at least a single bond
and the outer atoms have 8 electrons in their outer
shell (except for H)
Rearrange the lines (electron pairs) so that every
period 2 atom has 4 pairs of electrons. The outer
atoms already have 4 pairs, so this should normally
only involve moving lone pairs so that they become
bonding pairs of eectrons.
Write the Lewis structure
NF3 , CO32- , NO2- , O3
 Strength
Triple bonds > Double bonds > Single bonds
 Length
Single bonds > Double bonds > Triple bonds
Bond Type
Length
(nm)
Strength
(kJmol-1 )
C-C
0.154
348
C=C
0.134
612
CΞC
0.120
837
 In
diatomic molecules (e.g. H2 ,Cl2) both
atoms exert an identical attraction.
 When the atoms are different (e.g. HCl) with
one more electronegative than the other, a
polar bond is formed.
 Relative polarity is predicted from
electronegativity values.
Element
F
O
N
Cl
C
H
Electronegativity
4.0
3.5
3.5
3.0
2.5
2.1
 C-O
is more polar than C-Cl since the
difference in E value for C-O is greater than
that for C-Cl.
 The
-
-
shapes of simple molecules and ions can
be determined by using the Valence Shell
Electron Repulsion (VSEPR) theory.
Electron pairs around the central atom repel
each other
Bonding pairs and lone pairs arrange
themselves to be as far apart as possible
The shape of a molecule
depends on the number of
electron pairs in the outer
shell of the central atom.
The 5 basic molecular shapes
show the arrangement of
the electron pairs (charge
centres) that result in
minimum repulsion
between the bonding and
lone pairs of electrons.
Focus on 3
Basic shape : Tetrahedral arrangement
Tetrahedral
arrangement
Trigonal pyramidal
Bent
 Draw
the Lewis structure for the molecule or
ion.
 Count up the number of electron pairs
(bonding and lone pairs) in the outer shell of
the central atom. Multiple bond is counted
as single electron pair because the electrons
occupy the same space region.
 Check the table (page 100) to get the basic
shape. Draw the 3-D shapes.
 A lone pair is just an electron pair in the
outer shell of an atom - cannot be “seen”
 State the actual shape of the molecule.
Consider
CH4 , NH3 , H2O , CO2 , SO2
Order of repulsion :
lone pair-lone pair > lone pair- bonding pair > bonding pair – bonding pair
Lone pairs are held closer to the nucleus than the bonding pairs.
The distance between the lone pair electrons and the bonding pairs of
electrons is shorter than the distance between the bonding pairs to
each other.
Repulsion due to lone pairs causes the bond angles to become smaller
Order of repulsion :
lone pair-lone pair > lone pair- bonding pair > bonding pair – bonding pair
Methane, CH4
Bond angle is 109.50
Ammonia, NH3
Greater repulsion by lone
pair of electrons.
Bond angle is smaller
than 109.50(1050)
Water, H2O
Even greater repulsion by
two lone pair of electrons.
Bond angle is even smaller
(1050)
Consider
NH4+ , H3O+
,
NO2-
NH4+
As the 4 negative charge centres repel each other
and take up positions in space to be as far apart as
possible, the electron pairs are distributed in
tetrahedral arrangement.
H3O+
As the 4 negative charge centres repel each other and take
up positions in space to be as far apart as possible, the
electron pairs are distributed in tetrahedral arrangement.
With one lone pair of electrons, the actual structure is
trigonal pyramid with a bond angle of 1070 for H-O-H bond.
Read page 104
NO2As the 3 negative charge centres repel each other and take
up positions in space to be as far apart as possible, the
electron pairs are distributed in trigonal planar arrangement.
With one lone pair of electrons, the actual structure of the
ion is bent with a bond angle of about 1170 for O-N-O bond.
Consider
N2H4 , C2H2
 Find
the number of electron pairs / charge
centres in the valence shell of the central
atom.
 Electron pairs / charge centres repel each
other to the positions of minimum energy in
order to gain maximum stability.
 Pairs forming a double or triple bond act as a
single bond
 Non-bonding pairs repel more than bonding
pairs.
Due to difference in electronegativity value between the 2
atoms in the bond.
 Unequal distribution of electron density results in small
charges on the atoms
( δ+ and δ- )
Example

A dipole is established when two electrical charge of opposite sign
are separated by a small distance.
Dipole moment
Non-polar Covalent bond
 No difference in
electronegativity value –
bond consists of
2 ____________ atoms.
 _______ net charge.
Examples :
Polar Covalent bond
 Due to the difference in
electronegativity value –
bond consists of
2 ____________ atoms.
 _______ net charge.
Examples:

-
Polarity of a molecule depends on the
the relative electronegativities of the atoms in
the molecule and the shape of the molecule
Which of the following is the most polar?
HF, HCl, HBr or HI
F, Cl, Br and I are in the same group (halogen).
Electronegativity decreases from F to I.
Since bond polarity is due to the difference in
electronegativity value, H-F is the most polar bond.
Name of molecule
Formula
Polarity of molecule
Hydrogen chloride
HCl
Polar
Water
H2O
Polar
Ammonia
NH3
Polar
Benzene
C6H6
Non-polar
Boron trichloride
BCl3
Non-polar
Methane
CH4
Non-polar
Bromobenzene
C6H5Br
Polar
Carbon dioxide
CO2
Non-polar
Sulfur dioxide
SO2
Polar
Tetrachloromethane
CCl4
Non-polar
For CO2 each C-O bond is polar since O is more
electronegative than C.
Why is the molecule non-polar?
 There
are some instances when the polar
bonds are arranged symmetrically so as to
give zero net direction of charge.
i.e. Overall dipoles cancel so that there is no
overall dipole.
For example,
carbon dioxide and carbon tetrachloride,
 Some
molecules have very low polarity - so
low as to be regarded as non-polar,
 Forces
between molecules.
 Does not exist in giant structure (ionic cpds,
metals & giant covalent structure)
Intermolecular
forces
Van der Waals
forces
Hydrogen bond
http://chemtools.chem.soton.ac.uk/projects/emalaria/index.php?page=13
 The
strength of the intermolecular forces
determines how easily the molecules will
separate and hence the melting and boiling
points.
Why is there an increasing boiling points of the noble gases as
you go down the group?
The boiling points of the noble gases are
Helium
Neon
Argon
Krypton
Xenon
Radon
-269°C
-246°C
-186°C
-152°C
-108°C
-62°C




Electrons can at any moment be
unevenly spread producing a
temporary instantaneous
(fluctuating) dipole.
An instantaneous dipole can induce
another dipole in a neighbouring
particle resulting in a weak
attraction between the two
particles.
The forces of attraction between
temporary or induced dipoles are
known as Van der Waals’ forces
(London Dispersion Forces).
Van der Waals’ forces increases
with increasing mass.
Because electrons are always moving around very
quickly, the charges switch around all the time.
 the more electrons in a molecule / atom, the
stronger these Van der Waals or London forces
are.
 This is seen in the increasing boiling points of the
noble gases as you go down the group.

The boiling points of the noble gases are
Helium
Neon
Argon
Krypton
Xenon
Radon
-269°C
-246°C
-186°C
-152°C
-108°C
-62°C
Cl2 : gas
I2 : solid
Iodine molecule is made up of larger atoms
with more electrons compared to chlorine.
With more electrons moving around, the
temporary dipole will be larger.
The larger atoms in the molecule means that
the valence electrons are less strongly held,
Hence the induced dipoles will be larger.
van der Waals forces are present between
covalent molecules with no H atom attached to
N, O or F.
 E.g. van der Waals forces are present between
HCl molecules.
 There’s also other intermolecular forces
beween the molecules :
permanent dipole - permanent dipole

 These
intermolecular forces between polar
molecules are stronger than between nonpolar molecules. (all things being equal)
 For polar substances with similar RMM, the
higher the dipole moment, the stronger the
dipole-dipole attractions and the higher the
boiling points.
Propane (C3H8) and ethanal (CH3CHO) both with RMM
= 44
 Ethanal has a higher bp.
Ethanal
- is a polar molecule
- has stronger intermolecular forces (van der waals &
dipole-dipole interactions) between the molecules of
ethanal than between the propane molecules.

It is not true that polar molecules have stronger
intermolecular forces and hence higher bp than
non-polar molecules.
Non-polar molecules with higher RMM might have higher bp.
Read further for a few exceptions

A hydrogen bond is a weak type of force that
forms a special type of dipole-dipole attraction
which occurs when a hydrogen atom bonded to a
strongly electronegative atom exists in the
vicinity of another electronegative atom with a
lone pair of electrons. These bonds are generally
stronger than ordinary dipole-dipole and
dispersion forces, but weaker than true covalent
and ionic bonds.
 Hydrogen
bonding is present between
covalent molecules with H atoms attached to
O, N and F
 Strength
of H bond
The larger the electronegativity of H and the
other atom (N, O or F), the stronger the H
bond.
Strength F > O > N
 Number of them that can be formed between
neighbouring molecules
Although the strength is such F > O > N, HF can only form 1 H bond to 1
neighbour.
H2O can form 2, thus promoting more intermolecular interactions .
The collective strength of the H bonds in water is greater than the
strength of the H bonds in HF because each O atom (with 2 lone
pairs) in the water molecule can form 2 H bonds with 2 other water
molecules, whereas each F atom in HF molecule can only form 1 H
bond with another HF molecule.
Ammonia molecule has 3 N-H bonds. N is
larger and < electronegative than F, has
far weaker H bonds due to lower electron
density on the N atom (only 1 lone pair)
compared to O and F.
 Between
1 single molecule
 Between
2 like molecules
 Between
2 unlike molecules
When the RMM is large, we expect the boiling
point to be high because larger molecules have
more space for electron distribution and more
possibilities for instantaneous dipole moment.
 However,

Compound
RMM
Boiling pt (K)
H2O
18
373
HF
20
292.5
NH3
17
239.8
H2S
34
212
HCl
36.4
187
PH3
34
185.2
Greater intermolecular
force because H2O, HF,
NH3 all exhibit hydrogen
bonding .
tend to have higher viscosity
than those that do not have
H bond.
Substances which have multiple
H bonds exhibit even higher
Viscosity.
 Electronegativity
Cannot occur without significant
electronegativity difference between H and
the atom it is boded to.
E.g. Both PH3 and NH3
have trigonal pyramidal shape but only NH3 has
H bonding.
 Atomic size
When the radii of the 2 atoms differ greatly,
their nuclei cannot achieve close proximity
when they interact resulting in weak
interaction.
Is there any hydrogen bonding between the
molecules if CH3F?
H
H
C
H
H
F
H C F
H
H is not joined directly to F in each molecule,
hence no hydrogen bonding between the molecules.
Is there any hydrogen bonding between the
molecules of ethanol?
Hydrogen bonding affects
 the boiling points of water, ammonia,
hydorgen fluoride and other molecules
 the solubility of simple covalent molecules
such as ammonia, methanol and ethanoic
acid in water
 the density of water and ice.
 the viscosity of liquids, e.g. the alcohols.
H2S, H2Se and H2Te molecules are held by van der waals forces.
Water molecules are held by hydrogen bond.
Hydrogen bond is much stronger than van der waals forces,
hence water has ahigher boiling point.
Types of Covalent Substances
Covalent substances can be divided into 2
categories as shown in the table below:
Simple Covalent
Molecules
Macromolecules
Examples: Hydrogen,
Examples: Diamond,
Oxygen, Water, Carbon Graphite and sand
Dioxide and Methane
(Silicon Dioxide)
The atoms can be either same like silicon and
carbon (graphite and diamond) or of 2 different
elements such as silicon dioxide.
 Allotropes are two (or more) crystalline forms of
the same element, in which the atoms ( or
molecules) are bonded differently.

Macromolecules
Molecules with giant molecular structures are called macromolecules.
Diamond
Graphite
Giant Molecular 1. Each carbon atom has 4 1. Three valence electrons
Structure
valence electrons.
in each carbon atom are
used for covalent bonding.
1. Each carbon atom is
2. The fourth electron is not
joined to 4 other carbon
used in chemical
atoms by strong
bonding.
covalent bonds in a
3. This gives hexagonal rings
of six atoms that join
tetrahedral
together to from flat layers.
arrangement.
4. These layers of carbon
atoms, lie on top of each
other and are held together
by van der Waal’s forces.
Diamond
Diagrammatical
Representation
Graphite
Structure of Graphite
Covalent Bonding
Properties
Uses
Diamond
1. Hard
2. Very high melting
point and boiling
point
3. Non-conductor of
electricity
Graphite
1. Soft and slippery
2. High melting point
and boiling point
3. Good conductor of
electricity
1. Gemstones
2. As tips of grinding,
cutting and
polishing tools.
1. In pencils
2. As a dry lubricant
3. Brushes for electric
motors
Fullerene
 60 C atoms are arranged in
hexagons and pentagons to give
a geodesic spherical structure
similar to a football.
 Following
the discover of Buckminsterfullerene
, many other similar carbon molecules have been
isolated.
 This has led to a new branch of science called
nanotechnology.
Recall:
In ionic compounds, the ions are held together by
strong ionic bonds in a giant ionic lattice.
In simple covalent molecules, the attractive
forces between the molecules are known as
intermolecular forces or van der Waal’s forces,
which is weaker than the ionic bonds.
In giant covalent molecules, the atoms are held
together by strong covalent bonds in a giant
covalent lattice.
4
Metallic Bonds
Metals consist of positive ions surrounded by a 'sea of moving
electrons'.
The negative 'sea of electrons' attracts all the positive ions and
cements everything together.
Metallic bonds are the results of the strong forces of attraction
between the negative electrons and the positive ions.
Hence, metals have high melting points and high boiling points.
Physical Properties of Metals
Physical
Properties
High Density
Explanation
The close packing of
atoms explains why
most metals have a
high density
Good Conductor Metals are good
of Electricity
conductors of
electricity in the solid
and molten state.
This is due to the sea
of delocalized
electrons.
Diagram
Physical
Properties
Malleable and
Ductile
Explanation
Metals are malleable
(can be bent or
flattened) and ductile
(can be drawn into
wires) because the
layers of metal ions can
slide over each other
when a force is applied.
High melting and Metallic bonds are
strong bonds.
boiling point
Except for: Mercury has
a low melting point and
is a liquid at room
temperature.
Similarly, sodium and
potassium have low
melting and boiling
point.
Diagram
Explanation

The valence electrons do not belong to any particular
atom, hence, if sufficient force is applied to the
metal, 1 layer of metals can slide over another
without disrupting the metallic bonding.

The metallic bonding in metal is strong and flexible
and so metals can be hammered into thin sheets
(malleability) or drawn into lonng wires (ductility)
without breaking.
If atoms of other elements are added by alloying, the
layers of ions will not slide over each other so readily.
The alloy is thus less malleable and ductile and
consequently harder and stronger.
‘Like tends to dissolve like’. Polar substances
tend to dissolve in polar solvents, such as water,
whereas non-polar substances tend to dissolve in
non-polar solvents, such as heptane or
tetrachloromethane.
 Organic molecules often contain a polar head
and a non-polar carbon chain tail. As the nonpolar carbon chain length increases in an
homologous series the molecules become less
soluble in water.
 Ethanol is a good solvent for other substances as
it contains both polar and non-polar ends.

 Water
will mix with polar liquids such as
ethanol. The oppositely charged ends of the
different molecules attract one another
forming hydrogen bonds.
 Gases
are generally slightly soluble in water.
 A small number of gases are highly soluble
because they react with water to release
ions.
Example,
SO2(g) + H2O (g)
H+(aq) + HSO3-(aq)
This solution is known as sulfurous acid , a
major component of acid rain