HSC chemistry 1st paper: Chapter -3: Chemical bond

1. C3:Periodic properties and Chemical bond
All elements are unstable and tend to combine with each other as well as with other elements
to form molecules. This force of attraction that binds the constituent atoms together in a
molecule are termed as chemical bond
Atoms combine to
(a) acquire a minimum energy state.
(b) acquire a stable noble gas configuration by completing their octet (i.e.
outermost shell) for most of the elements except hydrogen, helium, lithium and
Types of Chemical Bonds
 Ionic or electrovalent bond formed by the
to another to form.
 Covalent bond that results when atoms
 Normally, the electrons to be shared are contributed by both atoms. But if the electrons to
be shared between the two atoms are contributed by only one of these atoms, it is called
coordinate bond.
 Metallic bond forms when metals release their valence elec
(positive ions).
The atoms of different or same elements combine with each other to acquire stability by
completing their
 Octet: 8 electrons in the outermost valence shell or
 Duplet: 2 electrons in the outermost valence shell (
Lewis symbols are represented by placing dots (
of an atom. For example,
1
Chemistry 1st
paper
C3:Periodic properties and Chemical bond
Part: Chemical Bond
elements are unstable and tend to combine with each other as well as with other elements
to form molecules. This force of attraction that binds the constituent atoms together in a
chemical bond.
mum energy state.
(b) acquire a stable noble gas configuration by completing their octet (i.e.
outermost shell) for most of the elements except hydrogen, helium, lithium and
formed by the transfer of one or more electrons from one atom
bond that results when atoms share electrons.
Normally, the electrons to be shared are contributed by both atoms. But if the electrons to
be shared between the two atoms are contributed by only one of these atoms, it is called
forms when metals release their valence electron that bind metal ions
The atoms of different or same elements combine with each other to acquire stability by
8 electrons in the outermost valence shell or
2 electrons in the outermost valence shell (only in case of H, Li or Be)
Lewis symbols are represented by placing dots (.) as valence electrons around the symbols
elements are unstable and tend to combine with each other as well as with other elements
to form molecules. This force of attraction that binds the constituent atoms together in a
8 electrons in the
outermost shell) for most of the elements except hydrogen, helium, lithium and beryllium.
of one or more electrons from one atom
Normally, the electrons to be shared are contributed by both atoms. But if the electrons to
be shared between the two atoms are contributed by only one of these atoms, it is called
tron that bind metal ions
The atoms of different or same elements combine with each other to acquire stability by
only in case of H, Li or Be)
as valence electrons around the symbols

2. 2
IONIC OR ELECTROVALENT BOND
A bond formed by the complete transfer of one or more valence electrons from an electropositive
atom of an element to an electronegative atom of another element, so that both the atoms
complete their octet (duplet in special cases) is called an ionic bond or electrovalent bond.
Formation of Ionic compounds
NaCl (Sodium chloride)
First Atom (11Na) Second Atom (17Cl)
Electronic configuration 1s2
2s2
2p6
3s1
1s2
2s2
2p6
3s2
3p5
Transfer of electron(s)
Attained electronic
configuration after transfer
of electrons
Na+
: 1s2
2s2
2p6
(Stable
configuration of Ne)
Cl-
: 1s2
2s2
2p6
3s2
3p6
(Stable
configuration of Ar)
Ionic bond formation:
These two ions, Na+
and Cl-
, are held together by electrostatic forces of attraction.
Lattice energy
The amount of energy released when free ions combine together to form one mole of a crystal is
called lattice energy (V).
M+
(g) + X-
(g) MX(s) + lattice energy
Ionic compound
To sum up, the favorable conditions required for a stable ionic bond are the following:
(i) Elements (metals) having low ionization energy
(ii) Elements (non-metals) having high electron affinity (EA)
(iii) Small sized ions with higher charges i.e. high lattice energy.

3. 3
Properties of Ionic compounds
1. These compounds have ions as constituent particles.
2. They have high melting and boiling points.
3. They can conduct electricity in solution or in molten state.
 COVALENT BOND: Ionic bonding cannot result from a reaction between two non-metals
because their electronegativity difference is not great enough for electron transfer to take
place. However, reactions between two non-metals result in covalent bonding. Covalent
bond is a bond formed by the mutual sharing of electrons between the atoms of the same
or different elements to acquire a noble gas configuration. A covalent bond can be single
or multiple (double or triple)
 Single covalent bond: If two atoms share one electron pair, the bond is known as single
covalent bond.
Examples: Cl2, H2, NH3, HCl, CH4, H2O etc.
 Double covalent bond: If two atoms share two electron pairs, the bond is known as double
covalent bond.
Examples: Oxygen (O2), ethene (C2H4), carbon dioxide (CO2), etc.
 Triple covalent bond: If two atoms share three electron pairs, the bond is known as triple
covalent bond.
Examples: Nitrogen (N2), ethyne (C2H2), etc.
Formation of covalent Bond
Covalent bond formation:
Water (H2O)
First Atom (1H) Second Atom (8O)
Electronic configuration 1s1
1s2
2s2
2p4
No of electrons required
for completing octet
1
(Two atoms required)
2
Attained electronic
configuration after sharing
of electrons
1s2
(Stable configuration
of He)
1s2
2s2
2p6
(Stable
configuration of Ne)

4. Covalent bond formation:
> EXCEPTIONS TO OCTET RULES
Electron-deficient species (Incomplete octet)
The central atom in these species has less than eight electrons in its valence
less than four valence electrons usually has incomplete octet
The covalent compounds of group 13 elements form electron
Examples are boron trifluoride (BF
Electron-rich species: (super octet species)
Species in which the central atom has more than eight electrons in their valence shells are
termed electron-rich species. Non-
phosphorus, chlorine, etc. with an empty
Examples are phosphorus pentachloride (PCl
 NCl5 doesn’t form but PCl5 froms.
There are 3 electrons of each nitrogen and phosphorus is
produce NCl3 and phosphorus can produce PCl
orbital so it can extend valence electron by utilizing it. When one electron jump to 3d
orbital from 3s2
then total valence electron becomes 5.
But nitrogen has no vacant orbital so that NCl
1122
22221:)7( yx ppssN
2622
33221:)15( pspssP
1622
33221:)15( pspssP
 VALENCE SHELL ELECTRON PAIR REPULSION (VSEPR) THEORY
Basic Concepts of VSEPR Theory
The decreasing order of repulsive forces between different types of electron pairs is:
Lone pair-lone pair > Lone pair
4
(Incomplete octet)
The central atom in these species has less than eight electrons in its valence
less than four valence electrons usually has incomplete octet in its compounds.
group 13 elements form electron-deficient compounds.
Examples are boron trifluoride (BF3), aluminium chloride(AlCl3), beryllium ch
: (super octet species)
Species in which the central atom has more than eight electrons in their valence shells are
-metallic elements of third group and higher elements like silicon,
hosphorus, chlorine, etc. with an empty d orbital are capable of forming such species.
are phosphorus pentachloride (PCl5), sulphur hexafluoride (SF6).
froms.
There are 3 electrons of each nitrogen and phosphorus is unpaired so that nitrogen can
and phosphorus can produce PCl3 normally. As phosphorus has vacant 3d
orbital so it can extend valence electron by utilizing it. When one electron jump to 3d
then total valence electron becomes 5. So that phosphorus can form PCl
But nitrogen has no vacant orbital so that NCl5 doesn’t form.
1
2 zp
111
33 zyx ppp
1111
333 dppp zyx
VALENCE SHELL ELECTRON PAIR REPULSION (VSEPR) THEORY
Theory
The decreasing order of repulsive forces between different types of electron pairs is:
lone pair > Lone pair – bond pair > Bond pair- bond pair
The central atom in these species has less than eight electrons in its valence shell. An atom with
in its compounds.
deficient compounds.
, beryllium chloride (BeCl2).
Species in which the central atom has more than eight electrons in their valence shells are
metallic elements of third group and higher elements like silicon,
orbital are capable of forming such species.
unpaired so that nitrogen can
normally. As phosphorus has vacant 3d
orbital so it can extend valence electron by utilizing it. When one electron jump to 3d
So that phosphorus can form PCl5.
The decreasing order of repulsive forces between different types of electron pairs is:

5.  Theory of overlapping of atomic orbitals:
Sigma( ) bond: The covalent bond formed by head to head overlapping of two orbitals is called
sigma bond.
Pi ( ) bond: The covalent bond formed by side by side overlapping of two orbitals is called
pibond.
 HYBRIDISATION: hybridisation, which is defined as
The process in which two or more atomic orbitals having different energies mixed together
and produce equal number of new orbitals having equal energies is called hybridisation.
The orbitals are called hybrid orbita
> sp3
Hybridisationor Tetrahedral Hybridisation
 It involves the combination of one
four sp3
hybridised orbitals.
 The four sp3
hybridised orbitals are directed towards the corners of a regular tetrahedron.
5
Theory of overlapping of atomic orbitals:
) bond: The covalent bond formed by head to head overlapping of two orbitals is called
) bond: The covalent bond formed by side by side overlapping of two orbitals is called
hybridisation, which is defined as –
The process in which two or more atomic orbitals having different energies mixed together
and produce equal number of new orbitals having equal energies is called hybridisation.
The orbitals are called hybrid orbitals.
Hybridisationor Tetrahedral Hybridisation
It involves the combination of one s and three p orbitals and results in the formation of
hybridised orbitals.
hybridised orbitals are directed towards the corners of a regular tetrahedron.
) bond: The covalent bond formed by head to head overlapping of two orbitals is called
) bond: The covalent bond formed by side by side overlapping of two orbitals is called
The process in which two or more atomic orbitals having different energies mixed together
and produce equal number of new orbitals having equal energies is called hybridisation.
orbitals and results in the formation of
hybridised orbitals are directed towards the corners of a regular tetrahedron.

6.  The bond angle is 109.
50
.
Example: Methane (CH4), carbon tetracloride (CCl
(C2H6) etc.
Formation of methane (CH4)
(ground state) 1s2
2s2
2p2
Formation of a
 sp2
Hybridisation or Trigonal Hybridisation
 It involves the combination of one s and two
three sp2
hybridised orbitals.
 The three sp2
hybridised orbitals are directed towards the corners of a regular equilateral
triangle. Hence, the shape of the molecule becomes planar trigonal.
6
), carbon tetracloride (CCl4), water (H2O), ammonia (NH
Formation of a CH4 molecule
Hybridisation or Trigonal Hybridisation
It involves the combination of one s and two p orbitals and results in the formation of
hybridised orbitals.
hybridised orbitals are directed towards the corners of a regular equilateral
triangle. Hence, the shape of the molecule becomes planar trigonal.
O), ammonia (NH3), ethane
orbitals and results in the formation of
hybridised orbitals are directed towards the corners of a regular equilateral

7.  The bond angle is 120o
.
boron triflouride (BF
 sp Hybridisation or Linear Hybridisation
It involves the combination of one
hybridised orbitals.
The two hybridised sp orbitals have 50%
They are directed towards the two opposite ends of a straight line, hence, the bond angle is 180
They have a linear shape.
Examples: C2H2, CO2, BeCl
 Formation of Ammonia (NH3)
Nitrogen has three bonds around itself.
the time of bond formation.
N(7):1s2
2s2
2px
1
N in NH3
7
boron triflouride (BF3)
ar Hybridisation
It involves the combination of one s and one p orbital and results in the formation of two
orbitals have 50% of s character and 50% of p character.
two opposite ends of a straight line, hence, the bond angle is 180
, BeCl2 etc.
Nitrogen has three bonds around itself. So, it must have three unpaired electrons
2py
1
2pz
1
orbital and results in the formation of two sp
character.
two opposite ends of a straight line, hence, the bond angle is 180o
.
must have three unpaired electrons at

8. 1. We clearly see that the orbital with the lone pair is also involved in the hybridisation.
2. The lone pair of electrons will show an increased
hence, the shape of the molecule will no longer be a regular tetrahedron, but
will become a trigonal pyramidal and the angle
[why the bond angle of NH3 is 107
 Formation of water (H2O)
Oxygen has two bonds around itself, hence, it must have two
O(8): 1s2
2s2
2p4
O (in H2O)
1. We clearly see that the two orbitals
hybridisation.
2. The two lone pairs of electrons show maximum repulsion, hence, the bond angle is further
reduced to 104.5o
, instead of 109
3. The shape of the molecule becomes '
[why the bond angle of H2O is 104.5
8
1. We clearly see that the orbital with the lone pair is also involved in the hybridisation.
of electrons will show an increased repulsion as per the
hence, the shape of the molecule will no longer be a regular tetrahedron, but
a trigonal pyramidal and the angle will be reduced to 107
is 1070
?]
has two bonds around itself, hence, it must have two unpaired electrons around itself.
We clearly see that the two orbitals that are fully filled are also involved in the
The two lone pairs of electrons show maximum repulsion, hence, the bond angle is further
, instead of 109.
50
.
The shape of the molecule becomes ' ' bent.
O is 104.50
?]
1. We clearly see that the orbital with the lone pair is also involved in the hybridisation.
repulsion as per the VSEPR theory,
hence, the shape of the molecule will no longer be a regular tetrahedron, but
will be reduced to 107o
.
around itself.
involved in the
The two lone pairs of electrons show maximum repulsion, hence, the bond angle is further

9. sp3
d Hybridisation
 Shape is bipyramidal
 The bond angles are 90o
and 120
 Examples: PCl5, PF5 etc.
sp3
d2
Hybridisation
 The bond angles are 90o
.
 The shape is octahedral.
Examples: SF6, IOF5 etc.
sp3
d3
Hybridisation
 The only example of this kind is IF
 Shape is pentagonal bi-pyramidal
 Coordinate covalent bonds or Dative bonds:
formed between two atoms, by the donation of an electron pair (lone pair) by one atom, but
shared by both the atoms, as to co
bond or a coordinate bond or a
Here, the atom that provides the electron pair is known as the
receives the electron pair is known as the
Examples are sulphur dioxide (SO2
ammonium ion (NH4
+
) etc.
Ammonia and Boron Trifluoride Complex
a) The ammonia molecule has a lone pair of electrons with the nitrogen atom.
b) The BF3 molecule is short of an electron pair as already studied, due to an incomplete octet of the
c) The boron atom acts as an acceptor and the
coordinate bond is formed between the two molecules, resulting in
complex.
9
and 120o
.
kind is IF7
pyramidal
Coordinate covalent bonds or Dative bonds: A coordinate bond can be defined
by the donation of an electron pair (lone pair) by one atom, but
shared by both the atoms, as to complete their octets. This is known as a
coordinate bond or a dative bond.
provides the electron pair is known as the donor and the other atom, which
receives the electron pair is known as the acceptor.
2), sulphuric acid (H2SO4), hydronium ion (H
Ammonia and Boron Trifluoride Complex
a) The ammonia molecule has a lone pair of electrons with the nitrogen atom.
f an electron pair as already studied, due to an incomplete octet of the
boron atom acts as an acceptor and the nitrogen atom of ammonia acts as
coordinate bond is formed between the two molecules, resulting in an ammonia
coordinate bond can be defined as a bond
by the donation of an electron pair (lone pair) by one atom, but
coordinate covalent
and the other atom, which
), hydronium ion (H3O+
),
a) The ammonia molecule has a lone pair of electrons with the nitrogen atom.
f an electron pair as already studied, due to an incomplete octet of the boron atom.
nitrogen atom of ammonia acts as a donor. A
an ammonia-boron trifluoride

10. Ammonium ion (NH4
+
)
a) In NH3, the N atom has an ionic pair of electrons.
b) H+
is short of an electron pair to complete its octet.
c) The N atom acts as a donor and H
between the two.
> Covalent character of an ionic bond
Although the ionic bond is considered to be 100% ionic, actually it has some covalent
character, just as a covalent bond has some ionic character.
1. When a cation approaches an anion, the electron cloud of
towards the cation and hence, gets distorted. This effect is called
anion. As shown in the Figure 1 below.
2. The power of the cation to polarise the anion is called its
the anion to get polarised is called its
3. Greater the polarisation, more is the neutralisation of charges and this decreases the ionic
character and increases the covalent character.
4. The polarising power of the cation and polarisability of the anion and hence, the formation of
the covalent bond is favoured by the following (
10
the N atom has an ionic pair of electrons.
is short of an electron pair to complete its octet.
donor and H+
acts as an acceptor. A coordinate bond is formed
Covalent character of an ionic bond
Although the ionic bond is considered to be 100% ionic, actually it has some covalent
character, just as a covalent bond has some ionic character.
When a cation approaches an anion, the electron cloud of the anion is attracted
towards the cation and hence, gets distorted. This effect is called
anion. As shown in the Figure 1 below.
2. The power of the cation to polarise the anion is called its polarity power
the anion to get polarised is called its polarisability.
3. Greater the polarisation, more is the neutralisation of charges and this decreases the ionic
character and increases the covalent character.
ation and polarisability of the anion and hence, the formation of
covalent bond is favoured by the following (Fajan's rules):
A coordinate bond is formed
Although the ionic bond is considered to be 100% ionic, actually it has some covalent
ion is attracted
towards the cation and hence, gets distorted. This effect is called polarisation of the
and the tendency of
3. Greater the polarisation, more is the neutralisation of charges and this decreases the ionic
ation and polarisability of the anion and hence, the formation of

11. i) Small size of the cation: Smaller the cation, greater is the polarising power. For this
reason, LiCl is more covalent than
ii) Large size of the anion: Larger the anion, greater is its polarisability. This explains the
covalent character of Sodium halides NaI > NaBr > NaCl > Na
iii) Large charge on the cation and anion:
polarising power. So, the covalent character of chlorides is as follows: NaCl < MgCl
AlCl3. Similarly, greater the charge on the anion, the
iv) d orbital in valence shell:
orbital in valence shell. For example
 Intermolecular forces:
1. Dipole dipole bond
2. Hydrogen bond
3. Van Der Waals force
1. Dipole: The molecule consisting partial positive and negative end, is called dipole.
Dipole-dipole forces are attractive
molecule and the negative end of another polar molecule
2. Hydrogen Bond
In a molecule, whenever the hydrogen atom is directly linked to
(like F, O or N), the shared pair of
a slight negative charge ( ) and the hydrogen atom acquires a slight
There occurs a weak bond formation between the negative end of one molecule and the
end of the other. This bond is known as
lines. It involves dipole-dipole interactions.
For example:
11
Smaller the cation, greater is the polarising power. For this
reason, LiCl is more covalent than KCl.
Larger the anion, greater is its polarisability. This explains the
Sodium halides NaI > NaBr > NaCl > NaF.
iii) Large charge on the cation and anion: Larger the charge on the cation, greater is its
covalent character of chlorides is as follows: NaCl < MgCl
. Similarly, greater the charge on the anion, the more easily it gets polarised.
d orbital in valence shell: In between two cations polarisibility is more which
orbital in valence shell. For example CuCl>NaCl
Dipole: The molecule consisting partial positive and negative end, is called dipole.
forces are attractive forces between the positive end of one polar
molecule and the negative end of another polar molecule
hydrogen atom is directly linked to a highly electronegative atom
(like F, O or N), the shared pair of electron is pulled more towards this atom and
) and the hydrogen atom acquires a slight positive (
There occurs a weak bond formation between the negative end of one molecule and the
This bond is known as a hydrogen bond, which is represented by dotted
dipole interactions.
Smaller the cation, greater is the polarising power. For this
Larger the anion, greater is its polarisability. This explains the
Larger the charge on the cation, greater is its
covalent character of chlorides is as follows: NaCl < MgCl2<
more easily it gets polarised.
In between two cations polarisibility is more which has d
Dipole: The molecule consisting partial positive and negative end, is called dipole.
forces between the positive end of one polar
a highly electronegative atom
electron is pulled more towards this atom and it acquires
positive ( ) charge.
There occurs a weak bond formation between the negative end of one molecule and the positive
which is represented by dotted

12. 12
Strength of the Hydrogen Bond
The hydrogen bond exerts a strong intermolecular attractive force. Its strength lies in between van
der Waals forces of attraction and the covalent bond.
Covalent bond Hydrogen bond Van Der Waal forces of attraction
For the intermolecular H bond
i) Association: Molecules of carboxylic acids exist as dimers because of hydrogen bonding. For
example, acetic acid exists as a dimer because the electropositive H atom of one molecule forms
an H bond with the carboxyl electronegative oxygen atom of the other.
(ii) High melting and boiling points:
Existence of H2O as a liquid, whereas the H2S.
At room temperature H2O is liquid but H2S is gaseous. In H2O molecule the difference of
electonegativity between H and O is more which generates polarity. Water molecules combine
themselves such a way where hydrogen atom linked them. This is called hydrogen bond. Due to
present hydrogen bond H2O is liquid but H2S unable to produce hydrogen bond so that it is
gaseous.
b) NH3 has a higher boiling point than PH3.

13. c) Ethanol has a higher boiling point
hydrogen bond and ether does not.
(iii) Solubility: The compounds that can form hydrogen bonds with the covalent molecules are
soluble in such solvents, for example, lower alcohols are soluble in water due to the
of an intermolecular H bond between water and alcohol molecules.
(iv) Lower density of ice as compared to
Solid ice has a cage-like structure, in which each water molecule is tetrahedrally linked to four
water molecules as shown below.
low. Thus ice floats in water.
3. Van Der Waals Force:
The force of attraction between two weak dipoles is called Van Der Waals foce like
London dispersion force.
When a temporary dipole approach a non polar covalent molecule then non polar
become slightly polar and produces intermolecular force which is called London dispersion
force.
13
) Ethanol has a higher boiling point than diethyl ether, as ethyl alcohol has
hydrogen bond and ether does not.
that can form hydrogen bonds with the covalent molecules are
soluble in such solvents, for example, lower alcohols are soluble in water due to the
an intermolecular H bond between water and alcohol molecules.
as compared to water:
like structure, in which each water molecule is tetrahedrally linked to four
As a result vacant spaces created in water so density becomes
The force of attraction between two weak dipoles is called Van Der Waals foce like
London dispersion force.
dipole approach a non polar covalent molecule then non polar
become slightly polar and produces intermolecular force which is called London dispersion
l alcohol has an intermolecular
that can form hydrogen bonds with the covalent molecules are
soluble in such solvents, for example, lower alcohols are soluble in water due to the formation
like structure, in which each water molecule is tetrahedrally linked to four
As a result vacant spaces created in water so density becomes
The force of attraction between two weak dipoles is called Van Der Waals foce like
dipole approach a non polar covalent molecule then non polar
become slightly polar and produces intermolecular force which is called London dispersion