1. Syllabus
5
Chemical Bonding and
Molecular Structure
Kossel - Lewis approach to chemical bond formation, concept of ionic
and covalent bonds. Ionic Bonding : Formation of ionic bonds, factors
affecting the formation of ionic bonds; calculation of lattice enthalpy.
Covalent Bonding : Concept of electronegativity, Fajan’s rule, dipole
moment; Valence Shell Electron Pair Repulsion (VSEPR) theory and shapes
of simple molecules. Quantum mechanical approach to covalent
bonding: Valence bond theory - Its important features, concept of
hybridization involving s, p and d orbitals; Resonance. Molecular Orbital
Theory - Its important features, LCAOs, types of molecular orbitals (bonding,
a ntibonding), sigma a nd pi-bonds, molec ula r orbita l elec tronic
configurations of homonuclear diatomic molecules, concept of bond
order, bond length and bond energy. Elementary idea of metallic
bonding. Hydrogen bonding and its applications.
TYPES OF BONDS
1. Ionic Bond
2. Covalent Bond
3. Co-ordinate Bond
4. Hydrogen Bond
1. Ionic Bond
An ionic bond is formed by complete transfer of one or more electrons from the
valency shell of one atom to the valency shell of another atom. In this way both
the atoms acquire the stable electronic configurations of the noble gases. The
atom losing the electrons becomes a positive ion and the atom which gains
electron becomes negative ion.
The essential process which occurs in the formation of NaCl (s) from Na(s) and
Cl2(g) can be expressed in the form of a BORN HABER CYCLE
C H A P T E R
CHAPTER
INCLUDES
 Types of Bonds
 Ionic Bond
 Covalent Bond
 Co-ordinate or
Dative bond
 Hydrogen Bonding
 Fajan’s Rule
 Formal Charge
 Dipole Moment
(a) Na (s) Na (g), H sublimation = +108.5 kJ/mole  VSEPR Theory
 Hybridisation and
(b)(b)
1
Cl
2 2 (g) Cl (g),
1
H
2
BE
Cl2 =
1
× 243 = +121.5 kJ/mole
2
Shapes of
Molecules
(c) Na (g) Na+
(g) + e–
, HIonization Energy = + 495.2 kJ/mole  Molecular Orbital
(d) Cl (g) + e– Cl– (g), H Electron Affinity = – 348.3 kJ/mole Theory
(e) Na+
(g) + Cl–
(g) NaCl(s), HLattice Energy = – 758.7 kJ/mole
Net Reaction :
Na (s) +
1
Cl
2 2 (g) NaCl (s)

2. JEE main Chemical Bonding and Molecular Structure
S S
+
+ –
+ –
Antibonding – + – +
The net enthalpy change
= 108.5 + 121.5 + 495.2 + (– 348.3) + (– 758.7)
= – 381.8 kJ/mole
In this reactions the S is almost negligible hence feasibility which depends on G could be measured in terms
of H and hence H is – ve, G = –ve. So reaction is feasible.
2. Covalent Bond
G. N. Lewis was the first to suggest in 1916 that atoms may combine with one another by sharing of electrons
in their valency shell to get a noble gas configuration.
As discussed in ionic bond, there is decrease in energy in the formation of an ionic bond. As a matter of fact, the
formation of a chemical bond of any type is possible only if the approach of the atoms towards one another is
accompanied by decrease in energy.
Types of Covalent Bond
s – s overlap (Sigma Overlap)
+
1s 1s
Bonding
Antibonding
s1s
s*
1s
p – p (Sigma Overlap)
–
+
2pz
2pz
Bonding
2pz
*
2pz
p – p Overlap Lateral ( - overlap)
Bonding
2py 2py
2py
–
2py
Antibonding
2py
*2py
–
1s
(Sigma Overlap)
y1s
p – p overlap
+ –
– +
2pz
– + –
– +
2pz
+
–

–
–
+

3. Chemical Bonding and Molecular Structure JEE main
4 4
3. Coordinate or Dative Bond
It was proposed by Sidgwick. In this type of combination both the electrons needed for sharing are contributed
only by one atom. The atom which contributes the pair of electrons (lone pair) is known as donor and the atom
which accepts these electrons is called acceptor. The coordinate bond is usually represented by an arrow
pointing towards the acceptor. e.g.,
H F
H N. B F
H F
Coordinate bond is found in the compounds like SO2, SO3, O3, NH +
, H3O+
, NH4Cl, SO 2–
and H2SO4 etc.
4. Hydrogen Bonding
It is the force of attraction that exists between the hydrogen atom covalently bonded to highly electronegative
atom (N or F or O) in a molecule and the electronegative atom of the same or neighbouring molecule. the
bond is represented by a dotted line as shown below.
..... H – F ..... H – F ..... H – F .....
The hydrogen bond is of two types, intermolecular (formed between H atom of one molecule with
electronegative atom of neighbouring molecule) Intramolecular H-Bonding (H atom and electronegative atom
of the same molecule).
(a) Intermolecular hydrogen bonding decreases the volatility and increases the viscosity and surface tension
of a substance.
(b) Order of strength of hydrogen bonding
H ......F > H...... O > H.......N
Energy 10 kcal/mole 7 kcal/mole 2.0 kcal/mole
(c) Because of hydrogen bonding, water (H2O) has higher boiling point than that of H2S.
(d) Because of hydrogen bonding the boiling point of HF is higher than that of HCl.
(e) Because of hydrogen bonding, ethanol (C2H5OH) has higher boiling point than that of dimethyl ether or
methoxy methane (which does not involve hydrogen bonding). Further alcohols are soluble in water since
they form hydrogen bonds with water molecules.
(f) Ice has lower density than water due to the formation of open cage like structure because of formation
of hydrogen bonds in water (H2O) molecules.
FAJAN’S RULE
Covalent character in ionic bonds
When oppositely charged ions approach each other, there is not only the attraction between the positively
charged cation and the negatively charged anion but also simultaneous repulsion between their nuclei. Thus
there is distortion, or deformation or polarization of anions. The electronic charge of anion does not remain
spherical but gots distorted.
No polarization Polarization of anion
+ –
+ –

4. JEE main Chemical Bonding and Molecular Structure
2  2  2  cos
1 2 1 2
(1) Polarizing power of cation
(i) High charge on cation
(ii) Small size of cation
Both of these make the polarizing power of cation high
(2) Tendency of an anion to get polarized
(i) High charge on anion
(ii) Large size of anion
Both of these factors make the anion get easily polarized
This polarization of ion results in a electric charge concentration between the two nuclei resulting in a
covalent bond with a large degree of charge separation.
The extent of polarization depends upon
(1) Polarizing power of cations.
(2) Tendency of an anion to get polarized (polarizability)
FORMAL CHARGE
Formal charge
=
on atom
Total number of
valence Electrons –
in the free atom
Total number of non
bonding (lone pair
electrons)
Total number of
→ bonding shared
electrons
Formal charge can help us to predict the most reasonable arrangement of atoms and write the Lewis structure.
In case of polyatomic ions, the net charge is possessed by the ion as a whole and not by individual atom.
The formal charge do not indicate real charge separation within the molecule. Formal charge help in the selection
of the lowest energy structure from a number of possible Lewis structures.
Example :
Formal charge of Br and each Cl in BrCl3 is
Cl – Br – Cl
|
Cl
(1) Formal charge of Br
FC = 7 – 7 = 0
(2) Formal charge of each Cl
FC = 7 – 7 = 0
DIPOLE MOMENT (μ)
Dipole moment is product of the magnitude of charge developed on any of the atom and the distance between
the atoms. ( = q × d), where q = charge developed and d = distance between two atoms. The unit of dipole
moment is debye (D). 1 debye = 1 × 10–18
esu cm.
(a) Dipole moment is the vector quantity.
(b) Molecule with more than one bond will have more than one dipole moment. The resultant dipole moment
of the molecule is given as res.  .
(c) Dipole moment predict whether a molecule is polar or non-polar.
(d) Dipole moment is zero for symmetrical and planar species.

5. Chemical Bonding and Molecular Structure JEE main
..
(e) Dipole moment helps to determine percentage ionic character can be calculated.
% ionic character 
Observed dipole moment
Calculated dipole moment
 100
Dipole moment for some molecules

 

F  = 0 (2) O  C  O  = 0
(3)(3)
N
H
H H
F
F F
 = 1.46D (4)
F
 = 0.25D
F
S
(5)
F F
F
 = 0
(6) O  = 0
(7) H2  = 0 (8)
(9) CCl4  = 0 (10)
H F
Cl
Cl
VALENCE SHELL ELECTRON PAIR REPULSION THEORY (VSEPR)
(a) Anumberof physical andchemicalpropertiesof moleculesareaffectedbythegeometryandshapeof molecules.
(b) VSEPR theory gives the information of shape of molecules.
(c) The arrangement of bonded pair electron and lone pair electron is done to minimize the repulsion i.e.,
Lone pair – lone pair > lone pair – bond pair > bond pair – bond pair.
(d) Theserepulsion - effects result in deviations from idealized shapeand alterations in bond anglesin molecules.
HYBRIDIZATION & SHAPES OF MOLECULES
S.No. Type of
hybridisation
No. of
hybridised orbitals
Bond pair Lone
Pair
Geometry Shape Examples
1. sp 2 2 – Linear Linear CO2
3 – Trigonal Planar Trigonal Planar BF3
2. sp
2
3 2 1 Trigonal Planar V-shape SO2
4 – Tetrahedral Tetrahedral CH4
3. sp
3
4
3
2
1
2
Tetrahedral
Tetrahedral
Pyramidal
V-shape
NH3
H2O
 = 1.91
 = 0
F
(1) B
F
N
F
..
2

6. JEE main Chemical Bonding and Molecular Structure
5 – Trigonal Trigonal PCl5
bipyramidal bipyramidal
4 1 Trigonal See-saw SF4
4. sp
3
d 5
3 2
bipyramidal
Trigonal T-shape ClF3
bipyramidal
2 3 Trigonal Linear ICl
–
2
bipyramidal
6 – Square Square (octahedral) SF6
bipyramidal bipyramidal
5. sp
3
d
2
6 5 1 Square Square IF5
bipyramidal pyramidal
4 2 Square Square XeF4
bipyramidal Planar
7 – Pentagonal Pentagonal IF7
6. sp
3
d
3
7
6 1
bipyramidal
Pentagonal
bipyramidal
Distored Square XeF6
MOLECULAR ORBITAL THEORY
(a) Linear combination of atomic orbitals (LCAO) method is used for the formation of molecular orbitals.
(b) For two atomic orbitals A and B, whose wave function is given as A and B. Then molecular
orbitals (MO) are given as MO = A ± B.
(c) The molecular orbital  formed by the addition of atomic orbitals is called bonding molecular orbitals. And
molecular orbital * formed by the subtraction of atomic orbitals is called antibonding molecular orbital.
(d) Molecular orbital configuration for diatomic molecules with less than 14 electrons.
 1s *1s 2s *2s {2px, 2py} 2pz {*2px, *2py} *2pz
(e) Molecular orbital configuration for diatomic molecules with more than 14 electrons.
 1s *2s 2s *2s 2pz {2px, 2py} {*2px, *2py} *2pz
(f) Molecular orbitals are arranged with increasing order of energy level.
(g) 2px and 2py similarly *2px and *2py have same energy and follow Hund’s rule.
(h) Bond order =
1
[Bonding electrons – Antibonding electrons].
2
(i) Bond order  Bond strength 

e.g., N2 = 14 electrons
1
Bond length
.
 Molecular orbital configuration
1s2
*1s2
2s2
. *2s2 2p2
, 2p2
2p2
 * 2p ,  * 2p  *2p
B.O =
x y z
1
[10 – 4] = 3
2
x y z
As no unpaired electron hence N2 is diamagnetic.
  