CHEM 1411 – CHAPTER - 9
COVALENT BONDING ORBITALS
VSEPR Model of covalent Molecules (Valence Shell electron Pair Repulsion Model)
This model explains the 3 – dimensional geometry of covalent molecules. The following are the
important assumptions of this theory.
1. The central atom in a covalent molecule is surrounded by a definite number of electron pairs.
These electron pairs include bond pair and lone pair of electrons.
2. The electron pairs arrange themselves at the maximum possible distance around the central
atom to minimize repulsion. This arrangement gives a definite geometry to the molecule.
3. The shared e- pairs are called the bond pairs and the unshared e- pairs are called lone pairs.
4. The repulsion between the e- pairs is in the order lp-lp > lp-bp > bp-bp.
Geometry of molecules depends on the number of electron pairs on the central atom.
2 pairs of e-
1800
Linear
3 pairs of e-
1200
Trigonal planar
4 pairs of e-
109.50
Tetrahedral
5 pairs if e-
1200 and 900
Trigonal bipyramidal
6 pairs if e-
900
Octahedral
To find the geometry of molecules, first draw the Lewis structure and then find the number of
electron pairs. Using the above table assign a suitable geometry. This model considers double and
triple bonds as if they were single bonds
 Molecules in which central atom has no lone pair
Ex. BeCl2, BF3, CH4, PCl5, SF6
1. Methane
H
C
H
H
H
109.50
Tetrahedral
2. BeCl2
Cl _____Be_____Cl
1800
Linear
 Molecules in which central atom has one or more lone pair
Ex. SO2, NH3, H2O
Dipole Moment (
Dipole moment of a covalent molecule is the product of the magnitude of charge and the distance
between the charges.
Q x r where ‘Q’ is the magnitude of charge and ‘r’ is the distance.
Dipole moment is exhibited by polar molecules It is expressed in Debye unit (D)
1 D = 3.36 x 10-30 C. m
Dipole moment is vector quantity; it has both magnitude and direction.
A polar molecule contains one or more polar bonds and a resultant dipole moment.
It is not always necessary that a molecule must be polar if it contains polar bonds. There are molecules
containing polar bonds but the molecule as a whole is non-polar. The overall polarity depends on the
geometry of the molecule.
The following rules are to be kept in mind to decide a molecule polar or non-polar
 Symmetric molecules containing same kind of atoms around the central atom are
non polar, even if they contain polar bonds.
Ex. BeF2 ,CO2, CCl4
BeF2 and CO2 molecules are non polar though they contain polar bonds because of their
linear geometry. Their dipole moments cancel each other. CCl4 is tetrahedral.
 Symmetric molecules containing different kind atoms around the central atom are
polar if they contain polar bonds.
Ex. CHCl3,
 Asymmetric molecules are polar if they contain polar bonds.
Ex. NH3, NF3, H2O, SO2 etc
Valence Bond theory
According valence bond theory, a covalent bond is formed by the overlapping of atomic orbitals of the two
atoms.
Hybridization
The hypothetical concept of mixing of atomic orbitals of approximately equivalent energies to form equal
number new identical orbitals is hybridization. The new orbitals formed are called hybrid orbitals.
The concept of hybridsation is given to explain the observed valency of some elements like the
tetravalency of carbon.
Hybridization
sp3
sp2
sp
sp3d
sp3d2
# of orbitals involved
1s and 3 p
1s and 2 p
1s and 1 p
1s, 3p and 1d
1s, 3p and 2d
# of e- pair
4
3
2
5
6
Shape
Tetrahedral
Trigonal planar
Linear
Trigonal bipyramidal
Octahedral
Example
CH4
BF3
BeH2
PCl5
SF6
To find the hybridization of the central atom,
 Draw the Lewis structure of the molecule
 Treat double and triple bonds as single bonds to find the number of electron pairs
 Determine the number of the electron pairs surrounding the central atom
 Find the hybridization and shape corresponding to the number of electron pairs
Hybridization of molecules containing double or triple bonds
Ex. Ethene (C2H4), Ethyne (C2H2)
Molecular Orbital Theory of Covalent Molecules
 The electrons in a covalent molecule are present in new orbital called the molecular orbital
 A molecular orbital is formed by the combination of two atomic orbital of nearly same energies
 Combination of 2 atomic orbital result in the formation of 2 new molecular orbital, one with
lower energy and the other with higher energy.
 The molecular orbital with lower energy is called bonding molecular orbital and the one with
higher energy is the anti bonding molecular orbital.
 Bonding and anti bonding molecular orbital are of two types, sigma molecular orbital where the
electron density is concentrated symmetrically around the line joining the two nuclei and
pi molecular orbital where the electron density is concentrated above and below the line joining
the two nuclei. A sigma molecular orbital is more stable compared to a pi molecular orbital
 Electrons are filled in the molecular orbital in accordance with Hund’s rule, Pauli’s exclusion
principle and Aufbau principle.
Bond Order = ½ [Number of electrons in Bonding molecular obital – Number of
electrons in anti bonding molecular orbital]
Bond order is directly proportional to the bond dissociation energy and inversely proportional to the
bond length. Higher bond order makes the molecule more stable.
For diatomic molecules, the following is the order of increasing energies of molecular orbital.
s < *s < s < *s < 2px < py = pz <*py =
pz < *px
For lighter diatomic molecules B2, C2 and N2 , the order is slightly different from above.
s < *s < s < *s < py = pz <2px <*py =pz < *px