Valence Shell Electron Pair Repulsion Theory Notes

Postulates of VSEPR Theory

1.        The geometry of the molecule is determined by no electron pairs ( bond and lone pairs ) present in the valence shell of center atom

2.        The electron pairs tends to stay as far as possible from each other for minimum repulsion and maximum stability.

3.       The molecule having only a bond pair of electrons in the central atom have regular geometry

No of bond pair	Geometry	Bond angle	Example
2	Linear	1800	BeF2, BeCl2
3	Trigonal pair	1200	BF3, BCl3
4	Tetrahedral	109028’	CH4, CCl4

4.       If the molecules have both bond pair and lone pair electrons, they have distorted structure

No of e- pair	Lone pair	Bond pair	Geometry	Bond angle	Example
4	0	4	Tetrahedral	109028’	CH4
4	1	3	Tetrahedral	107048’	NH3
4	2	2	Tetrahedral	104028’	H2O

5.       The extinct of repulsion is in the order of,

                  Lp – Lp > Lp – Bp > Bp – Bp

Valence Bond Theory ( VBT)

    According to this theory, a covalent bond is formed by overlapping atomic orbitals having unpaired electrons.

Postulates of valence bond theory

1.        Half-filled atomic orbital of one atom overlaps with a half-filled atomic orbital of another atom to form a covalent bond

2.       Atomic orbitals undergoing overlap must be sufficiently close to each other with proper alignment

3.       The strength of the bond formed depends upon the extent of overlapping of atomic orbitals. The greater the overlapping of atomic orbitals, the stronger the covalent bond formed.

4.       Overlapping lowers the energy of the molecule and excess energy is released. The energy released per mole is called stabilization energy / bond energy.

5.       Number of unpaired electrons in an atom can increase at the time of reaction due to the excitation of electrons from one orbital to the orbital of slightly higher energy.

6.       Number of unpaired electrons in the ground state or excited state of an atom is called the covalency of the element.

Types of covalent bond

Sigma bond	Pie bond
It is formed by end to end overlapping of half-filled orbitals.	It is formed by sideways overlapping of half-filled orbitals.
Overlapping of orbitals takes place among the internuclear axis	Overlapping of orbitals takes place perpendicular to the internuclear axis
The extent of overlapping is large so the hand formed is strong.	The extend of overlapping  is small. So the bond formed is weak.
It may be present alone or along with pie-bonds.	It cannot be formed alone. It is formed when sigma bond is present.

 

 

Hybridization

           The process of mixing of atomic orbital of the same atom of different energy to obtain a new sheet of atomic orbital having equivalent energy and equal in a number of mixing orbitals is known as hybridization. Thus, obtain orbital are called hybrid orbitals.

Types of hybridization

1.       Sp³ hybridization

      The process of mixing of one ‘s’ and three ‘p’ orbitals is known as sp³ hybridization.

Formation of CH₄:

     There are four sp³ hybrid orbitals in carbon atoms where four hydrogen atoms overlap to form a sigma bond. Thus, the structure of methane is tetrahedral with a bond angle of 109⁰28’.

2.       Sp² hybridization

       The process of mixing one ‘s’ and two ‘p’ orbitals is called sp² hybridization.

 Formation of C₂H_­4:

The doubled bonded carbon in ethene is sp² hybridized. So, it forms a trigonal structure. The bond angle is 120⁰ and the carbon-carbon bond length is 1.34A⁰. The C-H bond length is 1.09A⁰.

3.       Sp hybridization

    The process of mixing one ‘s’ and one ‘p’ orbitals is called sp hybridization.

 Formation of C₂H₂:

  There are four unpaired electrons on carbon atoms and the other two in unhybrid P_y and P_z­ orbital. The carbon formed two sigma bonds with one Hydrogen and other carbon atoms. Its bond angle is 180^0.

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