pptx

1. Prepared By : 1-Dler Ismail Rasul 2-Saz Zirak Hamad
3-Nawa Masud Rasul 4-Nihad Khalil Khalid

2. CONTENT
• History
• Valence Bond Theory
• Applications of Valence
• σ-Bond
• π—Bond
• Comparison
• Illustration
• Hybridization
• Types of Hybridization
• sp,sp2,sp3

3. Valence bond theory came into account when
G.N Lewis proposed the geometrical structures of compounds
in 1916. He explained how the valence electrons are arranged
among the atoms in a molecule. Using Lewis theory, people
started drawing Lewis structures that allowed them to predict
many properties of molecules such as molecular shape and
polarity.
Later, in 1927 Hitler and London attempted to explain the
bonding properties by applying quantum mechanics. They took
the hydrogen atoms as an example to explain chemical bond
formation using Schrodinger's wave equation.
History

4. The valence bond theory was proposed by Heitler and
London to explain the formation of covalent bond
quantitatively using quantum mechanics.
Later on, Linus Pauling improved this theory by
introducing the concept of hybridization.
Valence bond (VB) theory assumes that all bonds are
localized bonds formed between two atoms by the
donation of an electron from each atom.

5. Valence Bond theory describes covalent bond
formation as well as the electronic structure
of molecules.
The theory assumes that electrons occupy atomic
orbital's
that
the
of individual atoms within a
molecule, and electrons of one atom are
attracted to the
nucleus of another atom.

6. 1. Valence bond theory is used to explain the formation of
covalent bonds between two atoms
2. They explained the overlap of atomic orbitals
3. This theory explained the structure of molecules based
on the hybridisation phenomenon
4. It gave the idea that maximum intersection gives rise to
the development of strong possible bonds
Applications of Valence

7. COMPARISO
N
σ- bond
This bond is formed due
to the overlap of pure
s- s;s-p;p-p (or) hybrid
orbitals of two atoms
along their internuclear
axis.
It is a strongest bond
because the extent of
overlapping of orbitals
in sigma bond is greater.
Electeron density of a
sigma bond is
symmetrical about the
line joining the two
nuclei.
π-bond
This bond is formed
due to lateral or side
wise or parallel
overlapping of pure
‘p’
orbitals of two atoms.
It is weaker than sigma
bond because the extent
of overlapping of
orbitals in pi bond is
lesser.
Electron density of pi
bond is
unsymmetrical.

8. COMPARISO
N
A sigma bond
can present
alone.
In sigma bond
free rotation of
atom is possible.
A sigma bond
possesses high
bond energy.
A sigma bond is
less reactive.
A pi bond is
always formed
in addition to
sigma bond.
In pi bond free
rotation is not
possible.
A pi bond
possesses low
bond energy.
A pi bond is
more reactive.

9. COMPARISO
N
A sigma bond
has greater
bond length.
Compound containing
sigma bond generally
undergo substitution
reactions.
A sigma bond
influence the
geometry of
molecule.
Examples:
CH₄,H₂,Cl₂
A pi bond has
lesser bond
length.
Compound
containing pi bond
usually undergo
addition reactions.
A pi bond generally
has no effect on
geometry.
Examples:
CH₂=CH₂,N≡N,O=O

10. SIGMA BOND
σs-s bond:

11. SIGMA BOND
σp-p bond:

12. SIGMA BOND
σs-p bond:

14. The electronic configuration of hydrogen atom in the
ground state is 1s1.
In the formation of hydrogen molecule, two half filled
1s orbital's of hydrogen atoms overlap along the inter-
nuclear axis and thus by forming a σs-s bond.

15. The electronic configuration of O in the ground state is
[He] 2s2 2px 2 2py 1 2pz 1 .
The half filled 2py orbital's of two oxygen atoms
overlap along the inter-nuclear axis and form σp-p
bond.
The remaining half filled 2pz orbital's overlap laterally
to form a πp-p bond.

16. Thus a double bond (one σp-p and one πp-p ) is
formed between two oxygen atoms.

17. The ground state electronic configuration of N is [He]
x y z
2s2 2p 1 2p 1 2p 1.
Aσp-p bond is formed between two nitrogen atoms due
to overlapping of half filled 2px atomic orbital's along
the inter-nuclear axis.

18. The remaining half filled 2py and 2pz orbital's form
two πp-p bonds due to lateral overlapping. Thus a triple
bond (one and two) is formed between two nitrogen
atoms.

20. The intermixing of two or more pure atomic orbital's
of an atom with almost same energy to give same
number of identical and degenerate new type of
orbital's is known as hybridization.
The new orbital's formed are also known as hybrid
orbital's.
During hybridization, the atomic orbitals with
different characteristics are mixed with each other.

21. sp
sp2
sp3

22. Intermixing of one 's' and one 'p'
orbital's of almost equal energy to
give two identical and degenerate
hybrid orbital's is called 'sp'
hybridization.
These sp-hybrid orbital's are arranged
linearly at by making 180 ⁰ of angle.
They possess 50% 's' and 50% 'p'
character.

23. For example:

24. Intermixing of one 's' and two 'p'
orbital's of almost equal energy to
give three identical and degenerate
hybrid orbital's is known as 'sp2'
hybridization.
The three sp2 hybrid orbital's are
oriented in trigonal planar
symmetry at angles of 120 ⁰ to each
other.
The sp2 hybrid orbital's have 33.3%
's' and 66.6% 'p' character.

25. For example:

26. In
sp3
hybridization, one 's'
and
'p' orbital's of almost
equal
three
energy intermix to give
four
identical and degenerate hybrid
orbital's.
These four sp3 hybrid orbital's are
oriented in tetrahedral symmetry
with 109 ⁰ 28' angle with each
other.
The sp3 hybrid orbital's have 25% ‘s’
and 75% 'p' character.

27. For example: