Valence bond theory is one of the main basic theory of chemistry to know about the geometry of molecules and bonding

1. Chapter 2
Theories of Covalent Bonding
Ilham

2. 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.
VALENCE BOND THEORY:

3. Valence Bond theory describes covalent bond formation
as well as the electronic structure of molecules.
The theory assumes that electrons occupy atomic orbital’s
of individual atoms within a molecule, and that the
electrons of one atom are attracted to the nucleus of
another atom.
A covalent bond forms when the orbitals of two atoms
overlap and the overlap region, which is between the
nuclei, is occupied by a pair of electrons.
(The two wave functions are in phase so the amplitude
increases between the nuclei.)
VALENCE BOND THEORY:

4. A covalent bond is formed by the overlapping of two half
filled valence atomic orbital's of two different atoms.
The electrons in the overlapping orbital's get paired and
confined between the nuclei of two atoms.
The electron density between two bonded atoms increases
due to overlapping. This confers stability to the molecule.
Greater the extent of overlapping, stronger is the bond
formed.
The direction of the covalent bond is along the region of
overlapping of the atomic orbital's i.e., covalent bond is
directional.
VALENCE BOND THEORY RULES:

5. There are two types of covalent bonds based on the pattern
of overlapping as follows:
σ-bond
A sigma bond (symbol: σ) is a covalent
bond formed via linear overlap of two
orbital’s.
π-bond
A pi bond (symbol: π) is a covalent bond
formed via parallel overlap of two orbital's.
Types of Overlap:

6. The covalent bond formed due to overlapping of
atomic orbital along the inter nucleus axis is called σ-
bond.
It is a stronger bond and cylindrically symmetrical.
Depending on the types of orbital's overlapping, the
σ-bond is divided into following types:
(i): σs-s bond, (ii): σp-p bond, (iii): σs-p bond:
σ-bond

7. (i): σs-s bond
σ-bond

8. (ii): σp-p bond
σ-bond

9. (ii): σp-p bond
σ-bond

10. The covalent bond formed by sidewise overlapping
of atomic orbital's is called π- bond.
In this bond, the electron density is present above and
below the inter nuclear axis.
It is relatively a weaker bond since the electrons are
not strongly attracted by the nuclei of bonding atoms.
π- bond

11. 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 internuclear axis and thus by forming a σs-s bond.
H2 Molecule

12. The electronic configuration of Cl atom in the ground
state is [Ne]3s2,3px
2,3py
2,3pz
1.
The two half filled 3pz atomic orbital's of two
chlorine atoms overlap along the inter-nuclear axis
and thus by forming a σp-p bond.
Cl2 Molecule

13. The electronic configuration of hydrogen atom in the
ground state is 1s1.
The electronic configuration of Cl atom in the ground
state is [Ne]3s2,3px
2,3py
2,3pz
1.
The half filled 1s orbital of hydrogen overlap with the
half filled 3pz atomic orbital of chlorine atom along the
internuclear axis to form a σs-p bond.
HCl Molecule

14. The electronic configuration of O in the ground state is: [He] 2s
2,
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.
Thus a double bond (one σp-p and one πp-p) is formed between two
oxygen atoms.
O2 Molecule

15. The electronic configuration of N in the ground state is: [He] 2s
2,
2px
1, 2py
1, 2pz
1.
A σp-p bond is formed between two nitrogen atoms due to
overlapping of half filled 2px atomic orbital's along the inter-nuclear
axis.
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.
N2 Molecule

16. The valence bond theory fails to explain the
tetravalency of carbon.
It fail to explain the paramagnetic behaviour of O2.
Structures of Xenon fluorides cannot be explained by
Valence Bond approach. According to the valence bond
approach, covalent bonds are formed by the
overlapping of the half-filled atomic orbital. But xenon
has a fully filled electronic configuration. Hence the
structure of xenon fluorides cannot be explained by
VBT.
N2 Molecule

17. We cannot use this direct overlap picture for CH4’s
bonding. The 2s and the three 2p orbitals on each C
do not fit into the CH4 molecule's 109o bond angles,
since the 2p orbitals are at 90° to each other
Valence Bond Theory states that HYBRID orbitals of
the outermost orbitals on an atom are formed from
the atoms’ atomic orbitals
Hybridization

18. 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 orbital's with
different characteristics are mixed with each other.
The number of hybrid orbitals obtained equals the
number of atomic orbitals mixed.
Hybridization

19. sp
sp2
sp3
sp3d
sp3d2
sp3d3
Types Hybridization

20. 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.
sp-Hybridization

21. The sp hybrid orbitals in gaseous BeCl2.
atomic
orbitals
on Be
hybrid
orbitals
orbital box diagrams

22. Figure 11.2 The sp hybrid orbitals in gaseous BeCl2(continued).
orbital box diagrams with orbital contours

23. 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’ character and
66.6% 'p' character.
sp2-Hybridization

24. The sp2 hybrid orbitals in BF3.
Note the three sigma
bonds formed between B
and each F.

25. In sp3 hybridization, one 's' and three 'p' orbital's of
almost equal 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’ character and
75% 'p' character.
sp3-Hybridization

26. The sp3 hybrid orbitals in CH4.

27. Figure 11.5 The sp3 hybrid orbitals in NH3.
You have to know how to draw
this energy hybrid formation.

28. Figure 11.5 continued The sp3 hybrid orbitals in H2O.
You have to know how to
draw this energy hybrid
formation.

29. In sp3d hybridization, one 's', three 'p'
and one 'd’ orbital's of almost equal
energy intermix to give five identical
and degenerate hybrid orbital's, which
are arranged in trigonal bipyramidal
symmetry.
Among them, three are arranged in
trigonal plane and the remaining two
orbital's are present above and below
the trigonal plane at right angles.
The sp3d hybrid orbital's have 20%
's', 60% 'p' and 20% 'd' characters.
sp3d-Hybridization

30. The sp3d hybrid orbitals in PCl5.

31. Intermixing of one 's', three 'p' and
two 'd' orbital’s of almost same
energy by giving six identical and
degenerate hybrid orbital's is called
sp3d2 hybridization.
These six sp3d2 orbital's are arranged
in octahedral symmetry by making 90
⁰ angles to each other.
This arrangement can be visualized
as four orbital’s arranged in a square
plane and the remaining two are
oriented above and below this plane
perpendicularly.
sp3d2-Hybridization

32. The sp3d2 hybrid orbitals in SF6.

33. In sp3d3 hybridization, one 's', three
'p' and three 'd' orbital's of almost
same energy intermix to give seven
sp3d3 hybrid orbital's, which are
oriented in pentagonal bipyramidal
symmetry.
Five among the sp3d3 orbital's are
arranged in a pentagonal plane by
making 72⁰ of angles.
The remaining are arranged
perpendicularly above and below this
pentagonal plane.
sp3d3-Hybridization

40. Figure 11.8
The conceptual steps from molecular formula to the
hybrid orbitals used in bonding.
Molecular
formula
Lewis
structure
Molecular
shape and e-
group
arrangement
Hybrid
orbitals
Figure 10.1
Step 1
Figure 10.12
Step 2 Step 3
Table 11.1

41. Figure 11.9 The s bonds in ethane(C2H6).
both C are sp3 hybridized s-sp3 overlaps to s bonds
sp3-sp3 overlap to form a s bond
relatively even
distribution of electron
density over all s
bonds

42. Figure 11.10 The s and p bonds in ethylene (C2H4).
overlap in one position - s
p overlap - p
electron density
Proper name is ethene.

43. Figure 11.11 The s and p bonds in acetylene (C2H2).
overlap in one position - s
p overlap - p