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CHAPTER

THEORIES OF CHEMICAL BONDING

Contents
❖ Valence Bond Theory with applications
❖ Molecular Orbital Theory with applications
❖ Valence Shell Electron Pair Repulsion Theory with applications
❖ Sigma and Pi bonds
❖ Hybridization
❖ Sp3 Hybridization
❖ Sp2 Hybridization
❖ Sp Hybridization
❖ Dsp3 Hybridization
❖ D2sp3 Hybridization

Saeed Anwar Lecturer in Chemistry
Government College Peshawar Pakistan
E-mail: saeedchemistry28@gmail
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VALENCE BOND THEORY
Valence bond theory was developed by W
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The main
postulates of valence bond theory are;
1
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2
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3
...

4
...

5
...

6
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Applications of valence bond theory
Valence bond theory (VBT) is also called ‘’electron pair theory’’ because a half-filled
atomic orbital both having opposite spins form a single covalent bond
...
The electrons which are not taking place in bond formation are
called ‘lone pair’ and remains inert
...
H2 Molecule
Hydrogen molecule (H2) is formed when 1s1 atomic orbital of one hydrogen atom and 1s1
atomic orbital of another hydrogen atom combines
...
Diagram is represented
as;

2
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When both half-filled atomic orbitals combine,
forming single covalent bond which is called sigma bond
...


3
...
The
electronic configuration of oxygen atom is: 1s2, 2s2, 2px1, 2py1, 2pz2
...
N2 Molecule
Nitrogen molecule (N2) is formed by combination of three half-filled atomic orbitals of
each nitrogen atom having three unpaired electrons each
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Nitrogen molecule having triple bond
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The other two bonds are called pi bonds (π) which is formed by 2py1 and
2py1 half-filled atomic orbitals of two nitrogen atoms
...
H2O Molecule
H2O molecule is formed by combination of tow half-filled atomic orbitals of oxygen atoms and
two half-filled atomic orbitals of hydrogen atoms
...
The
electronic configuration of oxygen atom is: 1s2, 2s2, 2px1, 2py1, 2pz2
...
HCl Molecule
HCl molecule is formed by combination of one half filled atomic orbital of Cl (3px1) and
one half filled atomic orbital of H (1s1) to form single sigma covalent bond
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Limitations of Valence bond theory (VBT)
1
...

2
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3
...

4
...
e
...
, (NO)
...
This theory cannot explain the bonding in electron deficient molecules
...
This theory cannot explain the non-existence of noble gases
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All
the sigma bonds have axial symmetry
...
Sigma bonds are stronger than pi
bonds
...

Sigma bond is formed by overlapping of s_s, s_px, s_py, s_pz, px_px atomic orbitals
...
Each 1s1 orbital of hydrogen
atom overlaps to form sigma bond in hydrogen molecule as shown below
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Fig: s__ p orbital overlap of HCl molecule
p___ p overlapping
Cl2 molecule is formed by combining one half filled atomic orbital of one Cl atom(3px1) with
one half filled atomic orbital of another Cl atom(3px1) containing p__ p overlapping and forms
sigma bond as shown below
...
Pi(ᴫ) bond
is weaker than sigma bond(σ)
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Nitrogen molecule (N2) is formed by combination of three half-filled atomic orbitals of
each nitrogen atom having three unpaired electrons each
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Nitrogen molecule having triple bond
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The other two bonds are called pi bonds (π) which is formed by 2py1 and
2py1 half-filled atomic orbitals of two nitrogen atoms
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MOLECULAR ORBITAL THEORY
Molecular orbital theory (MOT) was developed by Hund and Milliken in 1927 and extended
later by Lennard-Jones in 1929
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The atomic orbitals overlap to form molecular orbitals and the number of atomic
orbitals is always equal to the number of molecular orbitals
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Half molecular orbitals are lower in energy called bonding molecular orbitals
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3
...


4
...


5
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The no
...
of bonding and anti-bonding electrons
...
of electrons in bonding orbital - no
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Hydrogen molecule (H2)
Hydrogen molecule is formed by the overlap of two 1s atomic orbitals of two
hydrogen atoms
...
Electrons
are filled in these orbitals according to Aufbau principle
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The electronic configuration of H2 is represented as;
H(1s1) + H(1s1)

H2[σ(1s2), σ*(1s0)]

Bond order = 2-0 = 1, it means H2 molecule contain single covalent bond e
...
,H__H
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Helium molecule(He2)
Helium (He) has 1s2 electronic configuration
...

Two electrons go to the bonding M
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[σ(1s2)] and two electrons go to the anti
bonding M
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[σ*(1s2)]
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2
Diagram

3
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Oxygen (O2)
molecule is paramagnetic due to the presence of unpaired electrons in anti-bonding
ᴫ*2py1 and ᴫ*2pz1 orbitals
...

Bond order = 8-4/2 =4/2=2,it means Oxygen(O2) molecule contain double bond
...
g
...
Nitrogen(N2) molecule
Nitrogen (N) has electronic configuration 1s2, 2s2, 2px1, 2py1, 2pz1
...
Nitrogen (N2)
molecule has the following configuration which can be represented as;

N (1s2, 2s2, 2px1, 2py1, 2pz1) + N(1s2, 2s2, 2px1, 2py1, 2pz1)
N2[ (C2s2), (σ*2s2), σ2px2, ᴫ2py2, ᴫ2pz2, σ* 2px0, ᴫ*2py0, ᴫ*2pz0 ]
Bond order = 8-2/2 = 6/2 = 3, it means that Nitrogen (N2) molecule contain triple bond
...
g
...
Lithium (Li2) molecule
Lithium (Li) has electronic configuration 1s2, 2s1
...
One is called bonding molecular
orbital (σ2s) and the other is called anti bonding molecular orbital (σ*2s)
...

Chemical equation can be represented as;

Li2 [(σ2s2), ( σ*2s0)]

Li (2s1) +Li( 2s1)
Bond order = 2-0/2 = 1

It means that Li2 contains single covalent bond
...
g
...

The diagram is represented as;

Fig: Molecular orbital diagram of Li2 molecule
6
...
When two 2s2 atomic orbitals of two
Be atoms combine to form two molecular orbitals called bonding molecular orbitals (σ2s2)
and anti-bonding molecular orbitals (σ*2s2)
...

Chemical equation can be represented as;
Be (2s2) + Be (2s2)
Bond order = 2-2/2 = 0
It means that Be2 contains no covalent bond
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Fluorine (F2) molecule
The electronic configuration of F is 1s2, 2s2, 2px1, 2py2, 2pz2
...
Chemical equation can be represented as;
F (2s2, 2px1, 2py2, 2pz2) + F (2s2, 2px1, 2py2, 2pz2)

F2[σ2s2,σ*2s2,σ2px2,π2py2,π2pz2,σ*2px0,π*2py2,π*2pz2]

Fig: Molecular orbital diagram of F2 molecule
Heteronuclear diatomic molecules
1
...

Chemical equation can be represented as;
N (2s2, 2px1, 2py1, 2pz1 ) + O(2s2, 2px1, 2py1, 2pz1 )
NO[σ2s2, σ*2s2, σ2px2, π2py2, π2pz2, σ*2px0, π*2py1, π*2pz0]

Bond order = 8-3/2 = 5/2 = 2
...
Since there is one unpaired electron in NO molecule, therefore NO
is paramagnetic according to molecular orbital theory
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Carbon monoxide (CO)
The electronic configuration of C & O is:
C :1s2, 2s2, 2px1, 2py1, 2pz0(4 valence electrons)
O : 1s2, 2s2, 2px1, 2py1, 2pz2 (6 valence electrons)
The total number of valence electrons is 10
...
One is sigma
and two are pi bonds
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Therefore CO is
diamagnetic according to molecular orbital theory
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CN molecule
The electronic configuration of C and N is:
C : 1s2, 2s2, 2px1, 2py1, 2pz0(4 valence electrons)
N : 1s2, 2s2, 2px1, 2py1, 2pz1 (5 valence electrons)

The total number of valence electrons is 9
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5, CN contains triple covalent bond
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There is one unpaired electron, therefore CN molecule is paramagnetic
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Main postulates of VSEPR
1
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2
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Lone pair causes more repulsion than bond pair
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3
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4
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5
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Applications of VSEPR Theory
Shapes of molecules containing two electron pairs (AB2):
Beryllium chloride (BeCl2) is a typical example of molecules containing two electron pairs
...
This shows linear geometry having bond
angle 1800 and sp hybridization
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Structure: linear

Bond angle: 1800
Hybridization: sp
Bond pair: 02
Lone pair: 00
Fig: linear structure of BeCl2
Typical examples of this category are;
BeCl2, MgCl2, CaCl2, SrCl2, BaCl2, BeH2, MgH2, CaH2, SrH2, BaH2 etc
...

This shows equilateral triangle geometry having bond angle 1200 and hybridization sp2
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Structure: Equilateral triangular
Bond angle: 1200
Hybridization: sp2
Bond pair: 03
Lone pair: 00
Fig: Equilateral triangular structure of BF3
Typical example of this category are;
BF3, BCl3, BBr3, BI3, AlF3, GaF3, InF3, TiF3, AlCl3, GaCl3, InCl3, TiCl3, etc
...
All the pairs are bond pairs
...
50 and hybridization sp3
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Structure: Tetrahedral
Bond angle: 109
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Shapes of molecules containing five electron pairs (AB5):
Phosphorous pentachloride (PCl5)
PCl5 having five bond pairs and no lone pair therefore, it has trigonal bipyramidal structure
having two types of bond angles (900, 1200,) and hybridization (dsp3)
...

The structure of PCl5 is represented as;

PCl5

PCl5
Structure: Trigonal bipyrimidal
Bond angle:900,1200
Hybridization:dsp3
Bond pair: 05
Lone pair: 00

AB5

Typical examples of this category are;
PF5, PCl5, PBr5, PI5, AsCl5, SbCl5, AsBr5, AsI5
...
Therefore, it has octahedral
structure having bond angle 900 and hybridization d2sp3
...
The structure of SF6 is represented as;

SF6

SF6

AB6

Structure: Octahedral
Bond angle: 900
Hybridization: d2sp3
Bond pair: 06
Lone pair: 00
Typical examples of this category are;
SF6, SeF6, TeF6, SCl6, SeCl6, TeCl6, SBr6
...
Therefore, it has pentagonal
bipyrimidal structure having bond angle 720 and 900
...
The structure of IF7 is represented as;

IF7

IF7

AB7

Structure: Pentagonal bipyrimidal
Bond angle: 720,900
Hybridization: d3sp3
Bond pair: 07
Lone pair: 00
Typical examples of this category are;
ClF7, BrF7, IF7
...
The lone pair
present on Sn pushes the bond pair due to which the bond angle decreases from 1200 to 1160
...


Structure: Angular
Bond angle: 1160
Hybridization: sp2
Bond pair: 02
Lone pair: 01
Fig: Angular structure of SnCl2
Typical examples of this category are;
CCl2, GeCl2, SnCl2, PbCl2 etc
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N forms three covalent bonds with three
hydrogen atoms by sharing three electrons and other two electrons remain inert and cannot
form bond in this structure
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50 to 107
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N has valence shell 1s2, 2s2, 2px1, 2py1,
2pz1
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50 and hybridization sp3
...
50
Hybridization: sp3

Bond pair: 03
Lone pair: 01
Fig: Trigonal pyramidal structure of NH3
Typical examples of this category are; NH3, PH3, AsH3, SbH3, BiH3
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O forms two covalent bonds with two
hydrogen atoms by sharing two electrons and four electrons remain inert and not take part in
bond formation and forms two lone pairs on the central atom oxygen (O)
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50 to 105
...
This shows
angular structure having bond angle 105
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Structure: angular
Bond angle: 105
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Limitations of VSEPR Theory
1
...
e
...
, Li2O
should have the same structure as H2O but in fact it is linear
...
This theory cannot explain the shapes of molecules having extensive delocalized pi
electron systems
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This theory cannot explain the shapes of certain transition metal complexes
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This theory cannot explain the shapes of certain molecules having inert pair of
electrons
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The number of atomic orbitals is equal to the number of hybrid orbitals
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Atomic orbitals may have same or
different size, shape and energy but hybrid orbitals have same size, shape and energy
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sp3 hybridization 2
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sp hybridization 4
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d2sp3 hybridization
1
...
e
...
, Methane (CH4)
Methane forms sp3 hybridization by sharing four atomic orbitals with four atomic orbitals of
hydrogen atom
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50 bond angle and tetrahedral structure
having 25 % s character and 75 % p character
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2
...
e
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, Ethene (2HC==CH2)
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The fourth orbital (py1) forms pi bond with other carbon atom(py1)
...
Diagram is shown as;

Fig: Trigonal panar structure of Ethene

H

H

C
2s

2px 2py

2pz

2s

2px 2py

2pz

C

H

H

3
...
e
...
, Ethyne (HC==CH)
...
2s1 orbital combine with 1s1 orbital of hydrogen
and 2px1 orbital of one C combine with 2px1 orbital of other C to form sigma bond and 2py1,
2pz1 orbital of C combine with 2py1 and 2pz1orbital of another C to form two pi bonds
...
In sp hybridization, 50%
s character and 50% p character
...
dsp3 hybridization
Definition: The process of mixing of one s, three p and one d atomic orbitals to form five
equivalent dsp3 hybrid orbitals is called dsp3 hybridization
...
g
...

Phosphorous (P) atom has electronic configuration 1s2, 2s2, 2p6, 3s2, 3p3, 3d0
...
These five
atomic orbitals of P overlaps with 5 atomic orbitals of Cl atoms forming dsp3 hybridization
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d2sp3 hybridization
Definition: The process of mixing of one s, three p and two d orbitals to form six equivalent
d2sp3 hybrid orbitals is called d2sp3 hybridization
...
g
...
One electron of 3s2 and one
electron of 3p4 promotes to 3d orbitals, after promoting becomes3s1, 3px1, 3py1, 3pz1, 3d2
...
The diagram is shown as;

Fig: Octahedral structure of SF6
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֎֎֎END֎֎֎



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