How do bonds form?
Valence Bond Theory
Hybridization and VSEPR
Valence Bond Theory
Valence bond theory describes covalent bond
in terms of the overlap of atomic orbitals.
H2:
This type of end-to
end overlap of
orbitals produces
a sigma,σ bond
1s
1s
HF:
2p
1s
2p
1s
H
H
1s
Overlapping of the
1s orbitals
Other Points about Valence
Bond Theory
• In hydrogen fluoride the 1s orbital of the H will overlap
with the half-filled 2p orbital of the F forming a
covalent bond.
2p
+
+
F
1s
• The newly combined orbital will contain an electron
pair with opposite spin just like a filled atomic orbital.
2p
Example HF
H
Example: H2
Covalent Bond
H-H
F2 :
1s
• The valence bond model or atomic orbital
model was developed by Linus Pauling in
order to explain how atoms come together
and form molecules.
• The model theorizes that a covalent bond
forms when two orbitals overlap to produce a
new combined orbital containing two
electrons of opposite spin.
• This overlapping results in a decrease in the
energy of the atoms forming the bond.
• The shared electron pair is most likely to be
found in the space between the two nuclei of
the atoms forming the bonds.
Overlapping of the1s and 2p orbitals
+
• This theory can also be applied to
molecules with more than two atoms
such as water.
• Each covalent bond results in a new
combined orbital with two oppositely
spinning electrons.
• In order for atoms to bond according to
the valence bond model, the orbitals
must have an unpaired electron.
Covalent Bond
H-F

Bonding and Molecular Shape
• When two atoms like hydrogen come
together there is a precise distance between
the two orbitals that ensures maximum
overlap of the two orbitals
• The need for maximum overlap is responsible
for the different shapes of molecules found in
nature - VSEPR
Valence Bond Theory and
Molecular Geometry
1s
2p
1s
2s
2p
2s
2p
Orbital overlap
suggests that the bond
angle is 90, but we
know that the angle is
104.5, therefore there
must be a different
orbital configuration.
What’s a Hybrid?
• The distance between the two nuclei in a
bond is referred to as bond length
• The two shared electrons of opposite spins
spend most of their time between the two
nuclei
• Overlap of orbitals can be between like
orbitals (s and s) or unlike orbitals (s and p)
Hybrid Orbitals
Consider the H2O molecule:
Central atom O:
Bond length
• To make the
connection between
Quantum theory and
VSEPR shapes, we
theorize that the central
atom hybridizes the
available orbitals to
achieve the required
number of bonds and
the correct molecular
shape.
Hybrid Orbitals
• Hybrid orbitals are mixtures of s, p, and d
atomic orbitals with intermediate energies.
• The number of s, p, and d orbitals that have
combined, equals the number of hybrid orbitals.
s
+
p
sp
sp

sp2 Hybrid Orbitals in BF3
sp Hybrid Orbitals in BeF2
Linear geometry is achieved using two sp
hybrid orbitals.
Be atom:
Promotion:
Hybridization:
s
p
p
p
s
p
p
p
sp sp
p
p
Trigonal planar geometry is achieved using
three sp2 hybrid orbitals.
B atom:
Promotion:
sp
sp
Tetrahedral geometry is achieved using four
sp3 hybrid orbitals.
Promotion:
Hybridization:
s
p
p
p
s
p
p
p
sp3
sp3
sp3
sp3
p
p
s
p
p
p
sp2
sp2
sp2
sp2 sp2 sp2
p
Trigonal bipyramidal
geometry is achieved
using five sp3d hybrid
orbitals.
Octahedral geometry is
achieved using six sp3d2
hybrid orbitals.
sp3
sp3
sp3
p
sp3d and sp3d2 Hybridization
sp3 Hybrid Orbitals in CH4
C atom:
Hybridization:
s
sp3
Types of Bonds
Types of Bonds - σ
• There are 2 main types of covalent bonds
• End to end overlap of orbitals (s, p, d, f, or
hybrid) forms sigma, σ, bonds
• The first bond between the central atom and
a ligand is a sigma bond
– Single Bonds
• Areas of electron density are concentrated between
the nuclei of the bonding atoms (along the bond)
• Bond is extremely strong
– Double and triple bonds
• Areas of electron density are above and below the
plane of the molecule
• Bonds are highly reactive and easy to break

Covalent Bond Formation
A  (sigma) bond results from end-to-end
overlap of orbitals. The maximum electron
density lies along the bond.
Types of Bonds - π
• Double and triple bonds are pi, π, bonds
• A single bond is composed of 2 areas of electron
density above and below the sigma bond
• extremely reactive because they are so far away from
the influence of the nucleus and due to their location
• π bonds are formed from regular p orbitals
π bond
Covalent Bond Formation
Hybridization in Carbon
A  (pi) bond results from side-to-side overlap
of p orbitals. The electron density is zero along
the bond.
• Carbon is able to form several different hybrid
orbitals depending on how many other atoms
it is bound to
sp
sp2
Overlap of
py orbitals
in O2:
C
sp3
sp2 hybridization in Carbon, C2H4
sp2 hybridization in Carbon, C2H4
• Recall – you only need hybrid orbitals for
the first bond between 2 atoms!
• The rest of the orbitals (the ones that will
form the double bond) remain as normal p
orbitals
• This hybridization forms a planar molecule
with 120˚ angles between atoms.
C

sp hybridization in Carbon, C2H2
sp hybridization in Carbon, C2H2
• Each carbon needs 2 hybrid orbitals for the first
bond with the C and H atoms.
• The double and triple bond is formed from 2 normal
p orbitals which will form 2 π bonds above and
below the central plane of the molecule
• The unpaired electrons in the two p orbitals of
the two adjacent carbon atoms share
electrons by forming two π bonds.
sp hybridization in Carbon, C2H2
• This hybridization forms a linear molecule
with 180˚ angles between atoms.
UNIT 2 Chapter 4: Chemical Bonding and Properties of Matter Section 4.2
Determining the Hybridization of
the Central Atom of a Molecule or Ion
