orbitals. The s orbital is spherical in shape and p orbitals are in dumbbell
shapes.
Sigma overlap of a s orbital
with a p orbital
Pi overlap of two parallel p orbitals
Let us consider the formation of s overlap in the hydrogen molecule (H 2 ),
from two hydrogen atoms. Each hydrogen atom has one electron, which
occupies the 1s orbital. The overlap of two s orbitals, one from each of two
hydrogen atoms, forms a s bond. The electron density of a s bond is
greatest along the axis of the bond. Since s orbitals are spherical in shape,
two hydrogen atoms can approach one another from any direction resulting
in a strong s bond.
H .
H H
=
H 2
Hydrogen atoms,
each contains 1s
atomic orbital
Hydrogen atoms,
formation of bonding
molecular orbital
Hydrogen molecule
H
.
H H
:
:
2.4 Electronegativity and chemical bonding
Electronegativity is the ability of an atom that is bonded to another atom or
atoms to attract electrons strongly towards it. This competition for electron
density is scaled by electronegativity values. Elements with higher electronegativity values have greater attraction for bonding electrons. Thus, the
electronegativity of an atom is related to bond polarity. The difference in
electronegativity between two atoms can be used to measure the polarity of the
bonding between them. The greater the difference in electronegativity between
the bonded atoms, the greater is the polarity of the bond. If the difference is
great enough, electrons are transferred from the less electronegative atom to
the more electronegative one, hence an ionic bond is formed. Only if the two
atoms have exactly the same electronegativity is a nonpolar bond formed.
Electronegativity increases from left to right and bottom to top in the periodic
table as shown below (electronegativity is shown in parentheses).
1A
2A
3A
4A
5A
6A
7A
H (2.2)
Li (1.0)
Be (1.6)
B (1.8)
C (2.5)
N (3.0)
O (3.4)
F (4.0)
Na (0.9)
Mg (1.3)
Al (1.6)
Si (1.9)
P (2.2)
S (2.6)
Cl (3.2)
Br (3.0)
I (2.7)
2.4 ELECTRONEGATIVITY AND CHEMICAL BONDING
27