CHEMICAL
BONDING
LEWIS THEORY OF BONDING
CHEMICAL BONDS
Bond Type
Single
Double
Triple
# of e’s
2
4
6
Notation
—
=

Bond order
1
2
3
Bond
strength
Bond length
Increases from Single to Triple
Decreases from Single to Triple
TYPES OF BONDING CONDITIONS BETWEEN
ELEMENTS
Low
Electronegativity
and low Ionization
energy (Metals)
High
electronegativity
and High
Ionization energy
(Non-metals)
Low
Electronegativity
and low Ionization
energy (Metals)
Metallic bonding
Ionic bonding
(transferring of
electrons between
atoms)
High
electronegativity
and High
Ionization energy
(Non-metals)
Ionic bonding
Covalent bonding
(sharing of electrons
between atoms)
Electronegativity
STRUCTURAL DIAGRAMS FOR
COVALENT COMPOUNDS
Draw the Lewis Diagram for nitrogen trifluoride (NF3).
Step 1. Count the valence electrons
N=5
F=7
5 + 3( 7) = 26 valence electrons
STRUCTURAL FORMULA FOR
COVALENT COMPOUNDS
Step 2. Write a skeletal structure. Use the least
electronegative atom in the centre
Electronegativity: N = 3.0 & F = 4.0
F
N
F
F
= a pair of e(a single bond)
Step 3. Complete the octets for each terminal atom (except H)
..
:F..
N
:F
.. :
..
F
.. :
Step 4. Assign any additional electrons as lone pairs on the
central atom
..
:F..
..
N
:F
.. :
..
F
.. :
Example 2. COCl2 (24 electrons)
..
: Cl
..
C
:O:
..
..
Cl :
..
Step 5. Make multiple bonds where necessary to complete
the octets.
..
: Cl
..
C
:O:
..
..
:
Cl
..
Example 3. Chlorate ion, ClO3
-
((1 x 7) + (3 x 6) + 1) = 26
..
:O
..
..
Cl
: O:
..
..
O:
..
-
In some covalent compounds, the bonds
between atoms occur because one atom
has donated both electrons to the covalent
bond. This is called a coordinate covalent
bond.
+
H
H
N
H
H
: +
H+
H
N
H
H
Nitrogen supplies the two lone pair electrons to this
N-H bond. The H+ ion has no electrons.
To determine the number of coordinate
covalent bonds – subtract the bonding
capacity (lone valence electrons) from the
number of bonds the atom has.
+
H
H
N
H
H
Nitrogen
Bonds  4
Bonding capacity 3
Coordinate bonds  4-3=1
Exceptions to the Octet Rule
On occasion, both elements have the same electronegativity
or there may be two or more possible Lewis Structures.
• e.g. CS2
• (both electronegativities = 2.5)
• is it S=C=S or C=S=S ?
In such situations, one determines the
Formal Charge. The option with the lowest
formal charge has the most stable and
viable structure.
The Formal Charge for an atom is the
number of valence electrons in the free
neutral atom minus the number of valence
electrons assigned to the atom in the Lewis
structure.
Formal Charge = (# valence electrons)-(# of bonds)-(# of
unshared e-)
..
..
: C=S=S :
C
S
S
Valence electrons
4
6
6
Electrons assigned
6
4
6
Formal Charge
-2
2
0
..
..
:S=C=S :
S
C
S
Valence electrons
6
4
6
Electrons assigned
6
4
6
Formal Charge
0
0
0
In some structures the Lewis structure does
not represent the true structure of the
compound.
Bond order is the number of shared pairs of
electrons between two atoms. (i.e. – the
number of bonds between two atoms)
As the bond order increases. . .
• The length of the bond decreases.
• The energy associated with breaking the bond
increases.
Linus Carl Pauling (February 28, 1901 – August 19, 1994)
e-pairs
Notation
Name of VSEPR shape Examples
2
AX2
Linear
HgCl2 , ZnI2 , CS2 , CO2
3
AX3
Trigonal planar
BF3 , GaI3
AX2E
Non-linear (Bent)
SO2 , SnCl2
AX4
Tetrahedral
CCl4 , CH4 , BF4-
AX3E
(Trigonal) Pyramidal
NH3 , OH3-
AX2E2
Non-Linear (Bent)
H2O , SeCl2
AX5
Trigonal bipyramidal
PCl5 , PF5
AX4E
Distorted tetrahedral
(see-sawed)
TeCl4 , SF4
AX3E2
T-Shaped
ClF3 , BrF3
AX2E3
Linear
I3- , ICl2-
AX6
Octahedral
SF6 , PF6-
AX5E
Square Pyramidal
IF5 , BrF5
AX4E2
Square Planar
ICl4- , BrF4-
4
5
6