Covalent
Bonding
Bonding models for
methane, CH4. Models are
NOT reality. Each has its
own strengths and
limitations.
Covalent Bonds
Polar-Covalent bonds
 Electrons are unequally shared
 Electronegativity difference between .3
and 1.7
Nonpolar-Covalent bonds
 Electrons are equally shared
 Electronegativity difference of
0 to 0.3
Covalent Bonding Forces
 Electron – electron repulsive forces
 Proton – proton repulsive forces
 Electron – proton attractive forces
Bond Length and Energy
Bond
C - C
C = C
C  C
C - O
C = O
C - N
C = N
C  N
Bond length
(pm)
154
Bond Energy
(kJ/mol)
347
Double
Triple
Single
Double
134
120
143
123
614
839
358
745
Single
Double
Triple
143
138
116
305
615
891
Bond type
Single
Bonds between elements become shorter and
stronger as multiplicity increases.
Bond Energy and Enthalpy
H   Dbonds broken   Dbonds formed
Energy required
Energy released
D = Bond energy per mole of bonds
Breaking bonds always requires energy
Breaking = endothermic
Forming bonds always releases energy
Forming = exothermic
The Octet Rule
Combinations of elements tend to form so that
each atom, by gaining, losing, or sharing
electrons, has an octet of electrons in its
highest occupied energy level.
Monatomic chlorine
Diatomic chlorine
The Octet Rule and
Covalent Compounds
 Covalent compounds tend to form so that
each atom, by sharing electrons, has an octet
of electrons in its highest occupied energy level.
 Covalent compounds involve atoms of
nonmetals only.
 The term “molecule” is used exclusively for
covalent bonding
The Octet Rule:
The Diatomic Fluorine Molecule
F
F
1s
2s
2p
1s
2s
2p
F F
Each has
seven valence
electrons
The Octet Rule:
The Diatomic Oxygen Molecule
O
O
1s
2s
2p
1s
2s
2p
O O
Each has six
valence
electrons
The Octet Rule:
The Diatomic Nitrogen Molecule
N
N
1s
2s
2p
1s
2s
2p
N N
Each has five
valence
electrons
Lewis Structures
Lewis structures show how valence electrons
are arranged among atoms in a molecule.
Lewis structures Reflect the central idea
that stability of a compound relates to
noble gas electron configuration.
Shared electrons pairs are covalent bonds
and can be represented by two dots (:) or
by a single line ( - )
Comments About the Octet Rule
 2nd row elements C, N, O, F observe the octet
rule (HONC rule as well).
 2nd row elements B and Be often have fewer
than 8 electrons around themselves - they are
very reactive.
 3rd row and heavier elements CAN exceed the
octet rule using empty valence d orbitals.
 When writing Lewis structures, satisfy octets
first, then place electrons around elements
having available d orbitals.
Lewis Structures
Show how valence electrons are arranged
among atoms in a molecule.
Reflect the central idea that stability of a
compound relates to noble gas electron
configuration.
The HONC Rule
Hydrogen (and Halogens) form one covalent
bond
Oxygen (and sulfur) form two covalent bonds
One double bond, or two single bonds
Nitrogen (and phosphorus) form three covalent
bonds
One triple bond, or three single bonds, or one
double bond and a single bond
Carbon (and silicon) form four covalent bonds.
Two double bonds, or four single bonds, or a triple
and a single, or a double and two singles
Completing a Lewis Structure -CH3Cl
Make the atom wanting the most bonds central
Add up available valence electrons:
Join peripheral atoms
to the central atom
with electron pairs.
H
..
..
Complete octets on
H
atoms other than
hydrogen with remaining
electrons
C
..
H
Total = 14
..
Cl
..
..
C = 4, H = (3)(1), Cl = 7
Multiple Covalent Bonds:
Double bonds
H
H
C
H
H
C
H
C
H
H
Ethene
Two pairs of shared electrons
C
H
Multiple Covalent Bonds:
Triple bonds
H
C
C
H
H
C
C
H
Ethyne
Three pairs of shared electrons
Acetic Acid
H
Two electrons (one bond) per
hydrogen
O
Eight electrons (four
bonds) per carbon
H C C
H
O
H
Eight electrons (two bonds,
two unshared pairs) per
oxygen