Chapter 5. Covalent Compounds (Molecular Compounds)
heat
NaCl (ionic compound)
H2O (liquid)
Heat
NH3
Na+ + Cl- (gas)
H2O molecules (gas)
NH3 (molecules)
etc.
A molecular formula tells the # of atoms of each element
in a molecule of the compound
C2F4
C2H6O
1
A. Covalent bonds
Example
H
H+H
H
H2
Sharing of electrons
HH
or H H
2
H + Cl
H
HCl
Cl
H Cl
Lewis structure
Octet rule
Draw the Lewis structures of
H2O
3
NH3
CH4
4
Consider O2
Consider N2
Draw Lewis structures for the following compounds
CH2O
C2H4
5
B. Coordinate Covalent Bonds (less common)
BH3
H
H
H
H
B
H
B
Electron deficient compound
H
NH3
H
H
B
H
N
H H
H
= H
H
H
B
N H
H
H
Coordinate covalent bond
6
Draw Lewis structure for each of the following molecular
formulas in the most stable form (by pure sharing of electrons).
a) PCl3
b) C2F6
c) CH2O2
d) CH3N
e) C2H2Cl2
f) N2O2
7
Common elements in covalent compounds: C, O, N
C
C
N
N
O
O
C
C
N
8
For compounds containing C, H, O, N (the big 4), and F, try this
9
HCN
C3H4
CO2
10
C. Compounds not following the Octet Rule
NO
PCl5
11
E. Lewis structures of Polyatomic ions or molecules with a
central atom.
1. Calculate the total number of valence electrons.
2 Draw a single bond between the central atom and each of
the surrounding atoms.
3. Add nonbonding electrons to surrounding atoms such that
each has an octet of electrons (2 on H).
4. Place the remaining electrons on the central atom.
5. If the central atom does not have octet of electrons, use one
or two pairs of nonbonding e’s from the surround atoms
to form double or triple bonds with the central atom.
6. Check the total number of electrons.
NO212
Resonance
NO2-
O
N
O
-
-
N
N
O
O
O
Resonance structures
or resonance contributors
O
?
?
The real molecule or ion is a resonance hybrid of the
resonance structures.
Each resonance structure is less stable than the resonance
hybrid.
13
Neutral molecules with a central atom
14
Polyatomic ions
Examples:
NO3-
SO32-
15
Lewis dot structures of ionic compounds:
K2SO3
Ca(NO3)2
16
F. Electronegativity (EN)
Electronegativity of an element = the relative tendency
of its atoms to attract the bonding electron pair.
H Cl
or
H
Cl
EN of Cl > EN of H
17
Fig.5.11
Pauling Electronegativity Values
18
G. Polar covalent bond
Figure 5.12:
(a) (b) Nonpolar and Polar Covalent Bond
d-
d+
H Cl
H Cl
Polar covalent bond
19
The relative E.N. determines the bond type
Bond Type
Electronegativity Difference
Nonpolar Covalent
0.4 or less
Polar Covalent
Greater than 0.4 to 1.5
Polar Covalent
Between 1.5 and 2.0
(between nonmetals)
Ionic
Between 1.5 and 2.0
(metal and nonmetal)
Ionic
Greater than 2.0
Examples:
20
Exercise
Arrange the following bonds from most to least polar:
a) N-F
O-F
C-F
b) C-F
N-O
Si-F
c) Cl-Cl
B-Cl
S-Cl
21
H. Molecular Geometry
Valence shell electron pair repulsion (VSEPR) theory
H
CH4
H
C
H
All 4 bonds are equivalent
H
Lewis structure
Electron pair arrangement
H
H
Tetrahedral
109.5o
C
C
H
H
H
H
H
H
Molecular geometry
C
s
p
p
p
sp3 sp3 sp3 sp3
Four sp3 hybrid orbits 22
180o o
hybrid orbitals:
sp
120o
109.5o
sp2
sp3
23
Lewis structure
NH3
H
N
H
sp3
N
H
H
H
H
Electron pair arrangement: tetrahedral
Molecular geometry:
trigonal pyramidal
H
H2O
H
O
H
O
sp3
H
Electron pair arrangement: tetrahedral
Molecular geometry:
angular
24
BH3
H
B
H
H
sp2
B
H
Lewis structure
H
H
Trigonal planar
B
s
p p p
s
p
p
p
sp2 sp2 sp2
Three sp2 hybrid orbitals
Electron pair arrangement: trigonal planar
Molecular geometry:
trigonal planar
25
O
SO2
S
O
sp2
S
O
O
Electron pair arrangement: trigonal planar
Molecular geometry:
bent
BeH2
s
H
s
p
Be
H
sp
p
sp sp
Electron pair arrangement: linear
Molecular geometry:
linear
Two sp hybrid orbitals
26
180oo
hybrid orbitals:
sp
109o
120o
sp2
sp3
27
The Shape (Geometry)
of Molecules
28
Summary
# of groups electron pair makeup
of electrons
(density) of e- groups
around
arrangement
central atom
4
tetrahedral
4 bonding
3
2
Trigonal
planar
Linear
molecular
geometry
hybrid
orbitals
tetrahedral
3 bonding
1 nonbonding
trigonal
pyramidal
2 bonding
2 nonbonding
angular
(bent)
3 bonding
Trigonal
planar
2 bonding
1 nonbonding
Angular
2 bonding
Linear
sp3
sp2
sp
29
Examples:
Electron pair
arrangement
Molecular
geometry
Hybird
orbitals
HCN
30
I. Polarity of Molecules
H
Cl
One polar bond
Polar molecule
O
H
O
H
C
O
net
Polar molecule
Non polar molecule
The 2 polar bonds cancel each other
31
32
Summary:
• Draw the Lewis structure
• If all electron groups around the central atom are
connected to the same atom – nonplar
otherwise - polar
33
J. Naming of binary molecular compounds
monoditritetrapentahexaheptaoctaennea-(neno)
deca-
1
2
3
4
5
6
7
8
9
10
P2O5
diphosphorus pentaoxide
N2O4 dinitrogen tetraoxide
CO2
SO2
NO
34