1. Chapter 5. Covalent Compounds (Molecular Compounds)
heat
NaCl (ionic compound)
Na+ + Cl- (gas)
Heat
H2O (liquid)
H2O molecules (gas)
NH3
NH3 (molecules)
etc.
A molecular formula tells the # of atoms of each element
in a molecule of the compound
C2F C2H6O

2. A. Covalent bonds
Example H + H H2
Sharing of electrons H H
H H
or H H

3. Draw the Lewis structures of
+ Cl
HCl H
Lewis structure
Octet rule
Draw the Lewis structures of
H2O

4. NH3
CH4

5. Consider O2
Consider N2
Draw Lewis structures for the following compounds
CH2O
C2H4

6. = B. Coordinate Covalent Bonds (less common) BH3
Electron deficient compound
NH3 =
Coordinate covalent bond

7. Draw Lewis structure for each of the following molecular
formulas in the most stable form (by pure sharing of electrons).
a) PCl b) C2F6
c) CH2O2 d) CH3N
e) C2H2Cl2 f) N2O2

8. Common elements in covalent compounds: C, O, N

9. For compounds containing C, H, O, N (the big 4), and F, try this

10. HCN
C3H4
CO2

11. C. Compounds not following the Octet Rule
PCl5

12. 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.
NO2-

13. ? ? - - - Resonance O N O NO2- Resonance structures
or resonance contributors
The real molecule or ion is a resonance hybrid of the
resonance structures.
Each resonance structure is less stable than the resonance hybrid.

14. Neutral molecules with a central atom

15. Polyatomic ions
Examples:
NO3-
SO32-

16. Lewis dot structures of ionic compounds:
K2SO3
Ca(NO3)2

17. H Cl H Cl 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

18. Fig.5.11 Pauling Electronegativity Values

19. (a) (b) Nonpolar and Polar Covalent Bond
G. Polar covalent bond
Figure 5.12:
(a) (b) Nonpolar and Polar Covalent Bond d+ d-
H Cl
H Cl
Polar covalent bond

20. Electronegativity Difference
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
Between 1.5 and 2.0
(between nonmetals)
Ionic
(metal and nonmetal)
Greater than 2.0
Examples:

21. Arrange the following bonds from most to least polar:
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

22. Molecular Geometry
Valence shell electron pair repulsion (VSEPR) theory
CH4
All 4 bonds are equivalent
Lewis structure
Electron pair arrangement
Tetrahedral
109.5o
Molecular geometry C s
p p p
hybridize
sp3 sp3 sp3 sp3
Four sp3 hybrid orbits

23. o
180o
120o
109.5o
hybrid orbitals: sp sp sp3

24. Lewis structure
NH3
sp3
Electron pair arrangement: tetrahedral
Molecular geometry: trigonal pyramidal
H2O
sp3
Electron pair arrangement: tetrahedral
Molecular geometry: angular

25. Three sp2 hybrid orbitals
BH3
sp2
Lewis structure
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

26. Electron pair arrangement: trigonal planar Molecular geometry: bent
SO2
sp2
Electron pair arrangement: trigonal planar
Molecular geometry: bent
BeH2 s p
s p
Two sp hybrid orbitals sp
sp sp
Electron pair arrangement: linear
Molecular geometry: linear

27. o
180o
109o
120o
hybrid orbitals: sp sp sp3

28. The Shape (Geometry) of Molecules

29. Summary
# of groups electron pair makeup molecular hybrid
of electrons (density) of e- groups geometry orbitals
around arrangement
central atom
4 tetrahedral 4 bonding tetrahedral
3 bonding trigonal
1 nonbonding pyramidal sp3
2 bonding angular
2 nonbonding (bent)
3 Trigonal 3 bonding Trigonal
planar planar sp2
2 bonding Angular
1 nonbonding
2 Linear 2 bonding Linear sp

30. Examples:
Electron pair
arrangement
Molecular
geometry
Hybird
orbitals
HCN

31. I. Polarity of Molecules
H Cl
One polar bond
Polar molecule
net
Polar molecule
Non polar molecule
The 2 polar bonds cancel each other

33. Summary:
Draw the Lewis structure
If all electron groups around the central atom are
connected to the same atom – nonplar
otherwise - polar

34. J. Naming of binary molecular compounds
mono di
tri- 3
tetra- 4
penta- 5
hexa- 6
hepta- 7
octa- 8
ennea-(neno) 9
deca- 10
P2O5
diphosphorus pentaoxide
N2O4 dinitrogen tetraoxide
CO2
SO2 NO