Covalent Bonding
Sec. 8.3: Molecular Structures
Objectives
List the basic steps used in drawing
Lewis structures.
 Explain why resonance occurs and
identify resonance structures.
 Explain 3 exceptions to the octet rule
and identify molecules in which these
exceptions occur.

Molecular Structures
Molecular structures are used to predict the
arrangement of atoms in a covalent
compound
 5 models are used - see pg. 253, Fig. 13
 The Lewis structure model is the most useful;
letter symbols and bonds show the position of
atoms & “dots” show unpaired electrons.
 Electron dot diagrams must be drawn to
predict the structural formula.

Determining Lewis Structures of
Molecular Compounds
1. Predict the location of the atoms -
• Hydrogen is always a terminal or end
atom; it only shares 1 electron pair.
• The atom with the least attraction for
shared electrons is the central atom - the
central atom is usually the one closer to
the left in the periodic table.
• All other atoms become terminal atoms.
• For example, N is the central atom in NH3; C is
the central atom in CO2.
Determining Lewis Structures of
Molecular Compounds
2.Find the total number of electrons
available for bonding by adding up the
number of valence electrons of each atom
in the compound.
• In NH3, N has 5 and each H has 1, so the
total is 8.
• In CO2, C has 4 and each O has 6, so the total
is 16.
Determining Lewis Structures of
Molecular Compounds
3. Determine the number of bonding pairs
by dividing the total number of electrons
available by 2.
– For NH3: 8/2 = 4
– For CO2: 16/2 = 8
Determining Lewis Structures of
Molecular Compounds
4. Place one bonding pair between the
central atom and each terminal atom.
H
N
H
H
O C
O
Determining Lewis Structures of
Molecular Compounds
5A. Subtract the number of pairs you used
in step 4 from the total pairs you
determined in step 3.
– For NH3, there are 4 bonding pairs and 3 were
used in step 4. One bonding pair remains.
– For CO2, there are 8 bonding pairs and 2 were
used in step 4. 6 bonding pairs remain.
Determining Lewis Structures of
Molecular Compounds
5B. The remaining pairs include lone
pairs and pairs used in double and
triple bonds.
• First, place lone pairs around each terminal
atom (EXCEPT HYDROGEN) to satisfy the
octet rule.
• Next, assign any remaining pairs to the
central atom.
In NH3, 1 bonding pair
remained. Lone pairs cannot be
placed around H, so it must go
on the N.
H–N–H
H
In CO2, 6 bonding pairs
remained, so they are
placed around the O’s.
O–C-O
Determining Lewis Structures of
Molecular Compounds
6. Determine whether the central atom
satisfies the octet rule. If it doesn’t, you
MUST convert one or two of the lone pairs
on the terminal atoms into a double or
triple bond between the terminal and
central atoms. Keep in mind, carbon,
nitrogen, oxygen and sulfur often form
double and triple bonds.
O–C-O
Since the C in CO2 is only surrounded
by 2 electron pairs, a lone pair on each
O must be converted to a double bond.
O=C=O
Practice Problems

NF3

CS2

H2Se
Polyatomic ions
The atoms within the ion are covalently
bonded.
 Drawing Lewis structures is similar to
the procedure just practiced. The
difference is in step #2.

– Determine the number of electrons
available in the atoms of the ion
Then subtract the ion charge if it is
positive and add the ion charge if it is
negative
– Put a bracket around the ion structure with
the charge as a superscript
Example:
-3
PO4
The number of valence electrons available
are 29
 3 must be added to this
 Total of 32 electrons are available - this is
16 bonding pairs

O
0–P–0
O
3-
Practice Problems

ClO4-

NH4+

AsO4-3

TeO4-2
Resonance Structures

Using the same sequence of atoms, it is
possible to have more than one correct
Lewis structure when a molecule or
polyatomic ion has both a double bond
and a single bond.
Resonance Structures
Resonance is a condition that occurs
when more than one Lewis structure
can be written for a molecule or ion.
 Resonance structures are the 2 or more
correct Lewis structures for a molecule
or ion

– They differ ONLY in position of the lone &
bonding pairs, NEVER atom positions
Resonance Structures
Each actual molecule or ion that
undergoes resonance behaves as if it has
only one structure.
 The actual bond lengths are equal to each
other - the bond length is an average of
the bonds in the resonance structure.

– All bonds are shorter than single bonds but
longer than double bonds!
Practice Problems

Draw the resonance structures of
– O3
– NO2– SO4-2
Exceptions to the Octet Rule
Some molecules and ions do not obey the
octet rule.
 Three (3) reasons exist for these
exceptions.
– First, a small group of molecules has an
odd number of valence electrons and
cannot form an octet around each atom.

Exceptions to the Octet Rule
Exceptions to the Octet Rule
Second, some compounds form with fewer
than 8 electrons present around an atom
– a suboctet.
 This group is relatively rare; BH3 is an
example.

Compounds like BH3 tend to be reactive
- they will share an entire pair donated
by another atom.
 If this occurs, the bond formed is called
a coordinate covalent bond.

– Definition: a bond formed when an atom
donates a pair of electrons to be shared
with an atom/ion that needs 2 electrons to
become stable.
– Atoms/ions with lone pairs often form
coordinate covalent bonds with atoms/ions
that need 2 electrons
Exceptions to the Octet Rule
Third, some compounds have central
atoms that contain more than 8 valence
electrons.
 This electron arrangement is referred to
as an expanded octet.

– Expanded octets are explained by
considering the d orbitals of certain
elements
– Lewis structures are drawn by adding lone
pairs or more than 4 bonding pair to the
central atom
Exceptions to the Octet Rule