Drawing Lewis
Structures
Lewis Bonding
• Valence e- are the
players in bonding
• Valence e- transfer leads
to ionic bonds.
• Sharing of valence eleads to covalent bonds.
• e- are transferred or
shared to give each
atom a noble gas
configuration (a full
shell)
– the octet.
•Recall how to determine the valence electron for the
elements based on the elements position on the periodic
table.
•Lewis Dot Symbol
Ionic and Covalent Compounds
• Formation of sodium chloride:

 Na+ [ Cl

]


Cl

Na  +



• Formation of hydrogen chloride:


H Cl



Cl

H +



A metal and a nonmetal transfer electrons to
form an ionic compound. Two nonmetals
share electrons to form a covalent compound.
Ionic Compounds
Ionic compounds consist of a lattice of positive and
negative ions.
NaCl:
The Octet Rule
Atoms tend to gain, lose, or share electrons
until they have eight valence electrons.
Hydrogen is an
exception. It likes
to gain an electron
to have 2 electrons
(like Helium)
Lewis Structures For Covalent
Compounds
• A valid Lewis structure should have an octet for
each atom except hydrogen (which should have a
duet)
H  +  H  H H or H H

Nonbonding electrons
called “lone pairs”

Cl
Cl



Bonding
Electrons
 (2 per line)
Cl
Cl
 






Cl
+
Cl



Cl2:




H2:
Double and Triple Bonds
• Atoms can share four electrons to form a double
bond or six electrons to form a triple bond.


N 2:
N N

O
=O



O 2:
• Triple bonds are the shortest and strongest
whereas single bonds are the longest and
weakest
How to draw a Lewis Structure
•
•
•
•
•
•
•
•
•
1. Find the number of valence electrons you already Have. To do this,
determine the number of valence electrons for each element in the structure
and find the sum.
2. Find the total number of electrons Needed by all elements in the
structure. Every element needs 8 except hydrogen which needs 2.
3. Find the total number of electrons Shared between all elements. This is
found by Need – Have.
4. Find the number of bonds by dividing Shared by 2.
Other helpful hints:
A. You can calculate the number of Non Bonding Electrons by
Have - Shared
B. Connectivity - From the Chemical Formula, determine the atom
connectivity for the structure.
i. Given a chemical formula, ABn, A is the central atom and B flanks the A
atom. i.e., NH3, NCl3, NO2. In these examples, N is central in the structure.
ii. H and F are never central atoms.
• the number of bonds an atom prefers
depending on the number of valence electrons

Fam ily
H alogen s
F, B r, C l, I
C alcogen s
O, S
N itro gen
N, P
C arb on
C , Si
# C ovalent Bonds*
1 bond
O


N

3 bond often
C

4 bond
X
often
2 bond often
always
In general, these are the number of bonds formed by these atoms.
Using the hnsb System
•
•
•
•
•
An example: CH4
h 8
n 16
s 16 – 8 = 8
b 8/2 = 4
• nbe 8 –8 = 0
• conn C in middle
• bonds: C 4 and H 1
H H
H H
H
H C H
H
On the board we’ll do F2, N2, HCl, CO2
C
Lewis Structures
Do hnsb and draw Lewis structures for:


H N H

H

or
or
or


H O
H


NH3:


H2O:
H 
F

HF:


H F


H O
H


H N H
H
It’s because
he’s bald,
that’s why…
Drawing Lewis Structures



COCl2


O


Cl
C
Cl




H O
Cl


CH3OH
H

H C O
H

H

HOCl
Charged Polyatomic Groups
• For negative charges, add the correct
amount of electrons to your Have group
• For positive charges, subtract the correct
amount of electrons from your Have group
• Examples: CN- , H3O+, NO2-
• Other hints:
i. Break as little of the ‘rules’ as possible.
ii. If an atom has:
a. More bonds than it ‘should’ have then you
would add a ‘+’ charge per extra bond.
b. Less bonds than it ‘should’ have, then you
would add a ‘-’ charge per missing bond.
iii. -COOH is drawn as…
•
•
•
•
For polyatomic ions, when finished drawing
the Lewis structure, put square brackets
around it, and the ion charge in the top right
corner outside of the brackets.
Beyond the Octet Rule
• There are numerous exceptions to the octet rule.
• We’ll deal with two classes of violation here:
– Sub-octet systems
– Valence shell expansion
Beyond the Octet Rule
Sub Octet Systems
• Some atoms (Be and B in particular) undergo
bonding, but will form stable molecules that do
not fulfill the octet rule.
• Boron can exist with only 6 e- surrounding it.
• Beryllium can exist with only 4 surrounding e• Examples: BF3 , BeCl2
Beyond the Octet Rule
Valence Shell Expansion
• For third-row elements, the energetic proximity of
the d orbitals allows for the participation of these
orbitals in bonding.
• When this occurs, more than 8 electrons can
surround a third-row element.
• Cl and P can have 10e- around it and S can have
12e• Example: ClF3 (a 28 e- system), and SF6
F
F
Cl
F obey octet rule
F
Cl has 10e-
Resonance Structures
• We have assumed up to this point that there is one
correct Lewis structure.
• There are systems for which more than one Lewis
structure is possible:
– Different atomic linkages: Structural Isomers (organic
chemistry)
– Same atomic linkages, different bonding: Resonance
Resonance Structures (cont.)
• The classic example: O3.
O
O
O O
O
O
O
O
O
Both structures are correct!
Resonance Structures (cont.)
• In this example, O3 has two resonance structures:
O
O
O
• Each structure is perfectly valid, so conceptually,
we think of the bonding being an average of these
two structures.
• Electrons are delocalized between the oxygens
such that on average the bond strength is
equivalent to 1.5 O-O bonds.
• Try CO32-
Summary
• Remember the following:
– B and Be are often sub-octet
– Second row (Period 2) elements never exceed the octet
rule
– Third Row elements and beyond can use valence shell
expansion to exceed the octet rule, but we will only do
this for Cl, S, and P.
• In the end, you have to practice…..a lot!
Molecular Shapes
•
Both bonds and lone pairs (non-bonding
electrons) need space around an atom.
• Shapes depend on the arrangement around the
‘central’ atom.
1. Linear
-- bond angle of 1800
-- every two atom molecule is linear.
-- three atoms with the central atom having no lone
pairs is also linear.
Eg. CO, H2, CO2, BeCl2
2. Trigonal Planar
-- bond angle of 1200
-- when three atoms surround a central atom which has NO
lone pairs.
Eg. BF3, CO323. Trigonal Pyramidal
-- bond angle 1070
-- when three atoms surround a central atom which has one
lone pair.
Eg. NH3, AsH3, H3O+
4. Tetrahedral
-- bond angle 109.50
-- when four atoms surround a central atom which has NO
lone pairs.
Eg. CH4, CCl4, SiH4
5. Angular
-- bond angle of 104.50
-- three atoms present with one or two lone pairs on
the central atom.
Eg. H2O, ClO2-, O3
6. Other Shapes:
Trigonal Bipyramidal: PCl5
Octahedral: SF6
Trigonal Planar
Trigonal Pyramidal
Tetrahedral
Angular