Lewis
Structures
(The Localized Electron
Model)
G. N. Lewis
1875 - 1946
Using electron-dot symbols, G. N.
Lewis developed the Localized
Electron Model of chemical bonding
(1916) in which valence electrons exist
as lone pairs or as individual electrons
seeking to form a pairing in order to
achieve an octet.
Later, Linus Pauling would
expand the Localized Electron Model
to include Resonance and Orbital
Hybridization, Collectively known as
Valence Bond Theory (1930).
In 1957 VSEPR Theory was added to predict
molecular geometry, also describing any resulting
molecular polarity in molecules.
According to Lewis Theory, there are two types
of valence electrons:
•Non-bonding (or unshared) pairs
Localized Electron Model
In Lewis’s Localized Electron
Model, molecules are described as
being composed of atoms that are
bound together by sharing pairs of
electrons. He was able to show that
the arrangement of atoms in
Linus and Ava Helen
Pauling in Munich,
molecules could be predicted based on
with Walter Heitler
the arrangements of valence electrons
(left) and Fritz London
(right). 1927
of all atoms involved in the molecule.
Walter Heitler and Fritz London (1927) were the
first to solidify Lewis’s idea by linking atomic orbital
overlap to Schrödinger’s wave equation (1925) to show
how two hydrogen atom wave functions join together to
form a covalent bond.
We have seen how we can build models of
molecules by combining atoms according to
electron dot structures.
..
:Br:
•Bonding single (or unpaired) electrons
+ 3
•Boron has three unpaired electrons therefore it can
form three covalent bonds
•Bromine has three unshared pairs and one unpaired
electron, therefore it can only form one covalent bond.
•What about nitrogen?
=
..
: N Br
.. :
:Br
.. :
Today, we are going to learn a process by
which we will be able to draw a model of any
molecule.
1
Lewis Structures
Writing Lewis Structures
PCl3
5 + 3(7) = 26
Lewis structures are
representations of molecules
showing all electrons, bonding
and nonbonding.
Writing Lewis Structures
Writing Lewis Structures
• Things to consider
when building primary
skeleton:
2. Build a reasonable
skeletal structure for
the molecule using
only single bonds.
Keep track of
the electrons:
26 − 6 = 20
The central atom
should be the least
electronegative
element that isn’t
hydrogen.
Writing Lewis Structures
Keep track of the
electrons:
26 − 6 = 20 − 18 = 2
3. Subtract the total
number of
electrons used in
the primary bonds
from the available
valence electrons.
4. Fill the octets of the
outer atoms by
adding unshared
pairs
1. Find the sum of
valence electrons of all
atoms in the molecule
from the group
number or electron
dot structure.
Keep track of
the electrons:
26 − 6 = 20
Oxygen never bonds to
itself, except in O2 and O3
Carbon atoms are usually
bonded to each other
In molecules containing
both H and O, hydrogen is
usually bonded to oxygen
Writing Lewis Structures
5. Fill the octet of
the central atom.
6. Check to see that
all atoms have
and octet and
that the correct
number of
Keep track of the electrons:
valence electrons
were used
26 − 6 = 20 − 18 = 2 − 2 = 0
2
Writing Lewis Structures
Writing Lewis Structures
7. If you run out of electrons before the
central atom has an octet…
Example:
Try building a Lewis structure for HCN
1. Let’s try drawing the Lewis Structures
for the following molecules:
A. Carbon tetrachloride
B. Ammonia
C. Oxygen
D. Carbon dioxide
E. Dihydrogen carbon monoxide
F.
Ethanal (C2H4O)
2. Let’s draw the Lewis structure for
dihydrogen sulfate and for the
sulfate anion formed when
dihydrogen sulfate is placed in
water.
5. …form multiple bonds
until it does.
•Polyatomic ions are formed from a class of
molecules called Acids, or in some rare cases,
from Bases.
•Polyatomic ions are formed as acids or bases
loose or gain hydrogen atoms.
For example:
Hydrogen
nitrate
NO3nitrate ion
Hydrogen nitrate looses a hydrogen proton when placed in
water, resulting in the formation of the nitrate ion (notice the
1- charge)
•Lewis structures for polyatomic ions
must account for the loss or gain of
valence electrons
Cations – decrease valence electrons
by amount of charge
Anions – increase valence electrons
by amount of charge
•Lewis structures for polyatomic ions
are written in brackets [ ] with the
charge denoted as a superscript.
3
3.Try drawing the Lewis structure for
hydrogen nitrate and the nitrate ion.
4. Draw the Lewis Structure for
ozone, O3.
You may notice more than one Lewis structure
can be drawn for these species.
Notice that two L.S. can be drawn
correctly for ozone, O3
RESONANCE
Resonance theory, developed by Lewis (1928), is a
key component of valence bond theory and arises
when no single conventional model using only even
number of electrons shared exclusively by two
atoms can actually represent the observed
molecule. Resonance involves modeling the
structure of a molecule as an intermediate, or
average, between several simpler but incorrect
structures.
Resonance
• But this is at odds
with the true,
observed structure
of ozone, in
which…
…both O—O
bonds are the
same length.
Resonance
• One Lewis structure cannot accurately
depict a molecule such as ozone.
• We use multiple structures, resonance
structures, to describe the molecule.
• Resonance is denoted by a double headed
arrow separating the different Lewis
Structures:
Resonance
• In truth, the electrons that form the second C—O
bond in the double bonds below do not always sit
between that C and that O, but rather can move
among the two oxygens and the carbon.
• They are not localized, but rather are delocalized.
4
Resonance
Resonance
• Observe HCO2- :
Just as green is a
synthesis of blue and
yellow…
…ozone is a
synthesis of these
two resonance
structures.
• In truth, the electrons that form the second C—O
bond in the double bonds below do not always sit
between that C and that O, but rather can move
among the two oxygens and the carbon.
• They are not localized, but rather are delocalized.
5. Draw all three resonance structures
for the nitrate ion.
5