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Sec. 4.3: Lewis Bonding Theory:
Lewis Symbols for Atoms
Chapter Outline
 Element symbol = nucleus + core electrons
 Valence electrons are drawn as dots around
the symbol
••
•• •
•
 4.1 Types of Chemical Bonds
 4.2 Naming Compounds and Writing
Formulas (Lab)
 4.3 Lewis Structures
 4.4 Electronegativity, Unequal Sharing,
and Polar Bonds
 4.5 Vibrating Bonds and the Greenhouse Effect
 4.6 Resonance
 4.7 Formal Charge: Choosing among Lewis Structures
 4.8 Exceptions to the Octet Rule
 4.8 The Lengths and Strengths of Covalent Bonds
O
 Up to 4 valence electrons are placed around
the symbol one at a time; additional electrons
are paired up
 The result is up to 4 pairs of electrons = octet
H
•
 NOTE: hydrogen can not have an octet. When
forming bonds with other atoms, it can have a
maximum of 2 electrons in its valence shell
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Lewis Symbols and the Periodic Table
# of valence e-
Lewis Symbols and the Periodic Table
Group
e- configuration
Lewis Dot Symbol
1A
ns1
1
2A
ns2
2
3A
ns2np1
3
Main Group Elements:
4A
ns2np2
4
5A
ns2np3
5
6A
ns2np4
6
Members of same family
have same number of
valence electrons, and
similar bonding capacities.
7A
ns2np5
7
Lewis Structures of Ionic Compounds
Na
[Ne]3s1
e- + Cl
[Ne]2s22p5
[Ne]
-
Sample Exercise 4.8 (Modified)
Draw the Lewis symbols of the monatomic ions
formed by calcium and oxygen. Then draw the
Lewis structure of calcium oxide (CaO).
Na+ + e-
Cl
Unpaired dots = bonding
capacity.
Na+ + Cl
-
Na+ Cl
-
[Ne]3s23p6 = [Ar]
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Lewis Structures of Molecular
(Covalent) Compounds
A covalent bond is a chemical bond in which two or
more electrons are shared by two nonmetals,
resulting in an octet for both atoms. For example Lewis structure of F2
Lewis structure of H2O
Guidelines for Drawing Lewis Structures
(updated later on with the concept of “formal charge”)
4. Sum up the total number of valence electrons
(use the group number), and calculate the
number of pairs.
Guidelines for Drawing Lewis Structures
(updated later on with the concept of “formal charge”)
1. Hydrogen is always a terminal atom because it
can form only one bond.
2. The CENTRAL ATOM usually has the lowest
electron affinity (or electronegativity as defined
later)
3. Arrange the atoms geometrically and
symmetrically.
e.g. CHCl3
Multiple Bonds – sharing more than one
pair of electrons
Double bond – two atoms share two pairs of electrons
CO2
H2CO
5. Connect the atoms together so that each atom
has an octet (except H). You may have to form
multiple bonds.
Triple bond – two atoms share three pairs of electrons
N2
C2H2
Lewis Structures of Charged Species
ClO-
NO2+
Electronegativity, Unequal Sharing, and Polar Bonds
 Electronegativity ():
• Ability of an atom to attract bonding electrons.
• Periodic trend similar to ionization energy.
Electronegativities
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Ionization Energies
and Electronegativies
Polar Covalent Bonds
• Unequal sharing of electrons in a covalent bond
resulting in an uneven distribution of charge.
• Results from differences in electronegativity.
• Dipole Moment = polarity indicated by arrow
pointing to more “partially negative” end, with a
“partially positive” charge on the opposite size
ߜା
EN increases
across a row.
EN decreases
down a column.
Polar Covalent Bonds - Unequal sharing of electrons
resulting in an uneven distribution of charge.
+1
ߜା
ߜି
-1
Difference
Bond Type
Cl2
0 - 0.4
Covalent
HCl
0.4 - 2
Polar Covalent
2
NaCl
ߜି
H
Cl
Electronegativity Trends
 As seen previously, electronegativity increases moving
up to the right in the periodic table. (Noble gases not
included.)
 Bond polarity increases as ∆EN increases.
Ionic
∆EN = 1.9
0.9
0.7
0.4
Sample Exercise 4.12
Rank, in order of increasing polarity, the bonds
formed between O and C
Cl and Ca
N and S
O and Si
Are any of these bonds considered ionic?
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