This document discusses Lewis symbols and the octet rule, ionic and covalent bonding, electronegativity, Lewis structures, and bond energies. The octet rule states that atoms bond to gain, lose, or share electrons to achieve a full outer shell of 8 electrons. Ionic bonds form when electrons are transferred between atoms. Covalent bonds form when electrons are shared between atoms. Electronegativity determines whether bonds are nonpolar, polar covalent, or ionic. Lewis structures show bonding and lone pairs of electrons. Formal charges indicate unequal electron sharing. Bond energies quantify bond strength.

2. 2
9.1 LEWIS SYMBOLS AND THE OCTET RULE

3. 3
Valence electrons are the outer shell electrons of an
atom. The valence electrons are the electrons that
participate in chemical bonding.
1A 1ns1
2A 2ns2
3A 3ns2
np1
4A 4ns2
np2
5A 5ns2
np3
6A 6ns2
np4
7A 7ns2
np5
Group # of valence e-
e-
configuration

4. 4
Lewis Dot Symbols for the Representative Elements &
Noble Gases
Octet rule states that an atom tends to gain, lose or share e-
until it has eight e- in the valence shell

5. 5

6. 6

7. 7

8. 8
9.2 IONIC BONDING AND THE LATTICE ENERGY

9. 9
Li + F Li+
F -
The Ionic Bond
1s2
2s1
1s2
2s2
2p5
1s2
1s2
2s2
2p6
[He]
[Ne]
Li Li+
+ e-
e-
+ F F -
F -Li+
+ Li+
F -
LiF
Ionic bond: the electrostatic force that holds ions together in an
ionic compound.

10. 10
Lattice energy can be calculated using Hess’s Law
Electrostatic (Lattice) Energy
Lattice energy (U) is the energy required to completely separate
the ion in one mole of a solid ionic compound into gaseous ions.

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9.3 COVALENT BONDING AND THE LEWIS STRUCTURES

12. 12
A covalent bond is a chemical bond in which two or more
electrons are shared by two atoms.
Why should two atoms share electrons?
F F+
7e-
7e-
F F
8e-
8e-
F F
F F
Lewis structure of F2
lone pairslone pairs
lone pairslone pairs
single covalent bond
single covalent bond

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8e-
H HO+ + OH H O HHor
2e-
2e-
Lewis structure of water
Double bond – two atoms share two pairs of electrons
single covalent bonds
O C O or O C O
8e-
8e-
8e-
double bonds double bonds
Triple bond – two atoms share three pairs of electrons
N N
8e-
8e-
N N
triple bond
triple bond
or

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9.4 ELECTRONEGATIVITY

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The Electronegativities of Common Elements
Electronegativity is the ability of an atom to attract the
electrons toward itself in a chemical bond.
Non metalmetal

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• The electronegativity of H is 2.1; Cl is 3.0.
• Since there is a difference in electronegativity
between the two elements (3.0 – 2.1 = 0.9), the
bond in H–Cl is polar.
• Cl is more electronegative, the bonding electrons
are attracted toward the Cl atom and away from
the H atom.
• Cl atom = slightly negative charge
• H atom = slightly positive charge.
δ+
H–Cl δ–

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H F
FH
Polar covalent bond or polar bond is a covalent bond
with greater electron density around one of the two
atoms (unequal sharing of e- )
electron
rich
region
electron
poor
region
e-
riche-
poor
δ+
δ-

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Nonpolar covalent bond results when two identical
non-metals equally share electrons between them
(equal sharing of e-). Eg. H2, N2, Cl2

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Covalent
share e-
Polar Covalent
partial transfer of e-
Ionic
transfer e-
Increasing difference in electronegativity
Classification of bonds by difference in electronegativity
Difference Bond Type
0 Nonpolar Covalent
≥ 2 Ionic
0 < and <2 Polar Covalent
•No electronegativity difference between two atoms leads to a pure
non-polar covalent bond.
•A small electronegativity difference leads to a polar covalent bond.
•A large electronegativity difference leads to an ionic bond.

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H-H = 0.0 H-Cl = 0.9 Na-Cl = 2.1

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Classify the following bonds as ionic, polar covalent, or
covalent: The bond in CsCl; the bond in H2S; and the NN bond
in H2NNH2.
Cs – 0.7 Cl – 3.0 3.0 – 0.7 = 2.3 Ionic
H – 2.1 S – 2.5 2.5 – 2.1 = 0.4 Polar Covalent
N – 3.0 N – 3.0 3.0 – 3.0 = 0 Non polar covalent

22. The formal charge is the charge that an atom seems to
have in a Lewis structure.
formal charge
on an atom in
a Lewis
structure
=
1
2
total number
of bonding
electrons( )
total number
of valence
electrons in
the free atom
-
total number
of nonbonding
electrons
-
The sum of the formal charges of the atoms in a molecule
or ion must equal the charge on the molecule or ion.
9.5 LEWIS STRUCTURE & FORMAL CHARGE
Lewis structures are diagrams that show the bonding
between atoms of a molecule and the lone pairs of electrons
that may exist in the molecule.
Why formal charge ?
A Lewis structure in which there are no formal charges is
preferred.

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H C O H
C – 4 e-
O – 6 e-
2H – 2x1 e-
12 e-
2 single bonds (2x2) = 4
1 double bond = 4
2 lone pairs (2x2) = 4
Total = 12
formal charge
on C
= 4 -2 -½ x 6 = -1
formal charge
on O
= 6 -2 -½ x 6 = +1
formal charge
on an atom in
a Lewis
structure
=
1
2
total number
of bonding
electrons( )
total number
of valence
electrons in
the free atom
-
total number
of nonbonding
electrons
-
-1 +1

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C – 4 e-
O – 6 e-
2H – 2x1 e-
12 e-
2 single bonds (2x2) = 4
1 double bond = 4
2 lone pairs (2x2) = 4
Total = 12
H
C O
H
formal charge
on C
= 4 -0 -½ x 8 = 0
formal charge
on O
= 6 -4 -½ x 4 = 0
formal charge
on an atom in
a Lewis
structure
=
1
2
total number
of bonding
electrons( )
total number
of valence
electrons in
the free atom
-
total number
of nonbonding
electrons
-
0 0

25. The enthalpy change required to break a particular bond in one
mole of gaseous molecules is the bond enthalpy.
H2 (g) H (g) + H (g) ∆H0
= 436.4 kJ
Cl2 (g) Cl (g) + Cl (g) ∆H0
= 242.7 kJ
HCl (g) H (g) + Cl (g) ∆H0
= 431.9 kJ
O2 (g) O (g) + O (g) ∆H0
= 498.7 kJ O O
N2 (g) N (g) + N (g) ∆H0
= 941.4 kJ N N
Bond Enthalpy
Bond Enthalpies
Single bond < Double bond < Triple bond
9.6 BOND ENERGIES

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BE = measure the strength
of a bond
the larger the BE,
the stronger the bond

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Use bond enthalpies to calculate the enthalpy change for:
H2 (g) + F2 (g) 2HF (g)
∆H0
= ΣBE(reactants) – ΣBE(products)
Type of
bonds
broken
Number of
bonds
broken
Bond enthalpy
(kJ/mol)
Enthalpy
change
(kJ/mol)
H H 1 436.4 436.4
F F 1 156.9 156.9
Type of
bonds
formed
Number of
bonds
formed
Bond enthalpy
(kJ/mol)
Enthalpy
change (kJ/mol)
H F 2 568.2 1136.4
∆H0
= 436.4 + 156.9 – 2 x 568.2 = -543.1 kJ/mol