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Chapter 8
Chemical Bonding I: Basic Concepts
8.1 Lewis Dot Symbols:
• Valence electrons determine an element’s
chemistry.
• Lewis dot symbols represent the valence
electrons of an atom as dots arranged around
the atomic symbol.
• Most useful for main-group elements
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Write Lewis dot symbols for the following:
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8.2 Ionic Bonding:
• Ionic bond: electrostatic force that holds oppositely
charge particles together
• Formed between cations and anions
Example:
o
m.p. = 801 C

and
H f = – 410.9 kJ/mol
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Microscopic View of NaCl Formation
• Lattice energy
= the energy required to completely
separate one mole of a solid ionic compound into gaseous
ions
• NaCl(s)  Na+(g) + Cl(g)
Hlattice = +788 kJ/mol
• (Lattice energies are always positive)
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Coulombic attraction:
When two charges Q1 and Q2 are separated by a distance d
The electrostatic force between the two charges is defined
as:
Q1  Q2
F
d2
When, Q1 and Q2 are similar F is repulsive force
Q1 and Q2 are different F is attractive force
• Lattice energy (like a coulombic force) depends on
• Magnitude of charges
• Distance between the charges
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Lattice energies of alkali metal iodides:
F is inversely proportional to the distance between the charges
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The ionic radii sums for LiF and MgO are 2.01 and 2.06 Å,
respectively, yet their lattice energies are 1030 and 3795
kJ/mol. Why is the lattice energy of MgO nearly four times
that of LiF?
F is directly proportional to the product of the charges
(The product of the charges in LiF is half of that in MgO)
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• Born-Haber cycle:
A method to determine lattice energies
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• Lattice energy
= the energy required to completely
separate one mole of a solid ionic compound into gaseous
ions
• NaCl(s)  Na+(g) + Cl(g)
Hlattice = +788 kJ/mol
• NaCl(s)  Na+(g) + Cl(g)
Hlattice = -788 kJ/mol
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8.3 Covalent Bonding:
• Atoms share electrons to form covalent bonds.
or H–H
• In forming the bond the atoms achieve a more
stable electron configuration.
• Octet Rule: Atoms will gain, lose, or share electrons
to acquire eight valence electrons
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Lewis Structures:
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Multiple Bonds:
- The number of shared electron pairs is the number
of bonds.
Bond strength and bond length:
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8.4 Electronegativity and Polarity:
• Nonpolar covalent bond = electrons are shared
equally by two (similar) bonded atoms
• Polar covalent bond = electrons are shared
unequally by two (different) bonded atoms
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Electronegativity: ability of an atom to draw shared
electrons to itself.
- More electronegative elements attract electrons
more strongly.
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Electronegativity: The Pauling Scale
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Electron density distributions:
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Polar and nonpolar bonds:
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Dipole moments and partial charges:
- Polar bonds often result in polar molecules.
- A polar molecule possesses a dipole.
See Problem 8.5:
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8.5 Drawing Lewis Structures:

The Lewis structure of a molecule shows how the
valence electrons are arranged among the atoms in
the molecule.

From experiment, it was shown that in forming a
stable compound, atoms try to achieve noble gas
electron configuration [8 electrons (lone-pairs plus
bonding shared electrons) around each atom of the
molecule].
Rules for writing Lewis structures:
1) Sum the valence electrons from all the atoms.
2) Use a pair of electrons to form a bond between
each pair of bound atoms. Form multiple bonds
when it is necessary (e.g. CO2 and O3).
3) Arrange the remaining electrons to satisfy the
octet rule (form single bonds for hydrogen).
4) If any electrons remain place them in pairs on the
central atom {e.g (NO3)-}.
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Examples:
1.
H2O (water):
1H
and 8O
1H
and 8O
Sum of valence electrons = 1 + 1 + 6 = 8
..
H  O H
..
2.
CO2 (carbon dioxide):
6C
and 8O
Sum of valence electrons = 4 + 6 + 6 = 16
..
..
..
..
OC O
3.
CN- (Cyanide ion):
6C
and 7N
Sum of valence electrons = 4 + 5 + 1 = 10
[CN]-
4.
NO+ (Nitrogen oxide ion) 7N and 8O
Sum of valence electrons = 5 + 6 – 1 = 10
[NO]+
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5.
N2
(Nitrogen)
7N
and 7N
Sum of valence electrons = 5 + 5 = 10
NN
6.
Cl3Sum of valence electrons = 3 x 7 + 1 = 22
..
....
..
..
..
..
(: Cl  Cl  Cl :) 
Note that there is 3 lone-paris of electrons around
the central Cl atom. This is a case of exception
that will be discussed in sec. 8.8.
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7.
What is the Lewis structure of NO3 ?
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8.6 Lewis Structures and Formal Charge:
Formal Charge:
The formal charge of an atom in a molecule is the
difference between the number of valence electrons
on the free atom and the number of valence electrons
assigned to the atom in a molecule.
Rules for Determining Formal Charges:
1. Take the sum of the lone-pair electrons and one
half the shared electrons.
2. Subtract the number of assigned electrons from the
number of valence electrons on the free atom.
3. The sum of the formal charges of all atoms equals
the overall charge.
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Example: Formal charges on the atoms in ozone:
v.es.f.a l.p.es sh.es/2
O O O
O  6  4  12 4
O
O





0
6  2  12 6
1
6  6  12 2
1
 Formal charge guidelines:
− A Lewis structure with no formal charges is
generally better than one with formal charges.
− Small formal charges are generally better than large
formal charges.
− Negative formal charges should be on the more
electronegative atom(s).
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Identify the best structure for formaldehyde below:
Identify the best structure for the isocyanate ion below:
Structure (b) is best structure of isocyanate ion
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8.7 Resonance:
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8.8 Exceptions to the Octet Rule:
Exceptions to the octet rule fall into three categories:
− Molecules with an incomplete octet
− Molecules with an odd number of electrons
− Molecules with an expanded octet
Incomplete Octets:
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Odd Numbers of Electrons:
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Expanded Octet:
− Elements of the 3rd period and beyond have dorbitals that allow more than 8 valence electrons.
..
: Cl
..
..
: Cl :
P
: Cl :
:
.. : Cl
..
..
:
Cl
..
40 electrons
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8.9 Bond Enthalpy:
•
Bond enthalpy is the energy associated with breaking a
particular bond in one mole of gaseous molecules.
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− Bond enthalpies for polyatomic molecules depend upon
the bond’s environment.
− Average bond enthalpies are used for polyatomic
molecules.
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Prediction of bond enthalpy:
Ho = BE(reactants)  BE(products)
(BE = bond energy)
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