Chapter 3: Chemical Bonding Compounds are formed from

The Octet Rule
Atoms tend to lose, gain, or share electrons until they have eight valence
electrons
Chapter 3: Chemical Bonding
Compounds are formed from chemically bound atoms or ions
Bonding involves only the valence electrons
Ionic compounds – ionic radii and lattice energies
Molecular compounds
– covalent (polar or non-polar) bond
– bond order
– bond strength
– Lewis structures
Lewis Symbols –
show the valence
electrons as dots around
the atomic symbol
1
2
3
4
1
Bond strength
Bond orders and bond distances
Bond dissociation energy (the energy required to break a bond)
Single bond distances
Bond distances of other bond types
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7
8
/mol
2
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10
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One more example: HNO3 (HONO2)
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O
Concept of Resonance
N
O
O
24e
O
H
O
O
H
More than one possible Lewis structures. The actual structure of a
molecule is taken as a blend of all the feasible Lewis structures
or
N
N
O
O
O
H
do not confuse with structural
isomers
O
N
O
O
H
Exceptions to the octet rule
Elements in the 2nd period can never have more than 8 electrons
Elements in the 3rd or higher periods can have more than 8 electrons
N
O
O
N
O
Odd-electron
molecules
Incomplete octet
Expanded octet
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4
What does writing resonance structures accomplish?
Electrons are often delocalized between two or more atoms.
Electrons in a single Lewis structure are assigned to a specific
atom. Therefore, a single Lewis structure is insufficient to electron
delocalization. Composite of resonance structures more accurately
depicts electron distribution and describes a molecule.
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Concept of formal charge
Take SCN- as an example illustrating the importance of
the formal charge concept
(Don't confuse the formal charge with the oxidation state
or the charge on an atom)
More examples to show the
application
(formal charge and expanded octet
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Octet
Lewis
Structure
Expanded
Octet Lewis
Structure
Formal
Atom Charge
Formal
Atom Charge
Expanded
to
SNF 3
F
S
F
S
2+
N
2-
F
S
Cl
S
2-
O
O
S
O
S
2+
O
1-
Cl
O
O
S
XeO3
O
0
0
S
O
S
0
O
0
Cl
12
12
O
O
S
2+
O
1-
O
S
O
S
0
O
0, 1-
12
O
O
SO32-
S
N
An atom that has a d subshell in
the valence electron shell can
accommodate more than an octet
of electrons
O
O
Cl
SO4
F
F
F
SO2Cl2
Expanded octet
N
N
O
O
S
1+
O
1-
Xe
3+
O
1-
O
O
S
O
S
0
O
0, 1-
10
O
Xe
O
IOF5
O
Xe
F
I
1+
F
F
O 1-
F
I
F
F
0
0
F
I
0
F
O
0
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Hypervalence:
Hypervalent molecules: species having more than an octet
around at least one atom
O
O
F
Xe
O
O
I
14
F
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VSEPR model (Valence Shell Electron Pair Repulsions)
Molecular geometry
The most stable arrangement of groups attached to a central atom
is the one that has the maximum separation of electron pairs
(bonded or non-bonded)
The properties of a compound are very much determined by the size
and shape of its molecules.
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Electron-pair geometry
Molecules in which the central atom has no lone pairs
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Molecules in which the central atom has one or more lone pairs
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Using the examples to illustrate the VSEPR model
SF4 , BrF3, XeF2, XeF4, OSF4
CH4
NH3
OH2
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Using the examples to illustrate the VSEPR model
SF4 , BrF3, XeF2, XeF4, OSF4
F
F
F
S
F
S
F
F
F
F
F
F
F
F
Br
Br
F
F
F Br
F
F
F
F
F
Xe
F
Xe
Xe
F
Xe
F
F
F
F
F
F
F
F
S
F
F
F
Xe
F
F
SF4 adopts a seesaw structure
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Guidelines for applying the VSEPR model
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Valence Bond Theory
describes covalent bond in terms of the overlap of atomic orbitals
Valence bond model of H2
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Hybridization -- Hybrid orbitals
Valence bond theory and molecular geometry
Hybrid orbitals are mixtures of atomic
orbitals with intermediate energy
The number of s, p and d orbitals in
the mixture equals the number of
hybrid orbitals
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Examples with an
sp3 hybridization
central atom
Examples:
PCl5, SF4, OSF4
Hybridization of s, p and d orbitals
Examples:
SF6, XeF4
Pentagonal bipyramidal
Geometry (sp3d3)
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IF7
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Hybridization in ethylene, acetylene, etc.
Concept of σ and π bonds
Bonds that result from σ overlap are called σ bonds.
Bonds that result from π overlap are called π bonds.
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Structure of
Ethylene
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Bond polarity and dipole moment
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Molecular Dipole Moments
Molecule must have polar bonds (necessary but not sufficient)
Need to know molecular shape because individual bond dipoles can cancel
Examples
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