Chapter Nine:
COVALENT BONDING:
ORBITALS
Assignment
• 1-85題中每5題裡任選1-2題
Chapter 9 | Slide 2
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Exercise
• Draw the Lewis structure for methane, CH4.
• What is the shape of a methane molecule?
• What are the bond angles?
9.1
Chapter 9 | Slide 3
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Concept Check
• What is the valence electron configuration of a
carbon atom?
• Why can’t the bonding orbitals for methane be
formed by an overlap of atomic orbitals?
9.1
Chapter 9 | Slide 4
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Bonding in Methane
• Assume that the carbon atom has four
equivalent atomic orbitals, arranged
tetrahedrally.
9.1
Chapter 9 | Slide 5
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Hybridization
• Mixing of the native atomic orbitals to
form special orbitals for bonding.
9.1
Chapter 9 | Slide 6
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sp3 Hybridization
• Combination of one s and three p orbitals.
• Whenever a set of equivalent tetrahedral
atomic orbitals is required by an atom,
the localized electron model assumes
that the atom adopts a set of sp3 orbitals;
the atom becomes sp3 hybridized.
9.1
Chapter 9 | Slide 7
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An Energy-Level Diagram Showing the
Formation of Four sp3 Orbitals
9.1
Chapter 9 | Slide 8
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The Formation of sp3 Hybrid Orbitals
9.1
Chapter 9 | Slide 9
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Tetrahedral Set of Four sp3 Orbitals
9.1
Chapter 9 | Slide 10
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Exercise
• Draw the Lewis structure for O2.
• What is the shape of an oxygen molecule?
• What is the approximate angle between lone
pairs of electrons on each of the oxygen atoms?
9.1
Chapter 9 | Slide 11
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Concept Check
• Why can’t sp3 hybridization account for the
oxygen molecule?
9.1
Chapter 9 | Slide 12
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sp2 Hybridization
• Combination of one s and two p orbitals.
• Gives a trigonal planar arrangement of
atomic orbitals.
• One p orbital is not used.
– Oriented perpendicular to the plane of the
sp2 orbitals
9.1
Chapter 9 | Slide 13
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Sigma (σ) Bond
• Electron pair is shared in an area
centered on a line running between the
atoms.
9.1
Chapter 9 | Slide 14
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Pi (π) Bond
• Forms double and triple bonds by
sharing electron pair(s) in the space
above and below the σ bond.
• Uses the unhybridized p orbitals.
9.1
Chapter 9 | Slide 15
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An Orbital Energy-Level Diagram for
sp2 Hybridization
9.1
Chapter 9 | Slide 16
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The Hybridization of the s, px, and py
Atomic Orbitals
9.1
Chapter 9 | Slide 17
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Formation of C=C Double Bond in
Ethylene
9.1
Chapter 9 | Slide 18
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Exercise
• Draw the Lewis structure for CO2.
• What is the shape of a carbon dioxide molecule?
• What are the bond angles?
9.1
Chapter 9 | Slide 19
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sp Hybridization
• Combination of one s and one p orbital.
• Gives a linear arrangement of atomic
orbitals.
• Two p orbitals are not used.
– Needed to form the π bonds
9.1
Chapter 9 | Slide 20
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The Orbital Energy-Level Diagram for
the Formation of sp Hybrid Orbitals on
Carbon
9.1
Chapter 9 | Slide 21
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When One s Orbital and One p Orbital
are Hybridized, a Set of Two sp Orbitals
Oriented at 180 Degrees Results
9.1
Chapter 9 | Slide 22
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The Orbitals for CO2
9.1
Chapter 9 | Slide 23
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Exercise
• Draw the Lewis structure for PCl5.
• What is the shape of a phosphorus
pentachloride molecule?
• What are the bond angles?
9.1
Chapter 9 | Slide 24
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dsp3 Hybridization
• Combination of one d, one s, and three p
orbitals.
• Gives a trigonal bipyramidal arrangement
of five equivalent hybrid orbitals.
9.1
Chapter 9 | Slide 25
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The Orbitals Used to Form the Bonds in
PCl5
9.1
Chapter 9 | Slide 26
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Exercise
• Draw the Lewis structure for XeF4.
• What is the shape of a xenon tetrafluoride
molecule?
• What are the bond angles?
9.1
Chapter 9 | Slide 27
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d2sp3 Hybridization
• Combination of two d, one s, and three p
orbitals.
• Gives an octahedral arrangement of six
equivalent hybrid orbitals.
9.1
Chapter 9 | Slide 28
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How is the Xenon Atom in XeF4
Hybridized?
9.1
Chapter 9 | Slide 29
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Concept Check
• Draw the Lewis structure for HCN.
• Which hybrid orbitals are used?
• Draw HCN:
– Showing all bonds between atoms
– Labeling each bond as  or 
9.1
Chapter 9 | Slide 30
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Concept Check
Determine the bond angle and expected
hybridization of the central atom for each of the
following molecules:
NH3
SO2
KrF2
CO2
ICl5
9.1
Chapter 9 | Slide 31
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The Localized Electron Model
• Draw the Lewis structure(s).
• Determine the arrangement of electron
pairs (VSEPR model).
• Specify the necessary hybrid orbitals.
9.1
Chapter 9 | Slide 32
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Molecular Orbital Model
• The electron probability of both molecular
orbitals is centered along the line passing
through the two nuclei.
– Sigma (σ) molecular orbitals
• In the molecule only the molecular orbitals
are available for occupation by electrons.
9.2
Chapter 9 | Slide 33
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Molecular Orbital Model (continued)
• MO1 is lower in energy than the s orbitals
of free atoms, while MO2 is higher in
energy than the s orbitals.
– Bonding molecular orbital – lower in energy
– Antibonding molecular orbital – higher in
energy
9.2
Chapter 9 | Slide 34
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Molecular Orbital Model (continued)
• The molecular orbital model produces
electron distributions and energies that
agree with our basic ideas of bonding.
• The labels on molecular orbitals indicate
their symmetry (shape), the parent atomic
orbitals, and whether they are bonding or
antibonding.
9.2
Chapter 9 | Slide 35
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Molecular Orbital Model (continued)
• Molecular electron configurations can be
written similar to atomic electron
configurations.
• Each molecular orbital can hold 2
electrons with opposite spins.
• Orbitals are conserved.
9.2
Chapter 9 | Slide 36
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Bonding in H2
9.2
Chapter 9 | Slide 37
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Sigma Bonding and Antibonding
Orbitals
9.2
Chapter 9 | Slide 38
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Bond Order
• Larger bond order means greater bond
strength.
B.O. = (# of bonding e- – # of antibonding e-)/2
9.2
Chapter 9 | Slide 39
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Homonuclear Diatomic Molecules
• Composed of 2 identical atoms.
• Only the valence orbitals of the atoms
contribute significantly to the molecular
orbitals of a particular molecule.
9.3
Chapter 9 | Slide 40
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Pi Bonding and Antibonding Orbitals
9.3
Chapter 9 | Slide 41
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Magnetic Properties of Liquid Nitrogen
and Oxygen
Chapter 9 | Slide 42
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Paramagnetism
• Paramagnetism – substance is attracted
into the inducing magnetic field.
– Unpaired electrons (O2)
• Diamagnetism – substance is repelled
from the inducing magnetic field.
– Paired electrons (N2)
9.3
Chapter 9 | Slide 43
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Apparatus
Used to
Measure the
Paramagnetism
of a Sample
Chapter 9 | Slide 44
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Molecular Orbital Summary of Second
Row Diatomic Molecules
9.3
Chapter 9 | Slide 45
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Heteronuclear Diatomic Molecules
• Composed of 2 different atoms.
9.4
Chapter 9 | Slide 46
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Orbital Energy-Level Diagram for the
HF Molecule
9.4
Chapter 9 | Slide 47
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The Electron Probability Distribution in the
Bonding Molecular Orbital of the HF
Molecule
9.4
Chapter 9 | Slide 48
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The Sigma System for Benzene
9.5
Chapter 9 | Slide 49
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The Pi System for Benzene
9.5
Chapter 9 | Slide 50
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Pi Bonding in the Nitrate Ion
9.5
Chapter 9 | Slide 51
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