Transcript No Slide Title

Covalent Bonding:
Molecular Geometry
Hybridization of Atomic Orbitals
Molecular Orbitals
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2
Valence shell electron pair repulsion (VSEPR) model:
Predict the geometry of the molecule from the electrostatic
repulsions between the electron (bonding and nonbonding) pairs.
Class
# of atoms
bonded to
central atom
# lone
pairs on
central atom
Arrangement of
electron pairs
Molecular
Geometry
AB2
2
0
linear
linear
trigonal
planar
AB3
3
0
trigonal
planar
AB4
4
0
tetrahedral
tetrahedral
AB5
5
0
trigonal
bipyramidal
trigonal
bipyramidal
AB6
6
0
octahedral
octahedral
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VSEPR
linear
trigonal bipyramidal
tetrahedral
trigonal planar
octahedral
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VSEPR
Cl
Be
Cl
linear
trigonal bipyramidal
PCl5
BF3
CH4
tetrahedral
trigonal planar
octahedral
SF6
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Effects of Lone Pairs
H
H
H
C
H
N
H
H
O
H
H
H
lone-pair vs. lone pair
lone-pair vs. bonding
bonding-pair vs. bonding
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>
>
repulsion
pair repulsion
pair repulsion
VSEPR
Class
# of atoms
bonded to
central atom
# lone
pairs on
central atom
AB3
3
0
AB2E
2
1
Arrangement of
electron pairs
Molecular
Geometry
trigonal
planar
trigonal
planar
trigonal
planar
bent
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VSEPR
Class
# of atoms
bonded to
central atom
# lone
pairs on
central atom
AB4
4
0
AB3E
3
1
Arrangement of
electron pairs
Molecular
Geometry
tetrahedral
tetrahedral
tetrahedral
trigonal
pyramidal
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VSEPR
Class
# of atoms
bonded to
central atom
# lone
pairs on
central atom
AB4
4
0
Arrangement of
electron pairs
Molecular
Geometry
tetrahedral
tetrahedral
AB3E
3
1
tetrahedral
trigonal
pyramidal
AB2E2
2
2
tetrahedral
bent
O
H
H
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VSEPR
Class
AB5
AB4E
# of atoms
bonded to
central atom
5
4
# lone
pairs on
central atom
Arrangement of
electron pairs
Molecular
Geometry
0
trigonal
bipyramidal
trigonal
bipyramidal
1
trigonal
bipyramidal
distorted
tetrahedron
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VSEPR
Class
AB5
# of atoms
bonded to
central atom
5
# lone
pairs on
central atom
0
AB4E
4
1
AB3E2
3
2
Arrangement of
electron pairs
Molecular
Geometry
trigonal
bipyramidal
trigonal
bipyramidal
trigonal
bipyramidal
trigonal
bipyramidal
distorted
tetrahedron
T-shaped
F
F
Cl
F
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VSEPR
Class
AB5
# of atoms
bonded to
central atom
5
# lone
pairs on
central atom
0
AB4E
4
1
AB3E2
3
2
AB2E3
2
3
Arrangement of
electron pairs
Molecular
Geometry
trigonal
bipyramidal
trigonal
bipyramidal
trigonal
bipyramidal
trigonal
bipyramidal
distorted
tetrahedron
trigonal
bipyramidal
T-shaped
linear
I
I
I
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VSEPR
Class
# of atoms
bonded to
central atom
# lone
pairs on
central atom
AB6
6
0
octahedral
octahedral
AB5E
5
1
octahedral
square
pyramidal
F
F
F
Arrangement of
electron pairs
Molecular
Geometry
Br
F
F
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VSEPR
Class
# of atoms
bonded to
central atom
# lone
pairs on
central atom
AB6
6
0
octahedral
octahedral
AB5E
5
1
octahedral
AB4E2
4
2
octahedral
square
pyramidal
square
planar
Arrangement of
electron pairs
Molecular
Geometry
F
F
Xe
F
F
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Predicting Molecular Geometry
1. Draw Lewis structure for molecule.
2. Count number of lone pairs on the central atom and
number of atoms bonded to the central atom.
3. Use VSEPR to predict the geometry of the molecule.
What are the molecular geometries of SO2 and SF4?
O
S
AB2E
bent
F
O
F
S
F
AB4E
F
distorted
tetrahedron
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Dipole Moments and Polar Molecules
electron poor
region
electron rich
region
H
F
d+
d-
m=Qxr
Q is the charge
r is the distance between charges
1 D = 3.36 x 10-30 C m
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Which of the following molecules have a dipole moment?
H2O, CO2, SO2, and CH4
O
S
dipole moment
polar molecule
dipole moment
polar molecule
H
O
C
O
no dipole moment
nonpolar molecule
H
C
H
H
no dipole moment
nonpolar molecule
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Does CH2Cl2 have
a dipole moment?
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Chemistry In Action: Microwave Ovens
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How does Lewis theory explain the bonds in H2 and F2?
Sharing of two electrons between the two atoms.
Overlap Of
Bond Dissociation Energy
Bond Length
H2
436.4 kJ/mole
74 pm
2 1s
F2
150.6 kJ/mole
142 pm
2 2p
Valence bond theory – bonds are formed by sharing
of e- from overlapping atomic orbitals.
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Change in electron
density as two hydrogen
atoms approach each
other.
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Valence Bond Theory and NH3
N – 1s22s22p3
3 H – 1s1
If the bonds form from overlap of 3 2p orbitals on nitrogen
with the 1s orbital on each hydrogen atom, what would
the molecular geometry of NH3 be?
If use the
3 2p orbitals
predict 900
Actual H-N-H
bond angle is
107.30
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Hybridization – mixing of two or more atomic
orbitals to form a new set of hybrid orbitals.
1. Mix at least 2 nonequivalent atomic orbitals (e.g. s
and p). Hybrid orbitals have very different shape
from original atomic orbitals.
2. Number of hybrid orbitals is equal to number of
pure atomic orbitals used in the hybridization
process.
3. Covalent bonds are formed by:
a. Overlap of hybrid orbitals with atomic orbitals
b. Overlap of hybrid orbitals with other hybrid
orbitals
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Bonding in Methane
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Fig. 10.7
Formation of sp3 Hybrid Orbitals
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Formation of sp3 Hybrid Orbitals
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Formation of a CH4 Molecule
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Formation of a NH3 Molecule
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Stylized Drawing of Valence Bond Theory
Predict correct
bond angle
CH4
NH3
Sigma bond (s) – electron density between the 2 atoms
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Formation of sp2 Hybrid Orbitals
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Formation of sp2 Hybrid Orbitals
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Formation of sp2 Hybrid Orbitals
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2pz orbital is perpendicular to the plane of hybridized orbitals
sp2 Hybridization of a C atom
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Bonding in Ethylene C2H4
H
H
C
H
C
H
Sigma bond (s) – electron density between the 2 atoms
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Bonding in Ethylene C2H4
Pi bond (p) – electron density above and below plane of nuclei
of the bonding atoms
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Bonding in Ethylene C2H4
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Formation of sp Hybrid Orbitals
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Formation of sp Hybrid Orbitals
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Formation of sp Hybrid Orbitals
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Bonding in acetylene
C 2H 2
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How do I predict the hybridization of the central atom?
Count the number of lone pairs AND the number
of atoms bonded to the central atom
# of Lone Pairs
+
# of Bonded Atoms
Hybridization
Examples
2
sp
BeCl2
3
sp2
BF3
4
sp3
CH4, NH3, H2O
5
sp3d
PCl5
6
sp3d2
SF6
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Sigma (s) and Pi Bonds (p)
1 sigma bond
Single bond
Double bond
1 sigma bond and 1 pi bond
Triple bond
1 sigma bond and 2 pi bonds
How many s and p bonds are in the acetic acid
(vinegar) molecule CH3COOH?
H
C
H
O
H
C
O
H
s bonds = 6 + 1 = 7
p bonds = 1
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Drawback of Valence Bond Theory
Experiments show O2 is paramagnetic
O
O
No unpaired eShould be diamagnetic
Molecular orbital theory – bonds are formed from
interaction of atomic orbitals to form molecular
orbitals.
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An analogy between light waves and atomic
wave functions
Amplitudes of wave
functions added
Amplitudes of
wave functions
subtracted.
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Energy levels of bonding and antibonding
molecular orbitals in hydrogen (H2)
A bonding molecular orbital has lower energy and greater
stability than the atomic orbitals from which it was formed.
An antibonding molecular orbital has higher energy and
lower stability than the atomic orbitals from which it was
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formed.
Energy levels of bonding and antibonding
molecular orbitals in boron (B2)
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Second-Period Homonuclear Diatomic Molecules
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Molecular Orbital (MO) Configurations
1. The number of molecular orbitals (MOs) formed is always
equal to the number of atomic orbitals combined.
2. The more stable the bonding MO, the less stable the
corresponding antibonding MO.
3. The filling of MOs proceeds from low to high energies.
4. Each MO can accommodate up to two electrons.
5. Use Hund’s rule when adding electrons to MOs of the
same energy.
6. The number of electrons in the MOs is equal to the sum of
all the electrons on the bonding atoms.
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Bond Order
1
bond order =
2
bond
order
½
(
Number of
electrons in
bonding
MOs
-
1
½
Number of
electrons in
antibonding
MOs
)
0
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MO for 2nd Period Homonuclear Diatomic Molecules
MO theory predicts that O2 is paramagnetic!
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Molecules with Resonance Structures
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Delocalized π Molecular Orbitals
Delocalized molecular orbitals are not confined between
two adjacent bonding atoms, but actually extend over three
or more atoms.
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Electron density above and below the plane of the
benzene molecule.
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Acknowledgment
Some images, animation, and material have been taken from the following sources:
Chemistry, Zumdahl, Steven S.; Zumdahl, Susan A.; Houghton Mifflin Co., 6th Ed., 2003;
supplements for the instructor
General Chemistry: The Essential Concepts, Chang, Raymon; McGraw-Hill Co. Inc., 4th
Ed., 2005; supplements for the instructor
Principles of General Chemistry, Silberberg, Martin; McGraw-Hill Co. Inc., 1st Ed., 2006;
supplements for the instructor
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