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PowerPoint Lecture Presentation
by
J. David Robertson
University of Missouri
Chemical Bonding II:
Molecular Geometry and
Hybridization of Atomic
Orbitals
Chapter 10
Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
10.1
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
B
B
10.1
MODEL SET COLORS
BALL COLOR
• BLUE & BLACK
= ATOM TYPE
= NON-METALS
• YELLOW/Red/GREEN = GROUP 1A
• RED (IIA)
= GROUP 11A metals
• RED (IIIA)
= GROUP 111A metals
HOMEWORK = PREPARE FOR EXAM!!
Cl
Be
Cl
lone pairs
on to
central
atom
20
atoms
bonded
central
atom
10.1
VSEPR
Class
# of atoms
bonded to
central atom
# lone
pairs on
central atom
Arrangement of
electron pairs
Molecular
Geometry
AB2
2
0
linear
linear
0
trigonal
planar
trigonal
planar
AB3
3
10.1
10.1
VSEPR
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
tetrahedral
AB3
3
0
trigonal
planar
AB4
4
0
tetrahedral
10.1
10.1
VSEPR
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
10.1
10.1
VSEPR
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
10.1
10.1
10.1
lone-pair vs. lone pair
lone-pair vs. bonding
bonding-pair vs. bonding
>
>
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
10.1
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
10.1
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
10.1
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
10.1
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
10.1
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
10.1
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
10.1
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
10.1
10.1
Bond Polarity and Polar Molecules
electron poor
region
electron rich
region
H
F
d+
d-
Predicting Molecular Geometry
1. Draw Lewis structure for molecule.
2. Count number of electron pairs (bonding & lone) 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 CF4?
F
O
S
O
F
C
F
F
S
F
C
F
F
F
Bond Polarity and Polar Molecules
Polarity refers to have two oppositely charged ends
(aka: dipoles)
Bond Polarity refers to a specific BOND that has two
oppositely charged ends
H +– - F
or
O -–+ C
Bond Polarity and Polar Molecules
Molecular Polarity refers to an ENTIRE MOLECULE
having two oppositely charged ends
Bond Polarity and Polar Molecules
Polar Molecules NEED Polar Bonds!
HOWEVER
Non-Polar Molecules may have either Polar or
Non-Polar bonds
Three types of bonds exist between atoms:
1. Non-Polar Covalent
2. Polar Covalent
3. Ionic
Bond Polarity and Polar Molecules
1. Non-Polar Covalent Bond
Atoms Share Electrons Equally….
H – H
Each atoms pulls equally hard
on the bonding electrons
Bond is essentially neutral…no + or – ends!!
Bond Polarity and Polar Molecules
2 & 3. Polar Covalent or IONIC Bond
Atoms DO NOTShare Electrons Equally….
H +– - F
Fluorine atoms pulls harder on
the bonding electrons
H
F
Bond Polarity and Polar Molecules
PREDICTING MOLECULAR POLARITY:
STEPS:
1. Draw Lewis Structure as REALISTICALLY as possible.
2. Calculate the ELECTRONEGATIVITY DIFFERENCE between
the atoms making the bonds.
3. Determine the TYPES OF BOND using the following scale:
0
0.4
Non-polar covalent
2.0
Polar covalent
Ionic
4. In your Lewis Structure draw in the DIRECTION OF
ELECTRON SHIFT within the bonds (IF POLAR OR IONIC).
5. Make a judgment (based off symmetry) if the molecule will
have a dipole.
More HW Answers:
Part A
β = 105
φ = 119.5
 = 109.5
 = 107
 = 120
 = 180
Part B
Ga
Ba
P
N
C
He
Si
Br
Cs
Al
Part D
1. Lose electrons
2. Valence electrons, they’re involved in bonding
More HW Answers:
Part D
3. Sharing of valence electrons…polar: e- pulled
to one side, non-polar: e- shared equally.
4. Triple: more shell orbitals need to overlap, so
atoms move closer together.
5. Triple: more energy required to hold nuclei
close together….well compressed spring.
6. EN ≥ 2.0
7. Already have filled outer-most shell
Part D
More HW Answers:
8. Polar molecules need opposite charges on
opposite ends of itself
9. Because + & - ends act like N & S poles
10. In order to give away a valence electron they
need to make a bond
Part F:
1. Octet, noble
2. Ionic
3. Cation, Anion
Part F
More HW Answers:
4. Sharing
5. NON-POLAR
6. IONIC
7. PULL or HOLD
8. IONS
9. UNEVENLY
10. ZERO
Part G
More HW Answers:
1. Al2S3
6. NH4Cl
2. Al2(SO4)3
7. MgSO4·5H2O
3. Al2(SO3)3
8. Fe2O3
4. MgO
9. Na3PO4
5. Ca(CN)2
10. KOH
Part H:
1. S2O7
4. P3O4
2. CO2
5. H2O
3. N2O
6. P2O5
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
10.2
EXAMPLE: HCl
STEPS:
1. Draw Lewis Structure as REALISTIC as possible.
2. Calculate the ELECTRONEGATIVITY DIFFERENCE between
the atoms making the bonds.
3. Determine the TYPES OF BOND using the following scale:
0
0.4
2.0
4. In your Lewis Structure draw in (if necessary) the DIRECTION
OF ELECTRON SHIFT within the bonds.
5. Make a judgment (based off symmetry) if the molecule will
have a dipole.
10.2
Does CH2Cl2 have
a dipole moment?
10.2
10.2
Chemistry In Action: Microwave Ovens
10.2
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.
10.3
10.4
Change in electron
density as two hydrogen
atoms approach each
other.
10.3
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
10.4
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
10.4
10.4
10.4
Predict correct
bond angle
10.4
Formation of sp Hybrid Orbitals
10.4
Formation of sp2 Hybrid Orbitals
10.4
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
10.4
10.4
10.5
10.5
Pi bond (p) – electron density above and below plane of nuclei
Sigma bond (s) – electron density between the 2 atoms
of the bonding atoms
10.5
10.5
10.5
10.5
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
O
H
C
O
H
s bonds = 6 + 1 = 7
p bonds = 1
H
10.5
Experiments show O2 is paramagnetic
O
O
No unpaired e-
Should be diamagnetic
Molecular orbital theory – bonds are formed from
interaction of atomic orbitals to form molecular
orbitals.
10.6
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
formed.
10.6
10.6
10.6
10.6
10.6
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.
10.7
1
bond order =
2
bond
order
½
(
Number of
electrons in
bonding
MOs
1
-
½
Number of
electrons in
antibonding
MOs
)
0
10.7
10.7
Delocalized molecular orbitals are not confined between
two adjacent bonding atoms, but actually extend over three
or more atoms.
10.8
Electron density above and below the plane of the
benzene molecule.
10.8
10.8
Chemistry In Action: Buckyball Anyone?
10.8