Transcript No Slide Title - HCC Learning Web

Chemical Bonding II:
Molecular Geometry and
Hybridization of Atomic Orbitals
Chapter 10
Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
Review
Valence Shell Electron Pair Repulsion
(VSEPR) Theory
Valence Shell Electrons
• the outer shell electrons of an atom, the ones involved in bonding
• for a given elements, # of valence electrons = Group Number
• C: Group 4A, 4 valence electrons
O: Group 6A, 6 valence electrons
Draw Lewis structure (pre-requisition)
NH3
H
N
H
H
– Key: Electrons are all negatively charged
– Action: electron pairs (bonding pairs & lone pairs ) around the central
atom repel each other to keep themselves as far away as possible
– Outcome: Maximum separation
Minimum repulsion
Symmetry (equal repulsion)
– Applications: predict the electron pair geometry
predict the bond angle
predict the molecular geometry
predict the hybridization of the central atom
predict the polarity of the molecule
– Electron pair geometry : the arrangement of all electron pairs (bonding pairs &
lone pairs ) surrounding the central atom of the molecule.
– Molecular geometry: the arrangement of only bonding pairs surrounding the
central atom of the molecule
VSEPR
Class
AB2
# of bonding
pairs
2
# of lone
pairs
0
Total # of
electron pairs
Electron pair
geometry
Molecular
Geometry
linear
linear
2
180o
B
A
B
B
A
B
VSEPR
Class
# of bonding
pairs
# of
lone pairs
Total # of
electron pairs
Electron pair
geometry
Molecular
Geometry
trigonal planar
trigonal planar
B
B
AB3
3
0
3
120o
A
A
B
B
B
trigonal planar
AB2E
2
1
B
bent
3
A
A
B
B
B
B
VSEPR
Class
# of bonding
pairs
# of
lone pairs
Total # of
electron pairs
Electron pair
geometry
Molecular
Geometry
tetrahedral
tetrahedral
B
AB4
4
0
4
A
B
B
B
tetrahedral
AB3E
3
1
A
2
2
A
B
B
B
B
trigonal pyramidal
4
B
AB2E2
B
109.5o
A
B
B
B
tetrahedral
bent
A
A
4
B
B
B
B
B
Class
AB5
# of bonding
pairs
5
Total # of
# of
lone pairs electron pairs
0
5
Electron pair
geometry
Molecular
Geometry
trigonal
bipyramidal
trigonal
bipyramidal
120
B
90o
AB4E
4
1
B
B
o
B
A
B
B
A
B
B
B
B
trigonal
bipyramidal
seesaw
B
B
5
B
B
A
A
B
B
AB3E2
AB2E3
3
2
B
B
trigonal
bipyramidal
T-shaped
B
B
5
B
2
3
5
A
B
A
B
B
trigonal
bipyramidal
linear
B
B
A
A
B
B
VSEPR
Class
AB6
# of bonding
pairs
6
# of
lone pairs
0
Total # of
electron pairs
6
Electron pair
geometry
Molecular
Geometry
octahedral
octahedral
B
B
B
90o
B
A
B
B
B
B
B
square pyramidal
B
AB5E
5
1
6
B
A
B
B
octahedral
AB4E2
4
2
6
B
B
B
B
octahedral
B
B
A
B
A
B
B
B
B
A
B
B
square planar
B
B
B
A
B
Bond angle
Periodic table
Electron pair
geometry
# of valance shell electrons
Lewis
structure
Molecular
geometry
Total # of electron pairs
around central atom
Hybridization of
the central atom
What is the electron pair geometry, bond angle, molecular geometry and
hybridization of the central atom of NH3?
electron pair geometry: tetrahedral
H
N
H
bond angle: 109.5o
molecular geometry: trigonal pyramidal
H
hybridization: sp3
Repulsive force
lone-pair vs. lone pair
lone-pair vs. bonding
bonding-pair vs. bonding
>
>
repulsion
pair repulsion
pair repulsion
Predicting Polarity of Molecules:
Dipole Moments and Polar Molecules
m=Qxr
Q is the charge
r is the distance between charges electron poor
region
1 D = 3.36 x 10-30 C m
H
d+
• A measure of the polarity of a molecule
• The dipole moment can be determined
experimentally
• Measure in Debye units
• Predict polarity by taking the vector sum of
the bond dipoles
electron rich
region
F
d-
10.2
Molecules containing net dipole moments are called polar
molecules. Otherwise they are called nonpolar molecules
because they do not have net dipole moments.
Diatomic molecules: Determined by the polarity of bond
Polar molecules:HCl, CO,NO
Nonpolar molecules: H2,F2,O2
Molecules with three or more atoms: determined by the
polarity of the bond and the molecular geometry
Dipole moment is a vector quantity, which has both
magnitude and direction.
• Symmetric molecules such as these are nonpolar because
the bond dipoles cancel
• All of the basic shapes are symmetric, or balanced, if all the
domains and groups attached to them are identical
C
O
C
O
Linear molecule
Nodipolar moment
O
O
bent molecule
Net dipolar moment
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
• Lewis structures and VSEPR do not tell us why
electrons group into domains as they do
• How atoms form covalent bonds in molecules requires
an understanding of how orbitals interact
Molecular Geometry & Hybridization
Covalent Bonding Theories
• Valence bond (VB) theory
– Bonding is an overlap of atomic orbitals
• Includes overlap of hybrid atomic orbitals
• Molecular orbital (MO) theory
– Bonding happens when molecular orbitals are formed
10.3
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
Predict the Hybridization of the Central Atom
Total # of
electron pairs
Electron pair
geometry
Hybridization
2
linear
sp
BeCl2
3
trigonal planar
sp2
BF3
4
tetrahedral
sp3
5
trigonal bipyramidal
sp3d
PCl5
6
octahedral
sp3d2
SF6
Examples
CH4, NH3, H2O
Total # of electron pairs = number of hybrid orbital (sum of the superscripts!)
Formation of sp3 Hybrid Orbitals
10.4
Formation of sp Hybrid Orbitals
10.4
Hybridization in molecules containing double and triple bonds
Two Kinds of covalent Bonds
• Sigma bond: bonding density is along the internuclear axis
– head to head overlap of two hybrid orbitals
– head to head overlap of p + hybrid
• any s character to the bond = sigma bond
• Pi bond: bonding density is above and below the
internuclear axis
– Sideway overlap p + p
10.5
Sigma bond
Pi bond
Multiple bonds
• Double bond: One sigma and one pi
bond between the same atoms
–Example: ethylene
• Triple bond: One sigma and two pi
bonds between the same atoms
–Example: acetylene
10.5
Review of 3 Types of bonds:
1. Sigma bond
e-density (overlap region) is along the internuclear axis
2. Pi bond
e-density (overlap region) is above and below the plane of
the internuclear axis.
Relative reactivity of bonds:
• Pi bonds are more reactive than sigma bonds
• Less energy is needed to break a pi bond than a sigma bond
3. Delocalized Bonds
• are not confined between two adjacent bonding atoms, but
actually extend over three or more atoms.
• less reactive than normal pi bonds
• Example: benzene
Summary: VB Theory
framework of the molecule is determined by the arrangement of
the sigma-bonds; Hybrid orbitals are used to form the sigma
bonds and the lone pairs of electrons.
Sigma (s) and Pi Bonds (p)
Single bond
1 sigma bond
1 sigma bond and 1 pi bond
Double 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
Valence bond theory
O
O
Experiments show O2 is paramagnetic
No unpaired e-
Should be diamagnetic
But experiment show
that there are two
unpaired electrons—
paramagnetic.
Molecular orbital theory – bonds are formed from
interaction of atomic orbitals to form molecular
orbitals.
10.6
Molecular Orbital (MO) Theory
• Bonds are formed from interaction of atomic orbitals to form
molecular orbitals.
MO theory takes the view that a molecule is similar to an atom
• The molecule has molecular orbitals that can be populated by
electrons just like the atomic orbitals in atoms
• # of MOs formed = # of atomic orbitals combined
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
Molecular Orbital (MO) Configurations
-MOs follow the same filling rules as atomic orbitals
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.(Electrons spread out as much as possible,
with spins unpaired, over orbitals that have 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