Transcript Ch 11 - Pial

Chapter 11
Theories of Covalent Bonding
11-1
Theories of Covalent Bonding
11.1 Valence Bond (VB) Theory and Orbital Hybridization
11.2 The Mode of Orbital Overlap and the Types of
Covalent Bonds
11.3 Molecular Orbital (MO)Theory and Electron Delocalization
11-2
The Central Themes of VB Theory
Basic Principle
A covalent bond forms when the orbtials of two atoms overlap
and are occupied by a pair of electrons that have the highest
probability of being located between the nuclei.
Themes
A set of overlapping orbitals has a maximum of two electrons
that must have opposite spins.
The greater the orbital overlap, the stronger (more stable) the
bond.
The valence atomic orbitals in a molecule are different from
those in isolated atoms.
11-3
Atomic Orbital Overlap
Orbital overlap and
spin pairing in diatomic
molecules
Hydrogen, H2
Hydrogen fluoride, HF
11-4
Fluorine, F2
Hybrid Orbitals
Key Points
The number of hybrid orbitals obtained equals the number of
atomic orbitals mixed.
The type of hybrid orbitals obtained varies with the types of
atomic orbitals mixed.
Types of Hybrid Orbitals
sp
11-5
sp2
sp3
sp3d
sp3d2
The sp hybrid orbitals in gaseous BeCl2
atomic
orbitals
hybrid
orbitals
orbital box diagrams
11-6
The sp hybrid orbitals in gaseous BeCl2 (continued)
orbital box diagrams with orbital contours
11-7
The sp2 hybrid orbitals in BF3
11-8
The sp3 hybrid orbitals in CH4
11-9
The sp3 hybrid orbitals in NH3
11-10
The sp3 hybrid orbitals in H2O
11-11
The sp3d hybrid orbitals in PCl5
11-12
The sp3d2 hybrid orbitals in SF6
11-13
11-14
The conceptual steps from molecular formula to the
hybrid orbitals used in bonding.
Step 1
Molecular
formula
Step 2
Lewis
structure
Figure 10.1
11-15
Step 3
Molecular shape
and e- group
arrangement
Figure 10.12
Table 11.1
Hybrid
orbitals
SAMPLE PROBLEM 11.1
PROBLEM:
Postulating Hybrid Orbitals in a Molecule
Use partial orbital diagrams to describe mixing of atomic
orbitals on the central atoms leads to hybrid orbitals in each of
the following:
(a) Methanol, CH3OH
PLAN:
Use the Lewis structures to ascertain the arrangement of
groups and shape of each molecule. Postulate the hybrid
orbitals. Use partial orbital box diagrams to indicate the hybrid
for the central atoms.
SOLUTION:
H
(a) CH3OH
H
11-16
(b) Sulfur tetrafluoride, SF4
C O
H H
The groups around C are
arranged as a tetrahedron.
O also has a tetrahedral
arrangement with 2 nonbonding
e- pairs.
SAMPLE PROBLEM 11.1
Postulating Hybrid Orbitals in a Molecule
continued
2p
2s
2p
sp3
single C atom
hybridized
C atom
2s
sp3
hybridized
O atom
single O atom
(b) SF4 has a seesaw shape with 4 bonding and 1 nonbonding e- pairs.
F
F S
F
F
3d
3d
3p
sp3d
3s
11-17
S atom
hybridized
S atom
Types of Covalent Bonds
Sigma () Bonds - Bonding that results from the endto-end overlap is called a sigma bond. It has the
highest electron density along the axis between the two
nuclei. Single bonds are sigma bonds.
Pi () Bonds - Bonds that result from the side-to-side
overlap of unhybridized p orbitals. The electron density
is above and below the axis between the two nuclei.
(This is why multiple bonds counted as one group of
electrons in VSEPR theory)
The multiple part of multiple bonds are  bonds. In a
double bond, there is one  and one  bond.
In a triple bond, there is one  and two  bonds.
11-18
The  bonds in ethane.
both C are sp3 hybridized
s-sp3 overlaps to  bonds
sp3-sp3 overlap to form a  bond
relatively even
distribution of electron
density over all 
bonds
11-19
The  and  bonds in ethylene (C2H4)
overlap in one position - 
p overlap - 
electron density
11-20
The  and  bonds in acetylene (C2H2)
overlap in one position - 
p overlap - 
11-21
SAMPLE PROBLEM 11.2
PROBLEM:
PLAN:
Describing the Bonding in Molecules with
Multiple Bonds
Describe the types of bonds and orbitals in acetone, (CH3)2CO.
Use the Lewis structures to ascertain the arrangement of groups and
shape at each central atom. Postulate the hybrid orbitals taking note of
the multiple bonds and their orbital overlaps.
sp2
SOLUTION:
sp2
sp3 hybridized
O
sp3
hybridized H
C
sp
H
C
C
H H H H
sp2 hybridized
H
2
sp3
H sp
2
sp2 C sp
C
3
sp H
sp3
H sp3
sp3
sp3
H
3
sp 3 H
sp
bonds
11-22
O
2
O
C
C
H3 C
CH3
bond
Restricted rotation of -bonded molecules
Rotation about the C-C bond can’t take place without
breaking the  electron overlap.
CIS
11-23
TRANS
The Central Themes of MO Theory
A molecule is viewed on a quantum mechanical level as a
collection of nuclei surrounded by delocalized molecular orbitals.
Atomic wave functions are summed to obtain molecular wave
functions.
If wave functions reinforce each other, a bonding MO is formed
(region of high electron density exists between the nuclei).
If wave functions cancel each other, an antibonding MO (*) is
formed (a node of zero electron density occurs between the nuclei).
11-24
An analogy between light waves and atomic wave functions.
Amplitudes of wave
functions added
Amplitudes of wave
functions subtracted.
11-25
Contours and energies of the bonding and antibonding
molecular orbitals (MOs) in H2.
11-26
The MO diagram for H2
11-27
MO diagram for He2+ and He2
*1s
1s
1s
Energy
Energy
*1s
1s
1s
1s
AO of
He
MO of
He+
1s
AO of
He+
He2+ bond order = 1/2
(exists)
11-28
AO of
He
MO of
He2
AO of
He
He2 bond order = 0
SAMPLE PROBLEM 11.3
PROBLEM:
PLAN:
Predicting Species Stability Using MO Diagrams
Use MO diagrams to predict whether H2+ and H2- exist.
Determine their bond orders and electron configurations.
Use H2 as a model and accommodate the number of electrons in
bonding and antibonding orbitals. Find the bond order.
SOLUTION:

1s
bond order
= 1/2(1-0)
= 1/2

+
AO of H
1s
11-29
1s
AO of H-
AO of H


MO of H2+
H2- does exist
H2 does exist
1s
AO of H
bond order
= 1/2(2-1)
= 1/2
configuration is (1s)1
MO of H2-
configuration is
(1s)2(2s)1
*2s
*2s
2s
Energy
2s
Li2
2s
Bonding in s-block
homonuclear
diatomic
molecules.
1s
1s
11-30
2s
Be2
*1s
*1s
1s
2s
2s
Li2 bond order = 1
(is observed)
1s
1s
1s
Be2 bond order = 0
(not observed)
Contours and energies of s and p MOs through
combinations of 2p atomic orbitals
11-31
Relative MO energy levels for Period 2 homonuclear
diatomic molecules.
without 2s-2p
mixing
with 2s-2p
mixing
MO energy levels
for O2, F2, and Ne2
MO energy levels
for B2, C2, and N2
11-32
MO occupancy
and molecular
properties for B2
through Ne2
11-33
SAMPLE PROBLEM 11.4
PROBLEM:
Using MO Theory to Explain Bond Properties
As the following data show, removing an electron from N2 forms
an ion with a weaker, longer bond than in the parent molecules,
whereas the ion formed from O2 has a stronger, shorter bond:
N2
N2+
O2
O 2+
Bond energy (kJ/mol)
945
841
498
623
Bond length (pm)
110
112
121
112
Explain these facts with diagrams that show the sequence and occupancy of MOs.
PLAN:
Find the number of valence electrons for each species, draw the MO
diagrams, calculate bond orders, and then compare the results.
SOLUTION:
N2 has 10 valence electrons, so N2+ has 9.
O2 has 12 valence electrons, so O2+ has 11.
11-34
SAMPLE PROBLEM 11.4
Using MO Theory to Explain Bond Properties
continued
N2+
N2
bonding e- lost
1/2(8-2)=3
11-35
O2 +
O2
2p
2p
2p
2p
2p
2p
2p
2p
2s
2s
2s
2s
1/2(7-2)=2.5
bond
orders
1/2(8-4)=2
antibonding
e- lost
1/2(8-3)=2.5
The lowest energy -bonding MOs in benzene and ozone.
11-36
The MO diagram for HF
Energy

1s
2px 2py
2p

AO
of H
11-37
MO of
HF
AO
of F
Energy
*
2p
The MO diagram for NO
s
*
p

p

*
possible Lewis
structures
p
s
2s
2s
AO of N
AO of O

s
MO of NO
11-38
2p
0
0
N
O
-1
+1
N
O
End of Chapter 11
11-39
Figure 10.1
The steps in converting a molecular formula into a Lewis structure.
Molecular
formula
Step 1
Atom
placement
Place atom
with lowest
EN in center
Step 2
Sum of
valence e-
Add A-group
numbers
Step 3
Remaining
valence e-
Draw single bonds.
Subtract 2e- for each bond.
Step 4
Lewis
structure
11-40
Give each
atom 8e(2e- for H)
Figure 10.12
Molecular
formula
The steps in determining a molecular shape.
Step 1
Lewis
structure
See Figure
10.1
Step 2
Electron-group
arrangement
Count all e- groups around central
atom (A)
Step 3
Bond
angles
Note lone pairs and double
bonds
Count bonding and
Step 4
nonbonding egroups separately.
Molecular
shape
(AXmEn)
11-41