11-1 Chapter 11 Theories of Covalent Bonding. 11-2 Theories of Covalent Bonding 11.1 Valence bond...
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Transcript of 11-1 Chapter 11 Theories of Covalent Bonding. 11-2 Theories of Covalent Bonding 11.1 Valence bond...
11-1
Chapter 11
Theories of Covalent Bonding
11-2
Theories of Covalent Bonding
11.1 Valence bond (VB) theory and orbital hybridization
11.2 The mode of orbital overlap and types of covalent bonds
11.3 Molecular orbital (MO) theory and electron delocalization
11-3
Figure 9.2
The three models of chemical bonding
11-4
Figure 9.11
Covalent bond formation in H2
11-5
Structure dictates shape
Shape dictates function
Key Principles
shape = conformation
Molecules can assume more than oneshape (conformation) in solution!
11-6
The Complementary Shapes of an Enzyme and Its Substrate
11-7
Valence-shell Electron-Pair Repulsion (VSEPR) Theory
A method to predict the shapes of molecules from their electronic structures (Lewis structures do not depictshape)
Basic principle: each group of valence electrons around a centralatom is located as far away as possible from the others in order to minimize repulsions
Both bonding and non-bonding valence electrons aroundthe central atom are considered.
AXmEn symbolism: A = central atom, X = surrounding atoms,E = non-bonding electrons (usually a lone pair)
11-8
Figure 8.12
A periodic table of partial ground-state electron configurations
11-9
Figure 10.12
The steps in determining a molecular shape
molecular formula
Lewis structure
electron-group arrangement
bond angles
molecular shape
(AXmEn)
Count all e- groups around the central atom A
Note lone pairs and double bonds
Count bonding and
non-bonding e- groups separately.
Step 1
Step 2
Step 3
Step 4
11-10
Figure 10.1
Steps to convert a molecular formula into a Lewis structure
molecular formula
atom placement
sum of
valence e-
remaining
valence e-
Lewis structure
Place the atom with the lowest EN in the center
Add A-group numbers
Draw single bonds and
subtract 2e- for each bond
Give each
atom 8e-
(2e- for H)
Step 1
Step 2
Step 3
Step 4
11-11
Figure 10.5
Electron-group repulsions and the five basic molecular shapes
Ideal bond angles are shown for each shape.
11-12
Figure 10.8
The three molecular shapes of the tetrahedral electron-group arrangement
Examples:
CH4, SiCl4,
SO42-, ClO4
-
Examples:NH3
PF3
ClO3
H3O+
Examples:
H2O
OF2
SCl2
11-13
Figure 10.10
The four molecular shapes of the trigonal bipyramidal electron-group arrangement
Examples:
SF4
XeO2F2
IF4+
IO2F2-
Examples:
ClF3
BrF3
Examples:
XeF2
I3-
IF2-
Examples:
PF5
AsF5
SOF4
11-14
VSEPR (Valence Shell Electron Pair RepulsionTheory)
Accounts for molecular shapes by assuming that electron groups tend to minimize their repulsions
Does not show how shapes can be explained fromthe interactions of atomic orbitals
11-15
The Central Themes of Valence Bond (VB) Theory
Basic Principle
A covalent bond forms when the orbitals of two atoms overlap and are occupied by a pair of electrons that have the highest probability of being located between the nuclei.
Three Central 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 (hybridization).
11-16
Figure 11.1
Orbital overlap and spin pairing in
three diatomic molecules
hydrogen, H2
hydrogen fluoride, HF
fluorine, F2
11-17
Linus Pauling
Proposed that valence atomic orbitals in the molecule are different from those in the isolated atoms
Mixing of certain combinations of atomic orbitalsgenerates new atomic orbitals
Process of orbital mixing = hybridization; generateshybrid orbitals
11-18
Hybrid Orbitals
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.
Key Points
sp sp2 sp3 sp3d sp3d2
Types of Hybrid Orbitals
11-19
Figure 11.2
The sp hybrid orbitals in gaseous BeCl2
atomic orbitals
hybrid orbitals
orbital box diagrams
VSEPR predicts a
linear shape
11-20
Figure 11.2
The sp hybrid orbitals in gaseous BeCl2 (continued)
orbital box diagrams with orbital contours
11-21
Figure 11.3
The sp2 hybrid orbitals in BF3
VSEPR predictsa trigonal planar
shape
11-22
Figure 11.4
The sp3 hybrid orbitals in CH4
VSEPR predicts a tetrahedral
shape
11-23
Figure 11.5
The sp3 hybrid orbitals in NH3
VSEPR predictsa trigonal
pyramidal shape
11-24
Figure 11.5
The sp3 hybrid orbitals in H2O
VSEPR predictsa bent (V) shape
11-25
Figure 11.6
The sp3d hybrid orbitals in PCl5
VSEPR predictsa trigonal bipyramidal
shape
11-26
Figure 11.7
The sp3d2 hybrid orbitals in SF6
VSEPR predicts anoctahedral shape
11-27
11-28
Figure 11.8
Conceptual steps from molecular formula to the hybrid orbitals used in bonding
molecular formula
Lewis structure
molecular shape
and e- group arrangement
hybrid orbitals
Figure 10.1
Step 1
Figure 10.12
Step 2 Step 3
Table 11.1
11-29
SAMPLE PROBLEM 11.1 Postulating Hybrid Orbitals in a Molecule
SOLUTION:
PROBLEM: Use partial orbital diagrams to describe how the mixing of atomic orbitals on the central atoms leads to hybrid orbitals in each of the following molecules.
PLAN: Use Lewis structures to establish the arrangement of groups and the shape of each molecule. Postulate the hybrid orbitals. Use partial orbital box diagrams to indicate the hybrid for the central atoms.
(a) methanol, CH3OH (b) sulfur tetrafluoride, SF4
(a) (a) CH3OH H
CH H
OH
The groups around C are arranged as a tetrahedron.
O has a tetrahedral arrangement
with two non-bonding e- pairs.
11-30
SFF
F
F
SAMPLE PROBLEM 11.1 (continued)
2p
2s
sp3 2p
2s
sp3
(b) SF4 has a seesaw shape with four bonding and one non-bonding e- pairs.
3p
3s
3d
S atomsp3d
3d
hybridized S atom
single C atom single O atom
hybridized O atomhybridized C atom
distortedtrigonal
bipyramidal
11-31
Figure 11.9
bonds in ethane, CH3-CH3
both carbons are sp3 hybridized s-sp3 overlaps to bonds
sp3-sp3 overlap to form a bond relatively even distribution of electron
density over all bonds
Covalent Bonds Between Carbon Atoms - Single Bonds
free rotation
~109.5o
11-32 Figure 11.10
and bonds in ethylene, C2H4
overlap in one position -
p overlap -
electron density
Covalent Bonds Between Carbon Atoms - Double Bonds
hindered rotation
~120o
11-33
Figure 11.11
and bonds in acetylene, C2H2
overlap in one position -
p overlap -
Covalent Bonds Between Carbon Atoms - Triple Bonds
hindered rotation
180o
11-34
Video: Hybridization
11-35
SAMPLE PROBLEM 11.2 Describing bonding in molecules with multiple bonds
SOLUTION:
PROBLEM: Describe the types of bonds and orbitals in acetone, (CH3)2CO.
PLAN: Use the Lewis structure to determine the arrangement of groups and the shape at each central atom. Postulate the hybrid orbitals, taking note of multiple bonds and their orbital overlaps.
H3C
C
CH3
O
sp3 hybridized
sp3 hybridized
CC
C
O
H
H
HHH
H
sp2 hybridized
bondsbond
CC
C
O
sp3
sp3
sp3
sp3
sp3
sp3
sp3
sp3
sp2 sp2
sp2
sp2
sp2sp2
H
HH
HH
H
11-36 Figure 11.12
Restricted rotation in -bonded molecules
cis trans
No spontaneous interconversion betweencis and trans forms (isomers) in solution at room temperature!
11-37
Limitations of VB Theory
Inadequately explains magnetic/spectral properties
Inadequately treats electron delocalization
VB theory assumes a localized bonding model
11-38
Molecular Orbital (MO) Theory
A delocalized bonding model
A quantum-mechanical treatment of moleculessimilar to that used for isolated atoms
Invokes the concept of molecular orbitals (MOs)(extension of atomic orbitals)
Exploits the wave-like properties of matter (electrons)
11-39
Central themes of molecular orbital (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).
11-40
Amplitudes of wave functions are added
Figure 11.13
An analogy between light waves and atomic wave functions
Amplitudes of wave functions are
subtracted
11-41
Figure 11.14
Contours and energies of the bonding and antibonding molecular orbitals in H2
11-42
Bonding MO: lower in energy than isolated atoms
Antibonding MO: higher in energy than isolated atoms
number of AOs combined = number of MOs produced
To form MOs, AOs must have similar energy and orientation
Sigma () and pi () bonds are denoted as before; a star (asterick)is used to denote antibonding MOs.
11-43
Figure 11.15
Molecular orbital diagram for the H2 molecule
MOs are filledin the same sequenceas for AOs
(aufbau and exclusion
principles, Hund’s rule)
11-44
The MO bond order
[1/2 (no. of e- in bonding MOs) - (no. of e- in antibonding MOs)]
higher bond order = stronger bond
Has predictive power!
11-45 Figure 11.16
MO diagrams for He2+ and He2
En
erg
y
MO of He+
*1s
1s
AO of He+
1s
MO of He2
AO of He
1s
AO of He
1s
*1s
1s
En
erg
y
He2+ bond order = 1/2 He2 bond order = 0
AO of He
1s
can exist! cannot exist!
11-46
SAMPLE PROBLEM 11.3 Predicting species stability using MO diagrams
SOLUTION:
PROBLEM: Use MO diagrams to predict whether H2+ and H2
- can exist. Determine their bond orders and electron configurations.
PLAN: Use H2 as a model and accommodate the number of electrons in bonding and antibonding orbitals. Calculate the bond order.
1s1s
AO of HAO of H
1s1s
MO of HMO of H22++
bond order = 1/2(1-0) = 1/2
HH22++ does exist! does exist!
MO of HMO of H22--
bond order = 1/2(2-1) = 1/2
H2- does exist!
1s1s 1s1s
AO of HAO of H AO of HAO of H--
configuration is (1s)2(
1s)1
AO of HAO of H+
configuration is (1s)1
11-47
*2s
2s
2s2s
1s
*1s
1s
1s
Figure 11.17
1s
*1s
1s
1s
2s 2s
*2s
2s
Li2 bond order = 1 Be2 bond order = 0
Bonding in s-block homonuclear
diatomic moleculesEn
erg
y
Li2Be2
11-48
Bonding and antibonding MOs for coreelectrons cancel = no net contribution to bonding
Only MO diagrams showing MOs created bycombining valence-electron AOs are important.
11-49
Figure 11.18
Contours and energies of and MOs through combinations of 2p atomic orbitals
end-to-endoverlap
side-to-sideoverlap
11-50
Relative energies
2p < 2p < *2p < *2p
More effective end-to-end interactionrelative to side-to-side in bonding MOs
11-51
Figure 11.19
Relative MO energy levels for Period 2 homonuclear diatomic molecules
MO energy levels for O2, F2 and Ne2
MO energy levels for B2, C2 and N2
without 2s-2p mixing
with 2s-2p mixing
11-52
Figure 11.20
MO occupancy and molecular properties for B2 through Ne2
11-53
Figure 11.21
The paramagnetic properties of O2
Explained byMO diagram
11-54
SAMPLE PROBLEM 11.4 Using MO theory to explain bond properties
SOLUTION:
PROBLEM: As the following data show, removing an electron from N2 forms
an ion with a weaker, longer bond than in the parent molecule, whereas the ion formed from O2 has a stronger, shorter bond.
PLAN: Find the number of valence electrons for each species, draw the MO diagrams, calculate bond orders, and compare the results.
Explain these facts with diagrams showing the sequence and Explain these facts with diagrams showing the sequence and occupancy of MOs.occupancy of MOs.
bond energy (kJ/mol)bond energy (kJ/mol)
bond length (pm)bond length (pm)
N2 N2+ O2 O2
+
945945
110110
498498841841 623623
112112121121112112
N2 has 10 valence electrons, so N2+ has 9.
O2 has 12 valence electrons, so O2+ has 11.
11-55
SAMPLE PROBLEM 11.4 (continued)
2s
2s
2p
2p
2p
2p
N2 N2+ O2 O2
+
bond orders
1/2(8-2) = 3 1/2(7-2) = 2.5 1/2(8-4) = 2 1/2(8-3) = 2.5
2s
2s
2p
2p
2p
2p
bonding e- lost
antibonding
e- lost
(weaker) (weaker)
11-56
En
erg
y
MO of HF
AO of H
1s
2px 2py
AO of F
2p
Figure 11.22
The MO diagram for HF
Heteronuclear DiatomicMolecules
lower in energythan 1s of H!
nonbonding MOs
11-57
In polar covalent compounds, bonding MOsare closer in energy to the AOs of the more
electronegative atom.
11-58
En
erg
yFigure 11.23
The MO diagram for NO
MO of NO
2s
AO of N
2p
*2s
2s
2sAO of O
2p
2p
2p
*2p
*2s
N O
0 0
N O
-1 +1
possible Lewis structures
bond order = 2.5
11-59
Figure 11.24
The lowest energy -bonding MOs in benzene and ozone
OO O
resonance hybrid