Download 1 Chem 1: Chapter # 11: Theories of Covalent Bonding: VALENCE

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Chem 1: Chapter # 11: Theories of Covalent Bonding:
VALENCE BOND THEORY:
DEVELOPED BY LINUS PAULING, who received the Nobel Prize in 1954 for his work
A view of chemical bonding in which bonds arise from the overlap of atomic orbitals on two
atoms to give a bonding orbital of electrons localized between the bonded atoms
RULE: Realize that Valence Bond Theory and all the others don't explain everything
VALENCE BOND THEORY:
The Central Themes of VB Theory
Basic Principle:
A covalent bond forms when the orbitals of two atoms overlap and the overlap region, which is
between the nuclei, is occupied by a pair of electrons.
(The two wave functions are in phase so the amplitude increases
between the nuclei.)
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.
There is a hybridization of atomic orbitals to form molecular
orbitals.
VALENCE BOND THEORY:
Ha to Hb: 1sa to 1sb
overlap radius = 74 pm
As overlap increases, strength of bond increases - both electrons are mutually attracted to both
atomic nuclei.
At optimum distance between nuclei with maximum overlap, a sigma σ bond (strong primary
bond) forms. Max electron density is along the axis of the bond
Ha to Fb: 1sa to 2pb
direct overlap or σ bond
F to F: the picture looks like a 2p orbital on one F is overlapping with a 2p orbital on the other F
atom, but actually each F is sp3 hybridized & electrons are localized between two atomic nuclei
We cannot use this direct overlap picture for CH4’s bonding. The 2s and the three 2p orbitals on
each C do not fit into the CH4 molecule's 109o bond angles, since the 2p orbitals are at 90° to
each other
Valence Bond Theory states that HYBRID orbitals of the outermost orbitals on an atom are
formed from the atoms’ atomic orbitals
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.
You write the types of hybrid orbitals here:
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Figure 11.2The sp hybrid orbitals in gaseous BeCl2(continued).
Figure 11.3:The sp2 hybrid orbitals in BF3.
Figure 11.4: The sp3 hybrid orbitals in CH4.Figure: 11.5: The sp3 hybrid orbitals in NH3. The sp3
hybrid orbitals in H2O.
Expanded Valence Shells have hybrid orbitals using s, p & d atomic orbitals. Example: PCl5 P:
[Ne]3s23p3
dsp3 hybridization results in 5 σbonds and trigonal bipyramidal geometry
(You can write these as dsp3 or sp3d)
Figure 11.6: The sp3d hybrid orbitals in PCl5.Figure 11.7: The sp3d2 hybrid orbitals in SF6.
Draw the figure 11.8 map here: The conceptual steps from molecular formula to the hybrid
orbitals used in bonding.
Problem 11.1: Use partial orbital diagrams to describe mixing of the atomic orbitals of the
central atom leads to hybrid orbitals in each of the following:
(a) Methanol, CH3OH
(b) Sulfur tetrafluoride, SF4
There can be more than one central atom, and each has its own hybridization and geometry
C2H6 and H2O2 and CH3COOH
C2H6: both C's are sp3 hybridized and can rotate around axis of bond.
H2O2: both O's are sp3 , etc.
VALENCE BOND THEORY: Multiple Bonds
H2CO: the Lewis structures shows a double bond between C and O, but we know it does not
have twice the bond dissociation energy of a single C-O bond
Pauling proposed that there was only one sigma bond between any two atoms, and the other
multiples were weaker pi bonds
If there are only 3 σ bonds around this carbon, it can't be sp3 hybridized - instead we have sp2
hybrid orbitals
sp2 hybridization results in only 3 σbonds, and trig planar geometry, with 120° angles
π bond is a sideways or parallel overlap of the p atomic orbitals rather than the direct overlap of
σ bonds
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Figure: 11.9: The s bonds in ethane(C2H6).Figure 11.10 The s and p bonds in ethylene (C2H4).
Problem 11.2 : Describe the types of bonds and orbitals in acetone, (CH3)2CO.
Look at acetylene: its geometry is linear. C is forming a triple bond to another C and a single
bond to H, so that's only two σbonds
Therefore sp hybridization results in only 2 σ bonds, and linear geometry
There are 2 π bonds from the parallel overlap of the 2p orbitals remaining on both C's
Figure 11.13 from 4th ed.
VALENCE BOND THEORY: RESONANCE
Resonance Structures and π Bonding:
π resonance structures involve an electron pair used alternately as a π bond or a LP
Ozone: O3 O==O--O or O--O==O
All are sp2, trig planar, each has 3 sp2 orbitals and a p orbital remaining.
VALENCE BOND THEORY
Benzene: C6H6 has carbons with sp2 hybrids and 120o angles, each C has 2 σ bonds to other C's,
1 σ bond to H, and 1 π bond electron available
-s
SUMMARY: draw the Lewis structure; determine arrangement of electron pairs using VSEPR,
specify the hybrid orbitals to accommodate the e- pairs
MOLECULAR ORBITAL THEORY:
- explains why H2 forms easily and He2 does not
- is an alternate way of viewing e- orbitals in molecules where pure s and pure p orbitals combine
to produce orbitals that are delocalized over the molecule
- they can have different energies and are assigned electrons just like we do in an atom - Pauli
exclusion principle and Hund's rule included
Pauling's Valence Bond Theory does not explain everything
MO Theory doesn't either, but it does correctly predict the electronic structure of certain
molecules that do not follow Lewis's approach, including the paramagnetism of certain
molecules, like O2
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.
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Atomic wave functions are summed to obtain molecular wave functions. The number of
molecular orbitals produced is always = # of atomic orbitals brought by the combining atoms
(only orbitals on different atoms are combined).
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).
The electrons of the molecule are placed in bonding or antibonding orbitals of successively
higher energy (just like Hund's rule).
Atomic orbitals combine most effectively with orbitals of the same type and similar energy (s
w/s, n=2 w/ n=2)
Figure: 11.13 An analogy between light waves and atomic wave functions.
Figure 11.14: Contours and energies of the bonding and antibonding molecular orbitals (MOs) in
H2.
MOLECULAR ORBITAL THEORY
BOND ORDER: the number of bonding e- pairs shared by 2 atoms in a molecule
Fractional bond orders are possible in MO Theory!
Silberberg method:
B.O. = ½(# of e- in bonding orbitals - # of e- in antibonding orbitals)
Figure 11.15: The MO diagram for H2.
Figure: 11.16 MO diagram for He2+ and He2.
Problem 11.3: Use MO diagrams to predict whether H2+ and H2- exist. Determine their bond
orders and electron configurations.
Figure 11.17: Bonding in s-block homonuclear diatomic molecules.
Figure 11.20: MO occupancy and molecular properties for B2 through Ne2
Figure 11.21 The paramagnetic properties of O2
MO Theory Practice
1. Draw the bonding and antibonding molecular orbitals for H2.
2. Do Valence Bond Theory (hybridization) and MO Theory for both O2 and O22-. Which
theory works better to explain the molecule and ion?
3. For N2, N2+ and N2- compare
a. Magnetic character
b. Net number of π bonds
c. Bond Order
d. Bond length
e. Bond strength