Transcript Theories of Chemical Bonding

Molecular Orbitals of Heteronuclear Diatomic Molecules
The MO for heteronuclear can be derived from combination of atomic
orbitals (AO) of participating atoms. Bear in mind, however, that the
energy levels for AO are different for heteronuclear molecules. The
variation of energy levels as a function of effective atomic number Zeff
and quantum number n is exemplified by the equation,
Zeff2
E = – RH ---------n2
From this, we place the energy levels of AOs according to their
relative heights. The AOs that are used to form MOs are represented
by some lines, and the energy levels are indicated accordingly.
The formation of MOs from AOs of simple molecules such as LiH and
LiF will be illustrated in the next few slides.
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The MOs for LiH
The 1s AO energy levels of H is lower
than the 2s AO of Li. Interactions of
these AOs lead to the formation of s
and s* bonds.
Please compare this energy level
diagram with that of H2
__
s*
_
_

s
AOLi
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AOH
2
MO for LiF
In the case of LiF, the energy of 2s AO of Li is higher the energy of 2p of F.
The MOs are derived
from the combination
of 2s AO of Li and 2p
AO of F.
A s and a s* are
derived from this
interaction, and the
two 2p orbitals are
not involved. These
are non-bonding
orbitals or pairs.
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MO for H2O
__ __
s*
_ _
  _ _
 
non-bonding sp3 hybrid AO
 
s
2H
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sp3 hybrid
 _ _
2p

2s
O
4
MO for Water, another approach
O–H = 97.3 pm
H-O-H = 105.6o
O charge=-0.782
2 H charge= 0.391
Dipole moment
= 2.15535 Debye
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MO Plots of O2
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MO
plots for
CO
C0 = 114.6 pm
C1 charge= –0 .014
O2 charge= 0.014
Dipole moment
= 0.13196 Debye
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