Chemistry Reference
In-Depth Information
the shapes of the s, p, and d orbitals as shown in
Figure 2.1
and the same symmetry arrangements as in CFT but with
the additional use of the molecular orbital theory (MOT) of chemical bonding. MOT combines the approximate
energies andwave functions of all of the component atomic orbitals to obtain the best approximations for the energies
and wave functions of the molecule. In other words, it makes use of covalency in the metal
ligand interactions.
During the formation of a molecule, the atomic orbitals of the individual nuclei interact. Such interactions may
be constructive or destructive depending on whether their wave functions add or subtract in the region of overlap.
Which orbitals can overlap effectively is dictated by symmetry considerations and only orbitals of matching
symmetry may interact. A constructive interaction will result in the formation of two types of bonding molecular
orbital: the
e
interactions will give rise to antibonding orbitals called
s
and the
p
s
* and
p
* with an associated decrease in electron density.
The bonds associated with
s
and
p
orbitals are called respectively
s
and
p
bonds. In simplistic terms direct,
'head-on' overlap of two suitably orientated orbitals result in a
-bond with uniform distribution of charge density
around the axis of the bond whereas 'side-ways' overlap will give rise to a
s
-bond with distribution of the charge
density above and below a plane crossing the axis of the bond. The electrons involved in the latter type of bonding
are spread out over a greater volume than those involved in the former type. A
p
p
-bond will hence be more readily
polarised than a
-bond and such bonds are said to be delocalised as sideways overlap occurs between all orbitals
in the vertical plane and all those in the horizontal plane. This is the case of alkynes and nitriles, both possessing
two sets of
s
-bonds perpendicular to each other. Delocalisation gives additional stability to a molecule as the
increase in the volume of the space occupied by the electrons involved lowers the potential energy of the system.
The bonds involved in coordination complexes can then be described as
p
s
-bonds (any lone pair donation from
p
p
a ligand to the metal) and
orbitals
of the ligand or from the p orbitals of the ligand to the metal d orbitals). In the octahedral environment of a central
metal atom with six surrounding ligands, the s, p
x
, p
y
, p
z
, d
z
2
, and d
x
2
-bonds (any donation of electron density from filled metal orbitals to vacant
2
valence shell orbitals of the central metal
atom have lobes lying along the metal
e
ligand bond directions and hence are suitable for
s
bonding. The
orientation of the d
xy
, d
xz
, and d
yz
makes such orbitals appropriate only for
y
p
-bonding. It is assumed that each
ligand system to give a bonding (maximum positive overlap) and an antibonding (maximum negative overlap)
molecular orbital. The simplified MO diagram
10
for the formation of a sigma-bonded octahedral ML
6
complex is
shown in
Figure 2.8
.
If a molecular orbital is closer in energy to one of the atomic orbitals used to construct it than to the other one, it
shall have more the character of the first one than the second one. Hence, the electrons occupying the six bonding
s
ligand possesses one
s
molecular orbitals will be largely “ligand” electrons with some metal ion character. Electrons occupying the
antibonding orbitals will be mainly “metal” electrons. During the complex formation the metal d electrons will go
either only to t
2g
or to both t
2g
and e
g
. In the absence of any
p
bonding any electrons in the t
2g
(which could
-bonding) will be purely metal electrons and the level is essentially nonbonding, whereas the e
g
levels participate in
contribute to
p
-bonding with the ligand. In other words, the central portion of the diagram closely
resembles the t
2g
and e
g
orbitals derived from crystal field theory (
Figure 2.7
)
, with one difference: the e
g
orbital is
now e
g
.
In terms of crystal field theory, the larger gap between t
2g
and e
g
energy levels in strong field ligands is
essentially a consequence of the raising of the e
g
energy levels by electrostatic interactions between the ligand and
the d electrons of the metal. However, the molecular orbital model shows how the difference in energy
s
D
could
p
bonding
8. Pronounced as sigma and pi from the Greek letters.
9. If the ligands possess also
p
-orbitals these have to be taken into account.
10. a
1g
, e
g
, t
1u
encountered in the diagram are symmetry symbols for the associated orbitals: a
1g
represents a single orbital which has the full
symmetry of the molecular system; e
g
and t
1u
represent a set of two and three orbitals, respectively, which are equivalent within the individual
set apart from their orientation in space. Subscripts g and u indicate whether the orbital is centrosymmetric or anticentrosymmetric.
