Inorganic Pharmaceutical Chemistry
Lecture No. 6
Date : 22/11 /2012
Dr. Mohammed Hamed
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Hybrid Orbitals
Hybridization was introduced to explain molecular structure when the
valence bond theory failed to correctly predict them. It is experimentally
observed that bond angles in organic compounds are close to 109o, 120o,
or 180o. According to Valence Shell Electron Pair Repulsion (VSEPR)
theory, electron pairs repel each other and the bonds and lone pairs
around a central atom are generally separated by the largest possible
angles.
Based on the valence bond theory, carbon would only be able to form two
covalent bonds, making CH2. However, as you will find out, we know that this
is not true and that in reality, it makes CH4. The hybridization of orbitals is also
greatly favored because hybridized orbitals are lower in energy compared to
their separated, unhybridized counterparts. This results in more stable
compounds when hybridization occurs. Also, major parts of the hybridized
orbitals, or the frontal lobes, overlap better than the lobes of unhybridized
orbitals. This leads to better bonding.
Carbon is a perfect example showing the need for hybrid orbitals. As you know,
Carbon's ground state configuration is:
6
C 1S2 2S2 2P2
According to the valence bond theory, carbon should form two covalent bonds,
resulting in a CH2. However, tests show that CH2 is highly reactive and cannot
exist outside of a reaction. Therefore, this does not explain how CH4 can exist.
However, you can excite a 2s electron and bump it into one of the 2p orbitals.
This would give you the following configuration:
While this would allow us to have four covalent bonds, resulting in CH4, it also
implies that the C-H covalent bonds would have different energies due to the
different levels of orbital overlap. However, with testing, it has been proven that
in CH4, any hydrogen can be removed with the same amount of energy. This
means that every C-H covalent bond should have equal energies. Once again,
this means that the valence bond theory fails to explain the existence of CH 4.
The only way it can be explained is if when we had the exited state above, the 2s
and the 3 2p orbitals fused together to make four, equal energy sp3 hybrid
orbitals. That would give us the following configuration:
Z
Z
Y
2S
X Y
2PX X Y
Z
2Py
Z
X
C
Y
2PZ
X
This explains how a carbon can have four equal energy bonds. The next section
will explain the various types of hybridization and how each type helps explain
the structure of certain molecules.
TYPES OF HYBRIDIZATION
During hybridization, the atomic orbitals with different characteristics are mixed
with each other. Hence there is no meaning of hybridization between same type
of orbitals i.e., mixing of two 's' orbitals or two 'p' orbitals is not called
hybridization.
However orbital of 's' type can mix with the orbitals of 'p' type or of 'd' type.
Based on the type and number of orbitals, the hybridization can be subdivided
into following types.
sp HYBRIDIZATION
* Intermixing of one 's' and one 'p' orbitals of almost equal energy to give two
identical and degenerate hybrid orbitals is called 'sp' hybridization.
* These sp-hybrid orbitals are arranged linearly at by making 180o of angle.
* They possess 50% 's' and 50% 'p' character.
sp2 HYBRIDIZATION
* Intermixing of one 's' and two 'p' orbitals of almost equal energy to give three
identical and degenerate hybrid orbitals is known as sp2 hybridization.
* The three sp2 hybrid orbitals are oriented in trigonal planar symmetry at angles
of 120o to each other.
* The sp2 hybrid orbitals have 33.3% 's' character and 66.6% 'p' character.
sp3 HYBRIDIZATION
* In sp3 hybridization, one 's' and three 'p' orbitals of almost equal energy
intermix to give four identical and degenerate hybrid orbitals.
* These four sp3 hybrid orbitals are oriented in tetrahedral symmetry with
109o28' angle with each other.
* The sp3 hybrid orbitals have 25% ‘s’ character and 75% 'p' character.
sp3d HYBRIDIZATION
* In sp3d hybridization, one 's', three 'p' and one 'd' orbitals of almost equal
energy intermix to give five identical and degenerate hybrid orbitals, which are
arranged in trigonal bipyramidal symmetry.
Among them, three are arranged in trigonal plane and the remaining two orbitals
are present above and below the trigonal plane at right angles.

The sp3d hybrid orbitals have 20% 's', 60% 'p' and 20% 'd' characters.
sp3d2 HYBRIDIZATION
* Intermixing of one 's', three 'p' and two 'd' orbitals of almost same energy by
giving six identical and degenerate hybrid orbitals is called sp3d2 hybridization.
* These six sp3d2 orbitals are arranged in octahedral symmetry by making 90 o
angles to each other. This arrangement can be visualized as four orbitals
arranged in a square plane and the remaining two are oriented above and below
this plane perpendicularly.
sp3d3 HYBRIDIZATION
* In sp3d3 hybridization, one 's', three 'p' and three 'd' orbitals of almost same
energy intermix to give seven sp3d3 hybrid orbitals, which are oriented in
pentagonal bipyramidal symmetry.
* Five among the sp3d3 orbitals are arranged in a pentagonal plane by making
72o of angles. The remaining are arranged perpendicularly above and below this
pentagonal plane.
In 1927, valence bond theory was formulated and it argues that a chemical bond
forms when two valence electrons, in their respective atomic orbitals, work or
function to hold two nuclei together, by virtue of effects of lowering system
energies. Pauling presented six rules for the shared electron bond,
1. The electron-pair bond forms through the interaction of an unpaired
electron on each of two atoms.
2. The spins of the electrons have to be opposed.
3. Once paired, the two electrons cannot take part in additional bonds.
4. The electron-exchange terms for the bond involves only one wave
function from each atom.
5. The available electrons in the lowest energy level form the strongest
bonds.
6. Of two orbitals in an atom, the one that can overlap the most with an
orbital from another atom will form the strongest bond, and this bond will
tend to lie in the direction of the concentrated orbital.
THE SHAPES OF COMPLEX METAL IONS
To describes the shapes of some complex metal ions. It goes on to look at some
simple examples of stereoisomerism (geometric and optical) in complex ions.
These shapes are for complex ions formed using monodentate ligands - ligands
which only form one bond to the central metal ion.
6-coordinated complex ions
These are complex ions in which the central metal ion is forming six bonds. In
the simple cases we are talking about, that means that it will be attached to six
ligands. These ions have an octahedral shape. Four of the ligands are in one
plane, with the fifth one above the plane, and the sixth one below the plane. The
diagram shows four fairly random examples of octahedral ions.
4-coordinated complex ions
These are far less common, and they can take up one of two different shapes.
Tetrahedral ions
There are two very similar ions which crop up commonly at this level: [CuCl 4]2and [CoCl4]2-. The copper(II) and cobalt(II) ions have four chloride ions bonded
to them rather than six, because the chloride ions are too big to fit any more
around the central metal ion
A square planar complex
Occasionally a 4-co-ordinated complex turns out to be square planar. There's no
easy way of predicting that this is going to happen. The only one you might
possibly come across at this level is cisplatin which is used as an anti-cancer
drug.
Cisplatin is a neutral complex, Pt(NH3)2Cl2. It is neutral because the 2+ charge
of the original platinum(II) ion is exactly cancelled by the two negative charges
supplied by the chloride ions
The platinum, the two chlorines, and the two nitrogens are all in the same plane.
We will have more to say about cisplatin immediately below
Stereoisomerism in complex ions
Some complex ions can show either optical or geometric isomerism.
Geometric isomerism
This occurs in planar complexes like the Pt(NH3)2Cl2 we've just looked at. There
are two completely different ways in which the ammonias and chloride ions
could arrange themselves around the central platinum ion:
The two structures drawn are isomers because there is no way that you can just
twist one to turn it into the other. The complexes are both locked into their
current forms
The terms cis and trans are used in the same way as they are in organic
chemistry. Trans implies "opposite" - notice that the ammonias are arranged
opposite each other in that version, and so are the chlorines. Cis implies "on the
same side" - in this instance, that just means that the ammonias and the chlorines
are next door to each other.
Optical isomerism
You recognise optical isomers because they have no plane of symmetry. In the
organic case, it is fairly easy to recognise the possibiliy of this by looking for a
carbon atom with four different things attached to it. It isn't qute so easy with
the complex ions - either to draw or to visualise!
The examples you are most likely to need occur in octahedral complexes which
contain bidentate ligands - ions like [Ni(NH2CH2CH2NH2)3]2+ or [Cr(C2O4)3]3-.
The diagram below shows a simplified view of one of these ions. Essentially,
they all have the same shape - all that differs is the nature of the "headphones". I
have deliberately left the charges off the ion, because obviously they will vary
from case to case. The shape shown applies to any ion of this kind.
If your visual imagination will cope, you may be able to see that this ion has no
plane of symmetry. If you find this difficult to visualise, the only solution is to
make the ion out of a lump of plasticene (or a bit of clay or dough) and three bits
of cardboard cut to shape
A substance with no plane of symmetry is going to have optical isomers - one of
which is the mirror image of the other. One of the isomers will rotate the plane
of polarisation of plane polarised light clockwise; the other rotates it anticlockwise
Coordination complexes in biochemistry
Approximately one-third of the chemical elements are present in living
organisms. Many of these are metallic ions whose function within the cell
depends on the formation of d-orbital coordination complexes with small
molecules such as porphyrins (see below). These complexes are themselves
bound within proteins (metalloproteins) which provide a local environment that
is essential for their function, which is either to transport or store diatomic
molecule (oxygen or nitric oxide), to transfer electrons in oxidation-reduction
processes, or to catalyze a chemical reaction. The most common of these utilize
complexes of Fe and Mg, but other micronutrient metals including Cu, Mn, Mo,
Ni, Se, and Zn are also important.
Geometry and magnetic nature of some complexes
Type
No. of Geome
Oxidation
Magnetic
of
unpaired try
state of
nature
hybrid
electrons shape
metal
(7)
ization
(6)
(5)
(3)
(4)
Paramag
netic
Paramag
netic
Diamagn
etic
Paramag
netic
Diamagn
etic
Ni2+(d8)
+2
[NiCl4]2–
+2
[Ni(CN)4]2+
0
Ni
0
Ni(CO)4
+2
[Ni(NH3)6]2
+2
Mn2+(d5)
1
Octah d2sp3(I
edral nner)
+2
[Mn(CN)6]4–
5
Tetrah
sp3
edral
+2
[MnCl4]2
+2
Cu2+(d9)
+2
[CuCl4]2–
+2
[Cu(NH3)4]2
2
0
Tetrah
sp3
edral
Squar
e
dsp2
planar
2
Tetrah
sp3
edral
Paramag
netic
2
Octah sp3d2(out
edral er)
Paramag
netic
5
Paramag
netic
Paramag
netic
Paramag
netic
Paramag
netic
Atom/ion/
complex
(1)
+2
2
0
Paramag
netic
Configuration
(2)
1
1
1
Tetrah
sp3
edral
Squar dsp2
e
planar
+
+
Paramagn
etic
3
Paramagn
etic
3
Paramagn
etic
3
Paramagn
etic
4
Paramagn
etic
4
Diamagnet
ic
0
Paramagn
etic
3
Paramagn
etic
Paramagn
etic
Diamagnet
ic
Paramagn
etic
Paramagn
etic
Paramagn
etic
3
Cr3+(d3)
Octahe d2sp3 (I
dral nner)
+3
[Cr(NH3)6]3+
Octahe sp3d2 (
dral Outer)
+3
[Cr(H2O)6]3+
+3
CO3+(d6)
Octahe sp3d2 (
dral Outer)
+3
[CoF6]3–
Octahe d2sp3 (I
dral nner)
+3
[Co(NH3)6]3+
+2
CO2+(d7)
+2
[Co(H2O)6]2+
+2
Fe2+(d6)
Octahe sp3d2 (
dral Outer)
4
0
Octahe d2sp3 (I
dral nner)
+2
[Fe(CN)6]4–
4
Octahe sp3d2 (
dral Outer)
+2
[Fe(H2O)6]2+
4
5
Paramagn
etic
1
Paramagn
etic
4
Diamagnet
ic
+3
o
Octahe sp3d2 (
dral Outer)
Octahe
dral
d2sp3 (I
nner)
Trigon
al
dsp3
bipyra (Inner)
midal
+2
Same
[Fe(NH3)6]2+
+3
Fe3+(d5)
+3
[Fe(CN)63–
o
Fe
o
Fe(CO)5