Inorganic Chemistry : STEROCHEMISTRY AND BONDING IN MAIN GROUP COMPOUNDS Content of chapter: 01. Hybridization 02. VSEPR 03. Walse Diagram 04. Bent Rule 05. Dπ - Pπ Bonds

1. PAPER CHNN-401
INORGANIC CHEMISTRY
UNIT 1
STEROCHEMISTRY AND BONDING
IN MAIN GROUP COMPOUNDS
Dr. Gaurang Rami
M.Sc. / Ph.D. / CSIR-NET / GATE / GSET
Assistant Professor
Shri Sarvajanik Science College (PG), Mehsana

2. 1.Hybridization
2.VSEPR
3.Walse Diagram
4.Bent Rule
5.Dπ - Pπ Bonds

3. Definition: The phenomenon of mixing up of atomic orbitals of similar
energies and formation of equivalent number of entirely new orbitals of
identical shape and energy is known as "hybridization" and the new
orbitals so formed is called as "hybrid orbitals".
Characteristic Features of Hybridization
1. The number of hybrid orbitals is the same as the number of atomic orbitals
intermixed.
2. Hybrid orbitals are equivalent in shape and have the same energies.
3. Hybrid orbitals form stronger bonds than their pure atomic counterparts.
4. Hybrid orbitals are directed in specific directions and therefore control the
geometry of the molecule.
5. Only valence shell orbitals take part in hybridization.
6. There should be only a small energy difference in the orbitals undergoing
hybridization.
7. Electron promotion is not a necessary condition to be followed.
8. In addition to the half-filled orbitals, fully filled and empty orbitals can
also undergo hybridization.

4. TYPE OF HYBRIDIZATION
Depending on the types of orbitals included in mixing, hybridisation could be
characterised as sp³, sp², sp, sp³d, sp³d², or sp³d³.
01. sp Hybridisation
Sp hybridisation is a type of hybridisation in which only the 1s and 1p orbitals of the
same element are involved. The production of the acetylene molecule is an instance.

5. 02. sp² Hybridisation
Sp² hybridisation combines and recasts nearly identical 1s and 2p orbitals of the same atom to
generate three new sp² hybrid orbitals with equal energies, maximal symmetry, and determinate
orientation in space. Consider the synthesis of an ethylene molecule.

6. 03. sp³ Hybridisation
sp³ hybridisation is a type of hybridisation in which 1s and 3p orbitals of the same element are
mixed and recast to generate a new hybrid orbital with the same energy, symmetry, or fixed
orientation in space. Types of hybridisation examples include the development of a methane
molecule.

7. 04. sp³d Hybridisation
The hybridisation is termed sp³d whenever 1s and 3p orbitals of the same element mix and recast
to generate hybrid orbitals with the same energies or equal orientation in space. PCl₅ is an
example of a molecule. The central phosphorus atom in PCl₅ experiences sp³d hybridisation,
resulting in the formation of five sp³d hybrid orbitals.

8. 05. sp³d² Hybridisation
sp³d² hybridisation occurs when 1s, 3p, and 2d-orbitals of the same elements combine and recast
to generate hybrid orbitals with the same energy and equal orientation in space. The formation of
SF₆ demonstrates this hybridisation.

10. 2. VSEPR
The VSEPR theory is used to predict the shape of the molecules from the
electron pairs that surround the central atoms of the molecule.
The theory was first presented by Sidgwick and Powell in 1940. The two
primary founders of the VSEPR theory are Ronald Nyholm and Ronald
Gillespie. This theory is also known as the Gillespie-Nyholm theory to
honour these chemists.
The VSEPR theory is based on the assumption that the molecule will take
shape such that electronic repulsion in the valence shell of that atom is
minimized.
The Valence Shell Electron Pair Repulsion Theory, abbreviated as VSEPR
theory, is based on the premise that there is a repulsion between the pairs
of valence electrons in all atoms, and the atoms will always tend to arrange
themselves in a manner in which this electron pair repulsion is minimalised.
This arrangement of the atom determines the geometry of the resulting
molecule.

11. Postulates of VSEPR Theory
The postulates of the VSEPR theory are listed below.
1. In polyatomic molecules (i.e., molecules made up of three or more atoms),
one of the constituent atoms is identified as the central atom to which all
other atoms belonging to the molecule are linked.
2. The total number of valence shell electron pairs decides the shape of the
molecule.
3. The electron pairs have a tendency to orient themselves in a way that
minimizes the electron-electron repulsion between them and maximizes
the distance between them.
4. The valence shell can be thought of as a sphere wherein the electron pairs
are localised on the surface in such a way that the distance between them is
maximised.
5. Should the central atom of the molecule be surrounded by bond pairs
of electrons, then the asymmetrically shaped molecule can be
expected.

12. 6. Should the central atom be surrounded by both lone pairs and bond pairs
of electrons, the molecule would tend to have a distorted shape.
7. The VSEPR theory can be applied to each resonance structure of a
molecule.
8. The strength of the repulsion is strongest in two lone pairs and weakest in
two bond pairs.
9. If electron pairs around the central atom are closer to each other, they will
repel each other. This results in an increase in the energy of the molecules.
10. If the electron pairs lie far from each other, the repulsions between them
will be less, and eventually, the energy of the molecule will be low.
11. The strength of the repulsion between a lone pair and a bond pair of
electrons lies in between the repulsion between two lone pairs and between
two bond pairs. The order of repulsion between electron pairs is as
follows:
Lone Pair- lone pair > Lone Pair- bond- pair > Bond Pair- bond pair

13. Following are the general rules for determining VSEPR geometries of a
given molecule.
1. Draw the Lewis structure of the molecule for which geometry has to
be evaluated.
2. Place the lone pairs (lp) and bonding pairs (bp) of electrons around
the central atom.
3. Electron–electron repulsions decrease in the sequence: lone pair–lone
pair (lp-lp) > lone pair–bonding pair (lp-bp) > bonding pair–bonding
pair (bp-bp).
4. If the central atom is connected to a surrounding atom by multiple
bond, then the electron–electron repulsions decrease in the order:
triple bond–single bond > double bond–single bond > single bond–
single bond.
5. Possible common shapes for a molecule or ion, AB n comprising of
central atom A surrounded by n atoms of B, are presented in below
Table.

14. The different geometries that molecules can assume in accordance with the
VSEPR theory can be seen in the illustration provided below.

15. Limitations of VSEPR Theory
Some significant limitations of the VSEPR theory include
1.Structures of IF7 and TeF7
- are isoelectronic with seven bonding pairs of
electrons around the respective central atoms. Hence, VSEPR predicts
pentagonal bipyramidal geometry for both of these structures.
VSEPR theory does not predict different bond distances for axial and
equatorial positions in pentagonal bipyramidal geometry.
Physical characterization establishes that the axial bonds of these structures
are slightly shorter than the respective equatorial bonds.
Additionally the geometry of TeF7
- in their crystallographically characterized
salts seems much distorted from the predicted pentagonal bipyramidal
geometry as the equatorial fluorine atoms are not coplanar.

16. 2. VSEPR theory successfully predicts the geometry of simple p-block
molecules but it is not appropriate to predict structures of the d-block
derivatives.
3. VSEPR theory does not take inert pair effect into account. Hence, it
does not explain structures of molecules derived from heavy elements of
periodic table. Crystallographic analyses has revealed that the species
such as [SeCl6]2-, [TeCl6]2- and [BrF6]- possess regular octahedral
geometry which is not justifiable using VSEPR theory.

17. 03 Walsh diagrams
Walsh diagrams were first introduced by Prof. A.D. Walsh to rationalize the
shapes adopted by polyatomic molecules in the ground state as well as in
excited states. W.D. are graphical representation depicting calculated orbital
binding energies of a molecule plotted against bond angles.
These diagrams predict geometries of small molecules and explain why a given
molecule is more stable in certain geometry than the other.
Walsh's rule states that a molecule will assume a structural geometry which
provides most stability to its highest occupied molecular orbital (HOMO).
W.D. is used to explain the regularity in the structures of related molecules
having identical number of valence electrons. For example, water (H2O) and
sulphur dioxide (H2S) exhibit similar structures. Walsh diagram also explains
the change in geometry of molecules with the change in their number of
electrons or their spin state.

18. Walsh diagrams often called angular coordinate diagrams or
correlation diagrams are representations of calculated orbital
energies of a molecule versus a distortion coordinate, used for making
quick predictions about the geometries of small molecules.
Walsh diagram for tri-atomic molecules
Simplified Walsh diagram for a triatomic molecule is depicted in
below figure which is energy versus bond angle plot.
It should be noted that the depicted energy levels are qualitative and
for actual system should be calculated by a suitable simulation.
MO levels drawn on the left are for the bent configuration with bond
angle of 90° whereas those on right are for the linear configuration
with bond angle 180°.

20. Walsh diagram for penta-atomic molecules:
For penta-atomic molecule, an imaginary model AH4 can be considered which can
assume a tetrahedral or square planar geometry. Methane (CH4) and sulfur
tetrafluoride (SF4) are the real examples for such categories.
For instance, for the formation of methane molecule one 2s and three 2p orbitals of
carbon and four 1s orbitals, one from each hydrogen atoms get involve in bonding.
Above mentioned eight orbitals result in four bonding and four antibonding orbitals
out of which bonding orbitals occupy all eight electrons involved in formation of
molecule.
In case of tetrahedral geometry significant overlap between the orbitals of carbon
and hydrogen is possible which reduces the energy of bonding orbitals whereas in
the square planar configuration the extent of overlap is very low which in turn
results in orbitals of considerably high energy as depicted in the given Figure
Hence, the CH4 molecule prefers the tetrahedral geometry rather than the square
planar or the intermediate distorted geometries.

22. 04. Bent's Rule
Bent's Rule describes and explains the relationship between the
orbital hybridizations of central atoms in molecules and the electro-
negativities of substituents.
The rule was stated by Henry Bent as follows:
"Atomic s character concentrates in orbitals directed toward electro-
positive substituents".
For example - molecular geometry can be explained and predicted by
changing the substituent group. In the molecule Me2XCl2, (X=main
group elements C, Si, Ge, Sn, Pb), the bond angle Cl-X-Cl is smaller
than C-X-C bond angle. With the highly electronegative halogen
substituent, Cl, more p character is concentrated on the central atom in
X-Cl than X-C bonds.
We know that the s- orbital is closer to the nucleus and so as the s-
character in a hybrid orbital increase, its electronegativity increases
too.

23. Bent's rule relates the orbital hybridization of central atom in a molecule
with the electro negativities of substituents and was originally stated by
Henry Bent as "Atomic s character concentrates in orbitals directed toward
electropositive substituents". Learners, you must recall that the ‘s’ orbital has
lower energy than the ‘p’ orbital. Moreover, a hybrid orbital having higher
‘s’ character has lower energy and ‘s’ orbital like shape.
On the other hand higher ‘p’ character results in higher energy and ‘p’
orbital like shape of the hybrid orbital. As the ‘s’ orbitals are closer to the
atomic nuclei, it stabilizes the lone pair of electron than the ‘p’ orbital does.
Hence, it is logical to say that more stable molecular model would allow the
orbitals rich in ‘p’ character for bonding purpose and those rich in ‘s’
character for accommodating the lone pairs of electrons.

24. It can also be stated that, electron density available on ‘s’ rich orbital is
closer to the nuclei hence, less available for bonding. Therefore,
electronegative substituent would withdraw electron density from ‘p’ rich
orbitals than the ‘s’ like orbitals.
We can illustrate this using structure and bonding in PCl3F2. Formation
of this molecule involves mixing of one s, three p and one d orbitals to
form five sp3d orbitals. The five hybrid orbitals are not identical but can
be grouped in set of two (axial) and three (equatorial) orbitals (Figure).
The equatorial set is resulted by s, px and py orbitals whereas axial set is
resulted by pz and d orbitals.
Experimental evidences indicate that the fluorine and chlorine atoms in
PCl3F2 molecule are situated at axial and equitorial positions respectively,
which is in agreement with Bent’s rule.

26. 05. dπ-pπ bond
Formation of inorganic molecules is different than that of organic molecules
in many aspects, one of which is the occurrence of dπ-pπ bonds.
Generally π bonds, as in case of organic molecules, form by lateral
overlapping of p orbitals present on two atoms such as carbon, nitrogen or
oxygen. Bonding interactions between two d orbitals resulting in δ bonds in
inorganic molecules are also prevalent.
However, bonding interactions in inorganic molecules can also make use
of suitably available d and p orbitals at once.
When a bond forms by lateral overlapping of p (or p*) and d orbitals
present on two different atoms, it is called dπ-pπ bond. Such bonds are
frequently observed in metal complexes such as carbonyls and nitrosyls.

27. A simple example is sulfur trioxide (SO3). Main group compounds such as
phosphine oxides and disiloxane can also feature the dπ-pπ bonds.
Presence of dπ-pπ bonding interactions usually result in shortening of bond
length and planar configuration of involved atoms.
However, observing such molecular features should not always be
attributed to dπ- pπ bonds as several other factors may also be playing role.
Hence, a careful evaluation of electronic and orbital symmetries must be
made.
Phosphine Oxide (R3P=O) offers an example of dπ-pπ bond in molecules
comprising of nonmetallic elements.
In this case all the p orbitals present on phosphorous are utilized in
hybridization and hence not available for lateral overlapping.

28. Empty d orbital on phosphorous accepts electron density from filled p
orbital available with oxygen atom.
This dπ-pπ bond causes tighter binding of both involved atoms which is
reflected in short bond distance (150 pm) and stability (bond energy 544
kJ/mol) of the bond.
Bonding in metal carbonyls is classical example of dπ-pπ bond, where empty
π* orbital present of carbonyl ligand accepts electron density from the
filled d orbital of metals which results in the increase in bond order of
metal-carbon bond