Once again, I welcome you all to my lecture
series on the chemistry of main group elements.
So far, I was discussing about the classification
of elements and periodic properties and, also
I gave introduction to the type of compounds
we come across among main group elements.
Before I proceed to elaborate the chemistry
of main group elements group wise, let me
bring another important concept that is very
important to understand the physical and chemical
properties and, also the reactivity of main
group compounds, that is structure and bonding
aspects.
Let me begin with Lewis structures. We are
all familiar with the Lewis structures.
In fact, Lewis structure was proposed by Gilbert
Newton Lewis in 1916 from the University of
California at Berkeley. And, besides working
on structural aspects related to main group
elements, he also added lot of information
about electrons in the periodic table and
also he worked on the purification of heavy
water and he proposed acid-base theory and
it is popularly known as Lewis acid-base concept,
and also he worked in the field of photochemistry.
In fact, he was nominated 41 times for Nobel
Prize and unfortunately, he was not given
Nobel Prize, and he found dead on March 23rd
1946, in his laboratory. At that time, he
was working with hydrogen cyanide, a deadly
poisonous gas.
So, what are the postulates one has to follow
before writing Lewis structure? So, one has
to determine, how the atoms are bonded in
a molecule. So, generally if there is only
one of one element and multiple copies of
another element the unique element is central.
Commonly, hydrogen is always peripheral bonded
to elements such as nitrogen, oxygen or sulphur
and then we have to do the counting of electrons.
Essentially, counting the valence electrons
of all atoms present in the molecule or ion
of knowing the basic structure of the molecule
we have to start placing the electrons around
atoms. The first step is to determine the
total number of electrons that are available
for making bonding in Lewis dot structure.
Here, we can use the help of group number
of an element to indicate the number of valence
electrons that it contributes to the molecule.
For example; oxygen is in group 16, it gives
6 electrons.
So, what are the postulates one should remember;
I will repeat again; first, determine how
the atoms are bonded in a molecule. And then,
identify the central atom based on what I
had said so far, and again I am telling you
hydrogen is always peripheral. Count the valence
electrons of all atoms in the molecule. First
establish the bonds between the central atom
and peripheral atoms with a pair of electrons
each and place the remaining electrons around
the atoms to complete their octet; begin with
the most electronegative atom(s). If needed
use double bonds to satisfy the octet. So,
these are the few points, one should remember
before initiating the process of writing Lewis
structure.
Let us consider an example. I will consider
a simple example, such as CO. In CO, we have
carbon and oxygen and we know the electronic
configuration of carbon, that is s2, 2s2 and
2p2 and we also know the electronic configuration
of oxygen 1s2, 2s2 and 2p4. So, while writing
the Lewis structure we should consider only
the valence electrons. So, we have here 4
and here 6. So, total for CO – we have 10
electrons. So, here first write a bond here
using two electrons and now we are left with
8 electrons and using these 8 electrons, let
us first try to satisfy the octet of oxygen.
Now, oxygen has 2 electrons. It requires 6
more electrons; let us place in this fashion.
So, 6 electrons are placed here. We are left
with 2 electrons. These 2 electrons can come
here. Now, if we look into carbon monoxide,
oxygen has satisfied octet, has 6 plus 2,
8 electrons, whereas carbon has only 4 electrons.
So, Lewis structure gives a satisfactory octet
for carbon and oxygen. So, in that case, what
one should do is, we can consider transferring
these electrons between oxygen and carbon
in this fashion.
So, now, if I start counting these electrons
along with these two; oxygen has 8 electrons.
In a same way we can also consider carbon
is satisfied having 8 electrons. So, thus
Lewis structure gives a satisfactory bonding
model for carbon monoxide. So, while writing
carbon monoxide one can write like this or
one can put a triple bond between C and O
and complete it. So, this is how carbon monoxide
has 3 bonds between carbon and oxygen. So,
here the octet is satisfied.
Let us look into another example such as,
[ClO2]- . So, later I will tell you, when
we look into a species by simply looking into
the Lewis dot structure, we should be able
to tell whether it exists in the cationic
form, anionic form or neutral form. So, here
its say anionic form, it is in anionic form.
So, [ClO2]- let us write in the same fashion.
The total number of valence electrons that
are available for making the bonds here, Cl
has 7 electrons and oxygen, totally we have
12 electrons, 6 + 6 and 1 charge is there.
So, we have a total of 20 electrons now. Here,
chlorine is the central atom. It should be
written like this and two are peripheral atoms.
First, we should make a covalent bond by utilizing
2 electrons here, another 2 electrons here.
So, we are left with now 16 electrons and
oxygen being the most electro negative we
should satisfy the octet of this one by adding
6 electrons for both the oxygen atoms.
So, now, we have totally used 16 electrons
and 4 electrons are left. These 4 electrons
I place on chlorine. So, now, we first start
counting all the 3 atoms have octet, that
is, 4 pair of electrons. So, here there is
no need to have any double bond. So, this
is how the structure of [ClO2]-can be explained
using Lewis structure. So, this can also be
written in another format or it can be. So,
one can write either with solid line like
this representing 2 electrons or one can simply
write using the dot structure to complete
this structure. So, this is how the Lewis
dot structure successfully explains bonding
in carbon monoxide and also [ClO2]-.
Let us look into some structures. In case
of some structures, Lewis structure cannot
match the experimental observations. For certain
atoms, do not match the experimental observations.
Let us look into the simple examples such
as formate [HCOO-] ion.
So, in case of formate ions the bond lengths
predicted by Lewis structure are incorrect.
So, according to Lewis structure, this is
how one can write the structure using Lewis
method for formate anion. So, here Lewis structure
predicts two types of bonding for CO; one
is C=O double bond, that is shorter, and one
C—O single bond, that is longer. But the
x-ray structure determination has proved that
both the bond distances are equivalent; that
means, the molecule shows resonance and this
resonance structure can be written in this
fashion. So, the combined structure can be
written in this form.
So, that means, if there are two or more identical
options for a molecule to have double bonds
the molecule will show resonance structures
and these two resonance structures are shown
here and because of this kind of resonance
structure, what happens, both the C O bond
distances are essentially equal, whereas Lewis
dot structure identifies them as 2 separate
entities and one is longer, one is shorter.
That means, there are some limitations as
far as Lewis model is concerned. What are
the failures or limitations of Lewis model,
we shall see now. A number of molecules with
odd number of electrons exist and in case
of Lewis structure we always count in pairs.
So, when odd electron comes, Lewis method
fails to give a satisfactory structure.
For example; if we take NO, in this case we
have 5 plus 6, 11 electrons. As a result,
a satisfactory Lewis structure cannot be written
for NO. An atom may not have enough electrons
to complete its octet without having formal
charges. This problem arises when you look
into the compounds of group 13 elements, where
we have s2p1 electronic configuration.
As a result, what happens, we will end up
with only 3 bonds and 6 electrons, example;
BF3 and another problem or limitation is when
we look into the molecules having more than
8 electrons, for example, sulphur hexafluoride,
where we have 12 electrons and it does not
say. how to put the remaining 4 electrons,
and another important limitation is predicting
the right structure of oxygen molecule. Oxygen
molecule is paramagnetic, but Lewis structure
does not depict this paramagnetic property.
More than all these things, the major drawback
of Lewis model is, it does not give any information
about geometry and shapes of the molecules.
So, in order to determine the molecular shapes
and the geometries of main group compounds
VSEPR theory was postulated. VSEPR is nothing,
but, the valence shell electron pair repulsion
theory. The expansion of abbreviation is self-explanatory;
valence shell, that means, we have to deal
with valence shell and electrons means we
are dealing with electrons present in valence
shell and then the term repulsion comes, that
means, when you are putting electrons into
the valence shell, there is some inter electronic
repulsion and that means, how to minimise
that one.
Valence shell electron pair repulsion theory
essentially gives an idea of how to minimise
the repulsion between the bonded pair as well
as lone pairs; that means, sovalence shell
electron pair repulsion theory is based on
the fact that the electrostatic repulsion
of the electrons is reduced to a minimum,
when the various regions of high electron
density assumes positions as apart as possible.
So, let us look into the concepts that are
used in VSEPR theory, and that means, essentially
it is based on the electrostatic repulsion
between the paired electrons, bonded electrons
and unpaired electrons and how to place them
surrounding the central atom to minimise these
interaction between the lone pairs and bond
pairs, and lone pairs and lone pairs, and
bond pairs and bond pairs.
First, this concept was proposed by Sidgwick
and Powell in 1940 and then in 1957, two more
scientists Gillespie and Nyholm refined it
and introduced the term called steric number
(SN). I would define what the steric number
is. First, we have to determine the steric
contribution on in the central atomby counting
all the electrons very similar to Lewis structure
and then we should arrive at a term called
steric number and steric number is nothing,
but the number of attached atoms plus number
of lone pairs and number of attached atoms
are essentially consuming 2 electrons to make
a covalent bond and then the number of lone
pairs. Once if determine this steric number,
the most stable geometry corresponded to that
can be obtained by maximizing the distance
between the steric points on the surface of
a sphere by keeping them as far away from
each other as possible. So, these are the
basic concepts used in valence shell electron
pair repulsion theory.
So, let us look into few points again. Lone
pairs of electrons; lone pair electrons are
held closer the nucleus. So, most prohibitive
repulsion is lone pair-lone pair, followed
by lone pair-bond pair and then bond pair-bond
pair. So, lone pairs will spread out as much
as possible, that means, essentially lone
pairs will occupy more space compared to bonded
pairs and hence, the angle between the lone
pairs will be always larger.
For predicting the molecular geometries, we
need two important terms. First, we have to
write Lewis structure for a given molecule
and then we have to count the total number
of bonding regions and lone pairs around the
central atom and determine the steric number.
Again, steric number is nothing, but, the
number of attached atom plus number of lone
pairs. So, arrange the bonding regions and
lone pairs in one of the standard geometries
to minimise electron – electron repulsion.
Multiple bonds count as a single bonding region.
So, let us consider the geometries with different
steric numbers. Steric numbers can be anywhere
between 2 to 6. Of course, more will also
come; let us first look into the steric numbers
up to 6.
When the steric number is 2, the molecule
will be linear and when the steric number
is 3, molecule will be trigonal planar, something
like this. So, when the molecule is linear
the steric number 2 and when the steric number
is 3. So, trigonal planar, the angles will
be 120 degree and when the steric number is
4, they are tetrahedrally disposed 
something like this. This is symmetric tetrahedral
with angles of 109.5 degree, something like
this. When the steric number is 5 the geometry
that should be used is trigonal bi-pyramidal,
this is something like this. Here we have
2 types of bonds; one is 90 degrees and another
one is 120 degrees. So, this is trigonal bi-pyramidal
geometry and with steric number 6 octahedral
geometry. So, this is how up to steric number
one can assume the geometries like this.
Molecular shapes: while determining the molecular
shapes we need two important terms, that is,
bond distances and bond angles. The distance
between the nuclei of two bonded atoms along
a straight line is nothing but the bond distance.
The angle between any two bonds containing
a common atom is essentially the bond angle.
So, in this, you consider this octahedral
molecule, the distance from the central atom
to the peripheral atom is called the bond
distance and angle between these 2 atom at
the central atom is called the bond angle.
This is bond angle and this is bond distance.
So, this is trigonal planar molecule with
coordination number three; with this trigonal
planar molecule with steric number 3, bond
angle is 120 and here it is a tetrahedral
and bond angle is 109.5 degree and, here we
have bond angles of 90 and 120, this is trigonal
planar.
And here bond angle is 90. Of course, with
steric number 5 we can have trigonal planar
as well as square pyramidal geometries. So,
one should remember main group compounds do
not prefer to have square pyramid geometry,
until and unless the sixth coordination is
a lone pair. Otherwise preferred geometry
for steric number 5 is trigonal bi-pyramidal.
With coordination number 6, it is octahedral
geometry, with bond angle 90 degree and with
coordination number 7, one can have pentagonal
bi-pyramidal geometry. With bond angles 90
and 72; 72 will be between the equatorial
ones and the axial one will make 90 degree
with the equatorial atoms. So, this is the
pentagonal bi-pyramidal geometry.
So, here I have listed different kinds of
molecules we come across and the corresponding
geometry and the angle with an example in
each case. AX2: steric number 2 and predicted
geometry according to VSEPR theory is linear
and of course, it is linear here, example;
BeCl2, beryllium dichloride, angle is 180
and when the molecule is AX3 type, where A
is the central atom and X are the peripheral
atoms, steric number 3. So, all are bonded
pairs, trigonal planar and example is BF3.
When we have steric number 4, example is AX4,
tetrahedral geometry; example is SiF4, methane
and other tetrahedral compounds having 4 bonded
pairs, here the angle is 109.5 and AX5, trigonal
bi-pyramidal angles are 90 and 120, example;
PF5, phosphorus pentafluoride and with AX6,
steric number 6, octahedral and here example
is sulphur hexafluoride. So, these are about
the steric number that accounts for only bonded
pair.
Let us look in to the combination of both
bonded pairs and lone pairs. We have numerous
examples of having both bonded pairs and lone
pairs. So, in this case, geometry and shapes
will be different, whereas in these cases,
what have shown here the geometry and the
shape are essentially the same. So far, we
have discussed about the molecules having
only bonded pairs starting from steric number
2 to steric number 6 and also 7. In my next
lecture, I will be discussing about the molecules
having both bonded pairs as well as lone pairs.
So, let me conclude today’s talk and have
pleasant chemistry learning and see you in
my next lecture to learn more about the VSEPR
theory.
Thank you very much.
