talking about the theories that explain
bonden chemical bundles and we said
there are several different theories
which try to explain the bonding
chemical bonding in molecules and also
the structure and one of them which we
will simply cover in this chapter is the
valence bond theory and the second one
we said is the molecular orbital theory
so we talked about the valence bond
theory and again the valence bond Theory
remember says that a chemical bond
between two atoms will form if the two
atomic orbitals from those two atom
overlaps and also provided that those
two atomic orbitals has one electron in
each and so when you have an overlap of
two atomic orbitals which has one
electron each and those two electrons
simply pair up between the bonds and a
bond is formed this is what how the
valence bond Theory simply explains the
abundant between two atoms as a matter
of fact and here the example if you
consider the formation of the hydrogen
molecule and the hydrogen hydrogen bond
in the hydrogen molecule so 1s orbital
from one hydrogen and 1s from the other
hydrogen which contains one electron in
each actually overlaps to form a bond
between two hydrogen atoms and also form
the molecule or if we have for example
the bond between hydrogen and chlorine
the valence bond theory says that the
atomic orbital of hydrogen 1s orbital
and the atomic orbital of chlorine 3 P
orbitals each has one electron remember
again those orbitals that overlap lead
into the bond should have one electron
in each as a matter of fact so we have a
bond between hydrogen and chlorine as we
have seen here or in a chlorine molecule
if we examine simply the bond between
two chlorine atoms
according to the valence bond Theory 3p
orbitals from the chlorine atoms overlap
which contains one electron in each
again then for bond simply form this is
actually nice but the problem was again
when it comes to explain the geometry of
the molecule and the bond angles
particularly which goes with the
geometry the valence bond Theory has
problems as try to as I try to explain
you for example if you consider the
beryllium difluoride molecule and
beryllium actually doesn't have any war
beetle which have one electron that
overlap and form bond so can't really
explain how the linear beryllium
difluoride molecule forms simply a bond
between fluorine and beryllium atoms and
the same thing is also true when it
comes to explain the angle between the
hydrogen oxygen bond in water molecule
actually and the valence bond theory
predicts that if it is purely overlap of
the atomic orbitals pure atomic orbitals
then the bond angle in water should be
ninety degrees and the same is also true
for ammonia as we have shown last time
and there also the angle has to be 90
degrees but in reality the angles are
not 90 degrees right
in both cases actually larger than 90
degrees and particularly when you try to
explain the reaction between hydrogen
and carbon actually according to the
valence bond Theory the molecule the
product that should form there should be
carbon one carbon atom two hydrogen
atoms simply CH 2 molecule just
considered in the overlap but again we
know that when you react hydrogen and
carbon the methane molecule form simply
ch4 molecule forms and it's a
tetrahedral so the only spawn Theory
cannot explain that right and we said
the valence bond Theory proposed what we
call hi
predation of the atomic orbitals in a
molecule to explain the geometry and the
angles actually and for example if we
simply try to look at the beryllium
cases beryllium difluoride molecule
they're simply you have one s two and
two s two electron configuration on a
beryllium and the hybridization model
which is proposed by the valence bond
Theory used by valence bond Theory
simply says that one electron from the
2's orbital of the beryllium goes to the
2p orbital and this way I have simply
one electron in two years of one
electron and 2p orbitals so those two
mixes together to form two hybrid
orbitals and again it says these hybrid
orbitals simply poets toward the
specific specific direction saw that the
electron-electron repulsions between
them will be minimal and when it simply
comes that in picture we know that to
get minimum repulsion between two
electron pairs or electron domains they
have to be simply in a linear fashion
right on the same line 180 degrees and
so on again like if you have simply df3
molecule in that case you have on boron
sp2 hybridization that simply means you
mix 1 2's orbital on boron and 2 in this
case P orbitals and by mixing those you
end up with three hybrid orbitals three
hybrid orbitals again you see what we
said last time the number of atomic
orbitals that you mix to form hybrid
orbitals must always be equal to the
number of hybrid orbitals that you end
up so if you mix two orbitals you end up
with two orbitals if you mix three you
end up with three so in the boron
trifluoride cases you have simply sp2
hybridization that means 1 s and 2 P
orbitals mixes to form 3 SP
two hybrid orbitals and these hybrid
orbitals also points toward the corner
of an equilateral triangle why to get
minimum reports minimal portions and now
these hybrid orbitals valence bond
theory says that overlap they overlap to
form a bond one hybrid orbital on a
central atom and atomic orbital on a
terminal atom overlaps to form a bond
and then you can also this way actually
by hybridization model valence bond
Theory explains the geometry of the
molecule that means the angle it becomes
exactly the same as the electron pair
valence shell electron pair repulsion
model right there also you try to find
out the number of electron pairs or
electron domains around the central atom
and then depending on the number of
electron domains those electrons will
simply point in three-dimensional space
in a such a way to get minimum
repulsions and those are if you have to
it will be linear if we have three
actually electron domains it will be
equal eighths with triangle shown simply
and if you have four tetrahedral 5/2
you're gonna lie pyramidal 6 octahedral
we have simply seen that as a matter of
fact before so this is what how the
hybridization explains the angle and the
geometry that is used by the valence
bond theory now let's see what we have
here in slides so again it says hybrid
orbitals from mixing of atomic orbitals
now we have to remember this this is
very important this mixing of atomic
orbitals these atomic orbitals should be
all on the same atom not the orbitals
atomic orbitals from different atoms
same atom right same atom and simply
they create orbitals equal energy called
degenerate degenerate simply means the
same energy these hybrid orbitals when
two orbitals mixes they create orbitals
when three orbiters mixes they create
three orbitals that means the number of
orbitals that you mix will be equal to
number of hybrid orbitals this is again
we talked about there's the SP
hybridisation on beryllium for example
in beryllium difluoride molecule and
here it shows how we simply get the s
and P 1p actually and one s hybridizes
and it forms the SP hybrid orbitals 2 SP
hybrid orbitals and this is again the
shape or how you can visualize these
hybrid orbitals the formation of hybrid
orbitals remember again this is again
3-dimensional this is a loop over here
loop over there too
so you mix those two atomic orbitals and
beryllium and you form two hybrid
orbitals one of them is this the other
one is this this one simply shows in
this direction and that one in that
direction that simply means these two
hybrid orbitals are this one is that
that one is this actually right so now
they have one electron in each and now
they can overlap with the terminal atoms
in this case which is the fluorine right
fluorine P orbitals which contains one
electron to form the hybrid orbitals and
the geometry in that case will be linear
just for the same reason like in the
electron very shell electron pair
repulsion theory to get the maximum
overlap simply the direction of the
terminal atoms that will approach to the
central atom is determined by the
positions off of the hybrid orbitals
that they will overlap right so this is
simply explains how the formations of
beryllium difluoride
by the overlaps and the hybridization
now if you have simply bf3 molecule in
that case 1 s and 2 P orbitals mixtures
to form three sp3 hybrid orbitals and
where are they they are over here this
is 1 this is 2 this is 3 how they
arrange themselves around the central
atom to get minimum repulsions like this
right again this is an equilateral
triangle and then the fluorine will
simply come from there and overlap then
bf3 will be
the planar trigonal planar molecule
geometry off the bf3 by the
hybridization now when you go to the
methane molecule formation of methane
molecule in that case one s and 3p
orbitals are hybridized right but and
you end up with four hybrid orbitals
these are simply called sp3 hybrids and
these sp3 hybrids arrange themselves
around the carbon in a such a way to get
minimum repulsion that means it will be
tetrahedral right tetrahedral geometry
and these are 109.5 actually and the
hydrogen's will simply come from that
directions to overlap with the hybrid
orbitals to form the bond right the
chemical bond and the molecule will be
simply tetrahedral molecule now in the
case of water here again we have an
oxygen sp3 hybridization right sp3
hybridization what but the oxygen has
simply a number of if we consider a
number of six valence electrons right
and two of those in 2s and four of those
are 2p orbitals so now we have to
remember the number of electrons that
comes come from the orbitals that we mix
to get hybrid orbitals will be in this
case six electrons so when we have
simply four hybrid orbitals around the
oxygen which simply points were the
corner of a tetrahedron then we put the
electrons in these hybrid orbitals how
many electrons we have six so we put one
electron in each hybrid orbital then we
have two more and then we have to fill
out the another hybrid orbital here two
and another two there again remember
just like atomic orbitals you can only
put two electrons maximum in a hybrid
orbitals two electrons and they are also
appeared so as a result actually if you
try to visualize the water
the molecule in the water then on the
oxygen I have four SP three hybrid
orbitals and two of those have two
electrons each and two of those has one
electron each now the orbitals hybrid
orbitals which have two electron each
cannot overlap with another atomic
orbitals remember that will be simply
three electron in the same one orbital
that's not possible
therefore the hybrid orbitals that will
make bond on the oxygen are the one
which contains one electron this one and
that one and when they have one electron
and they overlap with the hydrogen one
as here and why hydrogen 1s here form
the bonds right form the bond overlap
and those simply electrons pair up here
and since the orientations or the
geometry of the hybrid orbitals are
tetrahedral then the angle here will
also be close to the tetrahedral angle
in the case of water actually remember
it's 105 degrees less than tetrahedral
exactly but we know those are the effect
of those so this is how the
hybridization and the valence bond
Theory explains the bundle let's go to
now what happens if you have more than
four chemical bond around the central
atom right more than four electron pairs
after you simply hybridized orbitals you
put the electrons in the hybrid orbitals
if you need more than four orbitals to
overlap to form bond then you start
mixing now the D orbitals with s and P
for example you have one s orbital three
P orbitals if you mix those you get only
maximum is what four hybrid orbitals
that means if there is only for bond
around the central atom then it's fine
you can use those but if there's five
buns then you need another hybrid
orbitals so in this case we simply mix
one s and 3p
Wandy and it becomes SP d sp3d2 again
remember always the number here is 1
this is 3 means 3 P orbitals that's what
it means
now one simply means one D orbitals how
many orbitals we mix here actually five
atomic orbitals as a result we form 5 SP
3 D hybrid orbitals 5 of those now they
will simply point toward the corner of
trigonal bipyramidal geometry trigonal
bi-pyramidal geometry why to get the
minimum repulsion between the electron
pairs actually or electron domains you
can look at that and this is simply how
hypervalent molecules the bonding and
hypervalent molecules are explained
actually now if you have this is like
pcl5 right molecule now if I have an SF
6 molecule in this case I need 6 hybrid
orbitals unli sulfur to form simply 6
bonds right with 6 fluorine now the
hybridization in this case becomes sp3
and D 2 right sp3 and D 2 what does that
means 1 s 3 P and 2 D orbitals total of
6 atomic orbitals simply mixed and we
get 6 SP 3 D 2 hybrid orbitals now each
of these hybrid orbitals there again
points toward the corner of an
octahedron and the fluorines actually
will come in that directions to overlap
with these atomic or hybrid orbitals to
form a bond and as a result geometry of
the molecule becomes also octahedral
right okay he drew so this is how the
valence bond Theory simply explore
the bonding by hybrid explains the
geometry actually by hybridizations this
again shows the hybrid orbitals sort of
a summary it says simply draw the Lewis
structure use the valence shell electron
pair repulsion theory to determine the
electron domain geometry and then
specify the hybrid orbitals needed to
accommodate these electron pairs now
let's say let me give you an example
let's work out an example and consider
the phosgene molecule which is carbon
oxygen and Cl 2 molecule right all right
now by using the valence shell electron
pair repulsion theory we can easily
predict the geometry of this molecule
what we have to do first we have to
write down the lewis structure of this
and the lewis structure again as follow
you have the oxygen bonded to the carbon
double bond and then we have the
chlorine here and the chlorine here this
is a phosphine remember it's a very very
poisonous gas very poisonous actually
this is the Lewis structure now from the
Lewis structure I can simply predict
what will be the geometry of this
molecule what I should do I find out how
many electron domains or how many
electron pairs around the central atom
and in this case one two and three
electron domains or electron pairs and
what does that means that means that the
electron pair geometry let's say
electron pair geometry will be three
Kornel planar right trigonal planar
three guano planar now then the
molecular geometry will also be trigonal
planar right because there's no lone
pair electrons
no lone pair electrons so molecular
let's say molecular geometry will also
be trigonal planar they say now again
the angle here this angle is less than
120
greater than hundred and twenty degrees
I should say greater than 120 degrees
and this angle is less than 120 degrees
and we know that again because this two
electron pairs double bond actually
occupies more space and pushes the
bonding pair between the carbon and
chlorine closer together making this
angle less than hundred twenty degrees
this is a little bit greater than
hundred and twenty degrees this is what
we analyzed by the electron pair valence
shell electron pair repulsion theory
right we can get that informations now
let's look at this in terms of valence
bond Theory how the valence bond Theory
explains the bonding between the atoms
and what how the job geometry is also
predicted by the valence bond Theory
again what we will do is this around the
carbon right we have three electron
domains electron pairs now what does
that means three electron pair and three
electron domains in terms of
hybridization what kind of hybridization
I need to accommodate those three
electron pairs sp2 because I need three
hybrid orbitals three hybrid orbitals
will be obtained by mixing 1 s and 2 P
therefore the hybridization on carbon
let's put over here
valence bond Theory right first is
nation on carbon is sp2 type of
hybridization now what does that means
mixing one 2p orbital one yeah 2p
orbitals 1s orbitals remember there are
three P orbitals what does that means
that means there is one p orbitals which
is not involved in hybridizations right
not involved in itemization still there
empty or may have a net one electron
actually in this case it has a one
electron and it's there now how about
the hybridization around the oxygen hmm
SP no if you have sp3 then we actually
we don't need that why because we need
again we have three electron domains
that's it for three electron domains I
need three hybrid orbitals three hybrid
orbitals will be obtained again by s 1 s
and 2 P orbitals s P 2 so hybridization
on oxygen is also sp2 like this right
now this way actually when you have this
hybridization the electron pairs will be
on the corner of a trigonal planar in an
equilateral triangle that means there's
going to be minimum repulsion between
them right now let's say consider how
the bonding will take place now here the
carbon alright put that over here carbon
on the carbon I have three hybrid
orbitals SP two hybrid orbitals let's
put those where they will be one will be
here and the other one will be here
and the other one will be here remember
there are also little side on the other
here too for these not only
one side so these are three sp2 hybrid
sp2 sp3 sp2 angle 120 degrees right they
have to be simply form equilateral
triangle or they should simply point
toward the corners of an equilateral
triangle now on the oxygen here I also
have sp2 hybrid orbitals hybridization
and now I have simply overlap that and
then I have another sp2 hybrid orbital
here another sp2 hybrid over here so I
have simply one electron here from the
carbon one electron from the oxygen we
have the bunt right overlap awfully this
is let's say sp2 hybrid orbital this is
also sp2 hybrid orbitals how about this
this is also a speech ooh how about this
this is also sp2 and now you put three
electrons there right and two electrons
there like this and now in the Quarian
cases the chlorine will come from here
chlorine remember has one electron in
three P orbitals so it can simply use
that chlorine three P orbitals like this
and overlap and form a bond and another
chlorine here and uses the three P
orbitals this is actually three P
orbitals of chlorine and that's also
three P orbitals of the chlorine I'm
just trying to show you how we can
simply analyze the bundles and also
count the electrons coming from each
atoms so I have three electrons there
that is another bunt here overlap of the
hybrid orbital and pure 3 P orbitals
hybrid orbital 3 P hybrid orbital hybrid
orbital these are the hybrid orbitals
now when we count the valence electrons
and the P and the atomic orbitals
valence atomic orbitals here we have
made sp2 hybridization
and carbon and what does that means we
used 1 s and 2 P orbitals valence
orbitals that means there's one p
orbitals which is not involved in
hybridization what does that means it is
pure P orbitals perpendicular to this
play over here and it has one electron
right p orbital here perpendicular like
this and one electron
the same thing is also a true on the
oxygen on oxygen we also have the sp2
hybridization and there is one p orbital
which is not involved in hybridizations
as a result those two P orbitals P
orbitals here P orbitals they're on this
carbon and that oxygen they overlap
sidewise let's try to show you here just
to visualize how they will overlap this
is the carbon this is the oxygen and I
have simply ap orbital here
perpendicular to that plane and this p
orbital there perpendicular to that
plane and they now overlap like this to
form a bond by overlapping this is 3p
orbital pure and this is also true P
orbitals no this is also 2p orbits as
pure 2p orbitals and we have the overlap
the picture more looks like this carbon
oxygen and we have simply an overlap a
region where the electron there actually
and another region here and that is AB
bond and this bond is the second bond
here the first one is this second one is
this one and it is simply obtained by
the overlap of P orbitals on this carbon
and P orbitals on this oxygens which are
not involved in hybridization sideways
to get the bond now we name these two
kind of buns depending on their
symmetries and the symmetry of the this
bond here has a rotation
symmetry that means if I simply rotate
this around this axis nothing will
change those kind of bonds are called
Sigma bond right so this is simply a
sigma bond now if you ask me weird comes
from just name
that's a sigma bond Sigma bond me say
Sigma bond we know that it is it has a
rotational symmetry around the bound
axis that means if you rotate it around
the bond X as nothing will change right
has that symmetry or cylindrical
symmetry sometimes I said that's the
symmetry now this one is pi bond this is
called pi bond and the pi bond again do
not have rotational symmetry about the
bond axis right that means I have a loop
here another loop there you can call
this like the sausage here or sausage
there right and if you rotate it doesn't
have a symmetry right you can tell that
you rotate it so these kind of bond
which do not have rotational symmetry
about the bond axis is called pi bond
right PI bonds so now we classify the
bonds as Sigma bond and PI bond in
valence bond Theory and we know which
one is Sigma which one is pi now in
other things that we should always
remember the Sigma bond is always
stronger than a PI bond because the
Sigma overlap that means this fashion
overlap is always greater than this
fashion overlap in fact in the molecule
and other things we should always
remember that if I have a double bond
between two atoms the first one of those
actually double bond is a sigma bond the
second one is always PI bond right now
if I have three bonds between the two
atoms one is Sigma
the other two are PI bond
right now let's
how we can simply see those bonds they
will be here types of bond again we have
simply Sigma and PI type this is exactly
what I try to show you here and now the
pictures are follow right and when you
simply overlap these two orbitals like
this and this fashion you get as Sigma
bond why because this kind of overlap
has rotational symmetry around those
axes if you just rotate this around this
axis you wouldn't tell that it is
rotated right that's what symmetry means
symmetry axis means or symmetry
operations I should say now here this is
a PI bond
now there is one p orbital another p
orbital and they are sideway overlaps of
the 2p orbitals you call a sideway like
this and the result actually this is
what I tried to show you here this is
the PI bond and structure more looks up
probably the shape more looks like this
and it doesn't have rotational symmetry
around the bottom axis doesn't have a
rotation symmetry around the bond axis
okay now these are in other examples
this is the case where you have let's
say ethylene molecule and ethylene
molecule has a double bond like this so
if you try to analyze this double bond
here as we have done here one is Sigma
has to be the other one is PI now the
Sigma bond is formed which hybrid
orbital on this carbon tell me which
hybrid orbital on the carbon is used to
overlap sp2 right why because around the
carbon I have three electron domains
that I should have three hybrid orbitals
to get three hybrid orbitals I have to
have sp2 hybridization right
sp2 hybrid orbital from this carbon SP 2
hybrid from that carbon overlaps to
firma bomb just similar to this right
similar to that actually and now how
about the second one second one is the
PI bond as shown there now what are the
orbitals used to form this carbon
hydrogen bond on the carbon it uses sp2
hybrid orbitals on the hydrogen 1s is it
Sigma or Pi Sigma yes
single bonds are always Sigma remember
that and you can also visualize that by
considering the symmetry of the overlap
orbital actually and so that you can
simply analyze it nicely now we need to
have c2h2 acetylene molecule and in that
case the second bond here is PI bond the
actually that PI bond is formed by the
overlap of the P orbitals on this carbon
and that carbon and those P orbitals are
pure Peter McCord's they are not
involved in hybridizations right they're
not involved in hydrogen positions so in
the case of simply acetylene first of
all let me ask you this what kind of
hybridization I should have on carbon SP
just SP again two electron domains one
this and one here SP that's it so SP
horrible orbital used here to form the
bond sigma bond and used here one of
them with the SP again on this carbon to
form a sigma bond and the other two are
PI bond but those P again they are
obtained by the overlap of the pure
atomic orbitals which are not involved
hybrid on the carbons right and they are
those P orbitals are if two are in this
perpendicular to the plane two will be
in this direction again perpendicular
remember the p-orbitals are always
perpendicular to each other three of
those so as a result we end up with two
P orbitals and one Sigma orbitals and
the shape of those simply as here this
is one p orbital here and the other p
orbital there that means one of the P
orbitals is on this plane the plane of
the curtain here and the other P
orbitals is on a plane which is
perpendicular to the curtain here right
so two P orbitals one like this the
other one is perpendicular to that and
that's the triple bond
now sometimes the double bond
particularly or the electrons which are
in the PI bond I should say are not
localized between two atoms but they can
change simply positions in the molecule
and these are simply called delocalized
localized means they're always there
delocalized means they can change their
place right that's or space with a game
now let's look at the examples examples
are the benzene molecule in a benzene
molecule this simply this picture shows
you the Sigma orbitals Sigma bonding I
should probably say let's put it this
way Sigma type orbitals right in bonding
between the atoms carbon and hydrogen
and between the carbons like this okay
let me ask you this what will be the
hybridization if I ask you what is the
hybridization on each carbon in benzene
sp2 1 s and 2 P orbitals are mixed and I
have three hybrid orbitals right and
those are sp2 hybrid orbitals but and
then remember there's one more P
orbitals and that's not involved in
hybridization and that orbital is
radicular to the plane here on each
carbon and shown like this right these
are the P orbitals on each carbon that
not involved in hybridization right now
they can form pi bond by overlap in
sideways like this two to three PI bonds
but the PI bonds may be between those
two those two and those two or maybe
here and there and there and they can be
simply in any place so I can simply show
those PI bond 3 PI bonds actually in
benzene molecule this fashion which says
that it is delocalized pi bond right
it is delocalized pi but not between two
carbon atoms but goes over all the
carbon atoms and we indicated as in this
picture again that D localization PI
bond D localization
ok now the next Bond in theory that we
will talk here is the molecular orbital
Theory molecular orbital Theory and it's
all mo the abbreviations mo theory now
what is an mo Theory mo mo theory is
quite different than the valence bond
Theory as I will try to explain you
first of all we have to remember now the
mo theory says that the electrons of the
atoms in the molecule if atom is in a
molecule are in molecular orbitals right
if you if you think about the the atom
electrons are in atomic orbitals right
what kind of atomic orbitals s orbitals
P orbitals D orbitals F orbitals
dependents what element how many
electrons you have so those are the pure
atomic orbitals the molecular orbital
Theory says when you form a molecule now
you have molecular orbitals and those
electrons will
into the molecular orbitals and again it
says the molecular orbitals just the
same as atomic orbitals can accommodate
maximum of two electrons each so same as
the atomic orbitals and also it says
when you put two electrons in a
molecular orbital they have to pair up
that means spin off one electron is plus
1/2 the other one should be minus 1/2
right now
the question is how we can form these
molecular orbitals what are they right
what we mean by molecular orbitals now
again the theory says the molecular
orbitals in a molecule are obtained by
linearly mixing off the pure atomic
orbitals in the molecule now which I
tell me Corbett says these are these are
the atomic orbitals of the atoms which
present in the molecule which present in
a molecule right so so we we have simply
molecule and then we have the electrons
which simply forms the molecule for
example think about the water case or
just let's say the hydrogen molecule
right in hydrogen molecule again
molecular orbital theory says there two
electrons right coming from each
hydrogen these two electrons are in
molecular orbitals so and this molecular
orbital is simply obtained by linear
combination of the atomic orbitals of 1s
what from one hydrogen and 1s from the
other hydrogen right so that's how the
molecular orbital Theory approaches to a
molecule and abundant as I will try to
show you as a matter of fact so when we
simply write down the electron
configuration of an atom what we use we
use the atomic orbitals s P D F right
now according to the molecular orbital
if you have a molecule you can also
write down the electron configurations
of the molecule but the electrons now
will go into the molecular orbitals
molecular orbitals now let's let me just
say how we simply combine the I guess we
can show over here yes this it says
simply they different from atomic
orbitals because they represent the
entire molecule not a single atom
whenever two atomic orbitals on one on
one on each atom linearly combine
linearly combine two molecular orbitals
are formed one bond in one anti-bonding
bonding orbitals are constructive
combinations of atomic orbitals
anti-bonding orbitals are destructive
combinations of atomic orbitals
remember again and they're also nodal
plane means that the electron density is
zero remember again we said that the
orbitals are actually wave functions
right and those are simply the wave
functions that describes the electron
now when you have the wave function that
means you describe in a wave so you can
simply combine two waves linearly to
obtain simply a constructive type
combination interference or destructive
type interference right
remember again combine in two waves if
they are in phase then we have
constructive interference intensity of
the wave increases now if you have out
of phase than in that case the
destructive interference means that the
waves cancel each other right we have to
remember so this is more like that now
in the case of hydrogen I have simply 1s
let's call that hydrogen a and hydrogen
B here we have a bond this is the
molecule and in the hydrogen a what I
haven't won is a orbital and the other
one is let's call that one s B orbital
now in the
molecule molecular orbital theory says I
have two molecular orbitals one of them
is this mo which is equals to one s a
plus 1sb now this is a linear
combination that means I simply add up
to atomic orbitals pure atomic ellipses
two waves to get a molecular orbitals
now the second one I can also combine
those as follow 1s a minus 1s simply be
right these are again atomic orbitals so
this corresponds to constructive type
interference between two waves and this
corresponds to destructive type
interference between two waves now this
orbital is called bonding molecular
orbital bonding mo molecular orbital now
this one is called anti-bonding mo
anti-bonding mo anti band anymore and
usually the anti-bonding molecular
orbitals is indicated by a steric like
this on the top just to indicate that
said anti-bonding now what really
happens again as follow we have two
electrons in hydrogen molecule each
coming from each atom now the electron
goes into the molecular orbitals and
three electron in one molecular orbital
here and two electron there now since
this is a constructive interference
combinations and linear combinations of
the atomic orbitals any electron that
goes into this orbital stabilizes the
hydrogen molecule that's why it's called
bonding that means it keeps the two
hydrogen atom together any electron
which goes into anti-bonding try to
separate actually the two hydrogen atoms
that's why it's called anti-bonding so
electron in bonding molecular orbitals
will have high
probability of being between the two
atoms that are bundled in the molecule
the electron which is anti-bonding have
greater probability being more out of
the simply the inter nuclear distance
out again and between the two atoms we
have what we call the nodal plane this
is the nodal plane where where the
electrons density or probability of
finding electron is zero now let's see
all are here this is all what we said
again two 1s orbital combines this
fashion this one here linear combination
constructive combination and the bonding
molecular orbitals now remember again
the molecular orbitals comprises or
expand overall molecule it is not only
around an atom but it is all the atoms
in the molecule right so in this case
this molecular orbital simply expands
overall overall molecule and includes
both of the nuclei here two hydrogen
atoms like this and this is a
constructive interference now if you
look at the electron density in the
bonding molecular orbital which
corresponds to the bonding greater
electron density is between the two
nuclei right now this one is n-type
bonding that's a destructive type
interference of two waves you can think
about that way and now it it is this
this is the linear combinations one si
minus 1 SB now here the electron density
is more out of the inter nuclear
distance over here and there now I have
a nodal plane that means if I simply
pass a plane they are perpendicular that
will have zero electron probability
actually no electron density there now
let's look at the energy aspect of these
molecular orbitals this is the energy
aspects if I try to show you how the
energies of the molecular orbitals with
respect to the
pure atomic orbitals and show them on a
diagram like this it's called the
molecular orbital energy level diagram
actually shows the relative energies of
the atomic orbitals and the molecular
orbitals so what we see here is this
this is the 1s a I call that one si for
example suppose that hydrogen atomic
orbital periodic orbital this is the
other one they have the same energy
right
they're both hydrogen atoms in fact so
have the same energy now when they
combine they can form a molecular
orbital bonding molecular orbital which
is lower energy than this over here
right and this is a linear constructive
type combinations which corresponds to
the bonding molecular orbital that is
the bonding molecular orbitals and it
has a sigma 1s it is called Sigma 1s why
first of all it's called Sigma because
it has rotational symmetry around the
bond axis right just like the Sigma
orbital in the valence bond Theory by
the overlap and that's why it's called
Sigma and this one is called Sigma
bonding and this one is Sigma
anti-bonding that means it is this and
that and you see again it's called 1
Sigma 1s here that simply means that
molecular orbital is obtained by the
linear combinations of pure atomic 1s
orbital that's what it means
ok now as a result when you form the
bonding molecular orbital here molecular
orbital anti-bonding here then you put
the electrons now in molecular orbital
they're not anymore
in atomic orbitals molecular orbitals
now how you can place the electrons into
the molecular orbitals just the same as
we place the electrons in the atomic
orbital each molecular orbital will have
two electrons and then we start with
from the lowest and we fill it out right
in this case we only have two electrons
to put and we can put it over here Sigma
bonding Sigma 1s bonding orbital
this and this becomes empty right now
what really happens as a result of this
again molecular orbital Theory says and
we can also see here that the hydrogen
molecule is more stable than the two
hydrogen's infinitely apart right that
means when you have no pondan no
interactions between them how much more
stable twice as this energy what that
energy corresponds to because when you
have one electron and you take that
electron from there and put it there
what does that means that much energy is
released the system get that much more
stable and now another electron moving
in fact from there to there that much
energy is also released so that means
that the hydrogen molecule here is more
stable than the two hydrogen's
infinitely apart and not interacting
with each other now the question here is
how about the bond
can we talk about the number of bonds
right like in the valence bond Theory
single bond double bond triple bond and
so on know in molecular orbital theory
we never talk about the bonding or
number of bonds between the atoms at all
but what we simply have just to indicate
the overall stability of the molecule
the bond order and the bond order is
defined as the number of electrons in
bonding atomic orbitals minus the number
of electrons in anti-bonding atomic
orbits is divided by two and that gives
you what we call it bond order and if I
apply it here for the hydrogen molecule
number of electrons in bonding is two
number of electrons in an anti-bonding
zero divided by two is one so we say
that the bond order in a hydrogen
molecule is one that's what the
molecular orbital Theory says does that
corresponds to a single bond not
necessarily as we will see and we will
see that the bond order may be
fractional numbers it may be like 1.5 it
may be 3.5 right
so but what we can say that the higher
the bond order in a molecule the more
stable the molecule is what does that
means that means the stronger the
chemical bonds between the atoms in a
molecule you see how the bond strength
and the overall stability of the
molecule is expressed by the bond order
or related to the bond order the higher
the bond order means more stable the
molecule the higher the bond order
simply means the more electrons you have
in bonding orbitals and less electrons
in anti abundant because anti bond and
if you put one electron from there to
there you have to put that much energy
to the system and that's destabilizes it
as a matter of fact right so this is how
the molecular orbital Theory explains
the bonding between the atoms by using
the molecular orbitals again and the
bond order yes any questions
I guess the question is dismissed Thanks