Resonance Effect, Organic Chemistry Video Lecture - IIT JAM

I'm going to draw a molecule
of benzene.
And then we're going to think
about if anything interesting
might happen with
that molecule.
So let me draw it.
So we have 6 carbons
in a ring.
1, 2, 3, 4, 5 and 6
carbons in a ring.
What's interesting about
benzene, why it's different
than cyclohexane, is that
it has these 3 double
bonds in the ring.
So let's say we have these two
carbons are double bonded to
each other, these two carbons
are double bonded, and then
these two carbons over here
are double bonded.
Actually, I'll draw the
hydrogens here just so that we
remember that they're there.
But I'll it in a subtle color.
So this carbon right
here is going to be
bonded to how many hydrogens?
It has 1, 2, 3 valence electrons
already used up.
So it's going to have one
bonded to one hydrogen.
This one right here,
same thing.
Bonded to one hydrogen.
So it has 4 valence electrons.
This one, same thing.
I think you see the pattern.
That each of these are-- they
have three bonds to carbons,
one single bond to two carbons,
and then one extra
double bond.
And then the fourth bond
is to hydrogen.
So let me just draw all
of the hydrogens here.
I'm doing it in this dark color
so we don't have to pay
too much attention to it.
Now this right here,
this is benzine.
And you're going to see a lot
about benzine in the future.
But in this video, we're going
to study or try to understand
a particularly interesting
property of benzene, and
that's resonance.
And it's not a property of just
benzene, it's a property
of many organic molecules.
But benzene is kind of
the most fun version.
So let's think about what
might happen with this
molecule right here.
So I have this electron.
Let me do that in a
different color.
I have this-- let me
do it in this blue.
I have this electron
over here.
What if this electron moved over
to this carbon over here?
So this carbon is still going
to have the other
electron in the bond.
It's just going to kind of pivot
around a little bit.
So that electron moves
over there.
Now this carbon doesn't need 5
electrons, so this electron
goes to that carbon
right over there.
Now this carbon doesn't need 5
electrons, so that electron
goes back to the original
carbon that
lost that first electron.
So at the end of the
day, everyone has
kind of broken even.
If this happened, we might
end up with a structure
that looks like this.
And I'll draw a two-way arrow
because we can actually go in
both directions.
So let me draw just
the carbon chain.
So 1 carbon, 2 carbons, 3
carbons, 4 carbons, 5 carbons
and 6 carbons.
And then, over here, we had the
double bond over here, but
now it's moved over here.
So now the double bond-- and
actually, let me do it in a
blue color so we see
the difference.
So now the double bond
is over here.
This blue electron has
moved over there.
This blue electron has
moved up here.
Actually, let me color code it,
so it makes it very clear.
So let's say this is
a green electron.
Now the green electron has moved
from this carbon over to
that carbon.
We can imagine that
it's done that.
Then you would have this magenta
carbon or this magenta
electron that was with this
carbon, but now it's moved
over to this carbon over here.
And now the double bond
has shifted as well.
That's what this arrow showed.
We'll stick with the blue
carbon over there.
That blue carbon has moved down
to the original carbon.
And now the double bond
has shifted over here.
So we essentially have a very
similar, really, a very
similar molecule.
This is actually just a rotated
version of that.
But we have these double bonds
that could keep flipping back
and forth between this position
and that position
over there.
They can just keep
on doing it.
They can just keep flipping
either backwards or forwards.
And the reality of benzene is
that it's actually never in
either this structure
or this structure.
It's always, actually, in
something right in between.
The reality of benzene
actually looks
something more like this.
And I'll just draw it without
drawing all the carbons and
the hydrogens.
And obviously, in this case, let
me draw the hydrogens here
since I drew the hydrogens
up here.
This had the hydrogens
over here.
Don't want to forget those.
If you ever forget them
they're implicit.
I want to draw the hydrogens.
But if we just look at the
overall ring, we know that the
carbons and the hydrogens
are implicit.
The actual structure of benzene
is actually in between
that and that.
In reality, you kind of have a
half double bond between all
of the carbons.
So the reality is, is that it
looks something like this.
So you have half a double bond
there, half a double bond
there, half a double bond over
here, half a double bond over
here, and then half a double
bond over here.
And then we're almost done.
And then, half a double
bond over here.
The reality of benzene is that
these electrons are actually
spinning around the
whole ring.
It's not flip-flopping
between this
structure and this structure.
The actual structure, the lower
energy state structure,
is this right here.
Now these Lewis diagrams or
actually, I haven't drawn all
of the Lewis electrons.
But these are considered
contributing structures.
And you often draw these when
you're doing reaction
mechanisms. But the reality
is, is that resonance, the
resonance of these positions
creates-- the reality of
benzene is that it's actually
sitting in this
intermediate position.
Now, this doesn't happen
only with benzene.
Another example, there's going
to be many examples.
But just so that we're familiar
with maybe two of the
best examples, another example
that you'll see a lot in the
context of resonance is
the carbonate ion.
So carbonate ion.
You have a double bond to one
oxygen and then you have
single bonds to two
other oxygens.
And those two other oxygens
have extra electrons.
So if I were to draw this oxygen
over here it has 1, 2,
3, 4, 5, 6 valence-- or
actually, I should
say, it has 7 valence.
Let me make it very clear.
So it has 1, 2, 3, 4, 5,
6, 7 valence electrons.
It has one extra electron, so
it has a negative charge.
And the same is true
for this one.
It has 1, 2, 3, 4, 5, 6,
7 valence electrons.
One extra.
So it has a negative charge.
If you were to just look at
this, I guess you could call
it this resonance structure or
this contributing structure,
you'd say hey, maybe this
oxygen-- and this oxygen here
is neutral, so it has
6 valence electrons.
1, 2, 3, 4, 5, 6.
Maybe, just maybe, one of these
electrons can be given
to the carbon and then the
carbon would lose an electron
to this guy on top.
So maybe you could imagine a
situation where this electron
right here gets given
to the carbon.
And when that gets given to
the carbon, the carbon
releases-- it all happen
simultaneously.
The carbon releases this
electron and it goes back up
to that oxygen over there.
And so what's that going
to look like if
that were to happen?
So if that were to happen,
now our structure
will look like this.
We have a carbon.
Now this carbon only has
a single bond up here.
And then we have our oxygen.
The oxygen, it had its
6 valence electrons.
1, 2, 3, 4, 5, 6 valence
electrons.
But now it got this
extra blue one.
And now it got this extra blue
one, so now it has 7 valence
electrons, and it has
a negative charge.
Now this oxygen over here
gave one of its
electrons to the carbon.
Now it is bonded with it.
So now the carbon has
a double bond.
I'll actually do it
in that color.
Has a double bond with this
oxygen down here.
It gave an electron, so now it
only has 6 valence electrons.
1, 2, 3, 4, 5, 6.
And it is now neutral.
And this oxygen over here,
nothing really
new happened to it.
I could just copy
and paste it.
So let me copy and then
let me paste it.
So this one is just sitting
right like that.
But you could imagine a
situation where this oxygen
right here, then all of a
sudden-- and it could have
come from this oxygen up here
or it could come from this
oxygen right here.
This oxygen says hey, I have
an extra electron.
Let me give it to the carbon.
And then the carbon releases a
double bond with one of the
other oxygens.
In this case, it would
be this one.
Let me draw it.
So maybe this electron right
here gets given to the carbon.
Forms a double bond.
Then the carbon can let
go of an electron.
And so this electron right here
goes back to this oxygen.
And so what happens?
So if that were to happen, our
structure looks like this.
We have a carbon single bonded
to an oxygen up here that has
1, 2, 3, 4, 5, 6, 7
valence electrons.
That hasn't changed in-- we
could call it this resonance
reaction, or however you
want to call it.
So it still has a
negative charge.
We have this guy down here.
He took his electron back.
So now he has 7 valence
electrons again.
So 1, 2, 3, 4, 5, 6, 7 valence
electrons again.
And I can even show the
one that he got back.
That one's in purple.
So he now has a negative
charge.
And this guy now gave an
electron to the carbon.
So he forms a double bond,
a new double bond.
So this guy forms a double
bond with the carbon.
He gave an electron, so he
only has 1, 2, 3, 4, 5, 6
valence electrons, and
is now neutral.
Now these can all keep swapping
between each other.
You can even go from this
structure to that structure.
You can actually go from any one
of these structures to any
of the others.
And the reality of the
carbonate ion, let
me write this down.
This is the carbonate ion.
The reality of it is that its
true structure is some place
in between all of these.
So the true structure
of a carbonate ion
would look like this.
You would have a carbon and
you'd have three oxygens.
They have at least one
single bond with each
of those three oxygens.
And then you have 1/3 and then
you have-- actually, I should
say, you have 1/3 of a double
bond with each of them.
This is a 1/3 of a bond.
This is not standard notation,
but this is essentially going.
1/3 of the time, the electron
is on that bond.
And then the other 2/3 of the
time, each of these oxygens
have an extra electron.
You could imagine almost having
a negative 2/3 charge.
Now, people normally draw one
of these structures because
this is a nice-- kind of you're
dealing with whole numbers.
But the reality of carbonate
ions is that it's experiencing
this resonance.
That the electrons are actually
always floating in
between these forms. Actually
floating across
all of these bonds.
And that actually, makes this
molecule more stable.
This is at a lower energy state
than any of these forms.
And the same thing is
true with benzene.
This right here, where we're
in between these two
structures, is actually at a
lower energy state, a more
stable state than either
of these forms.