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.