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- 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.