Oldcroft

Transcription

We now know how to name aldehydes and ketones, and what I want to do in this video is show a mechanism for actually creating one. In particular, we're going to create a ketone. So let's say we've got ourselves some benzene, and in the first step of this reaction, the benzene is just going to sit and watch. We've got some benzene and we've got some of acetyl chloride. So it looks almost like an aldehyde or a ketone, but instead of having a carbon chain or a hydrogen, we're going to have a chlorine atom right over there. So this is acetyl chloride, sometimes called acyl chloride. This is acetyl chloride, and we're going to have an aluminum chloride catalyst. And a catalyst means that it participates in the reaction, but it enters the reaction and it exits the reaction as the same molecule. So it just catalyzes it. It doesn't disappear. It just changes halfway, but then goes back to what it was before. So we have some aluminum chloride and it's bonded to one, two, three chlorines. Now, the first step of this reaction is to turn this acetyl chloride into a good electrophile, turn it into something that's really good at nabbing electrons, so good that it can break the aromaticity of the benzene ring and essentially add itself to the benzene ring. This is actually the same mechanism we saw with electrophilic aromatic substitution. I always have trouble remembering the name, but I always imagine it's electrophilic substitution. Either way, but it's a very similar mechanism. And actually, what we're going to show in this video is called Friedel-Crafts acylation, because this right here is called an acyl group and we're essentially going to acylate the benzene ring. We're going to add this group right here to the benzene ring. So enough of what's going to happen. Let's actually see it happen. So the first thing to realize, this aluminum chloride, the aluminum in it is electron deficient. And at first, if you just look at the Periodic Table, you have these chlorines over here, pretty electronegative. Aluminum is in the same row, but chlorine's way more to the right, so it's more electronegative, so the chlorines are going to hog the electrons in this molecule. The chlorines are going to hog electrons, so the aluminum is going to have a partial positive charge. Chlorines will have slightly partial negative charge. On top of that, you see aluminum is a Group 3 element, one, two, three, so it has three valence electrons. You see that right here, one, two, three, nowhere close to the magic number of eight. Even when it covalently bonds with these chlorines, it can only pretend like it has six electrons. It can kind of pretend like it has these chlorine electrons over here, but that only gets it to six. So it would really like to have more electrons to get closer to that magic eight number. So what you can imagine is a situation where this chlorine on the acetyl chloride, it's already hogging this green electron from this carbon. It was already doing that. It's more electronegative, so this thing over here will actually be given to the aluminum. And so it will then have a bond with the chloride. So if that happened, what does our reaction look like? So if that happened, what does everything look like? Our aluminum, our aluminum chloride, or what was formerly aluminum chloride, now just gained an electron, and with it, it is now bonded to another chlorine. So it is now bonded to another chlorine, and since it gains an electron-- let me make it very clear. This is an L. My penmanship is deteriorating. That's an L. And since this aluminum gained an electron, it now has a negative charge. And normally, a negatively charged thing isn't that stable, but these guys are electronegative, so they might hog a lot of that negative charge. And on top of it, aluminum can now pretend like it has eight electrons. It has one, two, three, four, five, six, seven, eight. When you covalently bond to someone, you can kind of pretend like you have their electrons as well. So now you have this anion that was formed from the aluminum chloride, and now the acetyl chloride will look like this. Scroll down a little bit. Let me make it clear. We're in the next step of the reaction. What was formerly the acetyl chloride has now lost the chloride, so it's now really just an acyl group. So you have the carbonyl bonded to a CH3, a methyl, just like that. This guy lost his electrons, so now he has a positive charge. And this is actually not that stable. you're going to see it's actually highly reactive. It's a very good electrophile. It wants to steal other people's electrons. But it can exist for a short amount of time, especially because it is resident stabilized. You say, how is it resident stabilized? Well, this oxygen over here has two electron pairs that I didn't draw before. And let me draw the second one in a different color. It has two electron pairs like that. So you can imagine the situation. This carbon is already-- he has a positive charge, and his oxygen is more electronegative. It's already hogging his electrons. Maybe he wants to give back a little bit. Say, hey, this is positive. All of the electrons are hanging out here. They'd be attracted to the positive, and you could imagine one of these electrons being given back to the carbon. And if that happened then we have another residence form that looks like this. So this was our original molecule. That's our original, or what it looked like. We still have this double bond right over there or that pair of electrons. Now, this pair of electrons now forms another bond. This electron here is now on the oxygen end. This electron over here is now on the carbon end, and now they have a triple bond. And what happened? This positive carbon gained an electron, so it's now neutral. And the neutral oxygen lost an electron, so it is now positive. And you could imagine, this is not a very stable-- you wouldn't see this just floating around by itself, but it does stabilize this entire configuration. It stabilizes this molecule. So you can show that these are alternate residently stabilized structures right there. But as I said, these aren't super stable. This guy really, really, really wants to react. And now this is where benzene comes into the mix. And actually, let me draw a little dividing line here, just so we know that this was a separate stage of our reaction. So that was the first stage. Then we go over here and now benzene comes into the mix. The benzene was floating around. So we have our benzene floating around, just like that. And then I'm going to draw one of the hydrogens on one of the carbons. All of these carbons have hydrogens on them. I just won't draw them all. It just make things complicated. But this guy we said is a really good electrophile, and you have to be a really good electrophile to attract electrons from a benzene ring, to break it's aromaticity. But if it bumps into this guy in just the right way, at just the right angle, you could imagine that this electron on this carbon right here gets swiped by the acyl group. So then what do we have? So now I will go back in this direction. So you have what was a benzene ring. We can draw the double bonds here and here. And we, of course, have this hydrogen. But now this bond, which was a double bond there, is now bonded to the acyl group. So it has that blue electron that the acyl group nabbed. And let me draw the acyl group. And I'll flip it over so that we have the methyl on the right-hand side. So it's carbonyl bonded to a CH3. It was positive. It now gained an electron. It is now neutral. This carbon over here lost an electron, so it is now positive, so it is now positively charged. Now, we mentioned the aluminum chloride is a catalyst, so it won't just sit around there as the anion. It has to go back to being aluminum chloride, so let's bring the aluminum chloride back into the scene. So we have our aluminum chloride. Let me copy and paste it. So we have our aluminum chloride here, and so you can imagine that the benzene ring wants to go back to being aromatic, so this electron right here on the hydrogen might really want to go back to this carbon right over here, this carbocation. At the same time, if this anion now passes the hydrogen at just the right angle at the right time while this guy's attracted to this carbon, this chlorine can give this green electron to the hydrogen nucleus, which is really just a proton. And then the hydrogen's electron can be taken up by what was this carbocation. And then what do we have? Then we have a situation where our benzene ring is reformed. We have the aromaticity again. We have that double bond, that double bond, and now we have this double bond again, although now it's using the electron from the hydrogen. And now we've substituted this hydrogen with essentially this acyl group right over here. So we have a carbon double bonded to an oxygen bonded to a methyl group. And now the aluminum, or this anion, lost its electron, so it goes back to just being straight-up aluminum chloride. It goes back to just being straight-up neutral, electron-deficient aluminum chloride. And we're done. We've just acylized this benzene ring, and that's why this mechanism is called Friedel-Crafts acylation. And Friedel is actually a former president of MIT, and I did some reading on this. Apparently, he did not have a PhD, but because he discovered Friedel-Crafts acylation and this Friedel-Crafts alkylation as well, they said, hey, you know, this guy's a smart dude. Let's make him the president of MIT. But I just wanted to show you that this is a reaction for creating a ketone. So this ketone that we've created is acetophenone, which we've seen before, which we've learned is a common name for this molecule that we learned in the first ketone video. And I'll write it in purple. Acetophenone. And we're done!