Freshman Organic Chemistry I

CHEM 125a - Lecture 34 - Sharpless Oxidation Catalysts and the Conformation of Cycloalkanes

Chapter 1. Introduction for Professor Barry Sharpless [00:00:00]

Professor Michael McBride: So you remember this slide from several lectures ago. What are they trying to make? Remember what the point of this synthesis was, this particular reaction? The idea was to do a catalytic reaction with something that was chiral as the catalyst, to do an oxidation on sulfur, to make a single enantiomer of omeprazole. And remember the idea was that the product, after you've gone through these three steps that we showed you before, that the product then comes back and is the catalyst again, and you can do another cycle. So, in fact, there's this thing called a catalytic cycle. So what comes in is a peroxide and a sulfide, and what comes out is an alkoxide and the product that you want, the oxidized product; where an oxygen has been added to that unshared pair, that high HOMO. And so the idea of the chiral oxidizing agent was so you'd get a single enantiomer of the product. But do you remember what happened?

Student: It was racemic.

Professor Michael McBride: Right, it gave racemic material. So what did they do? Do you remember? Pardon me?

Student: [inaudible].

Professor Michael McBride: Yeah, ethyldiisopropylamine they added; these researchers added, in 2000. And they found out that that gave a 94% enantiomer excess, but for no obvious reason. Now it would be great if we knew what the reason for this was. And it turned out that, in fact, this kind of oxidative catalytic reaction, to be stereospecific, to give a single enantiomer, was the source of the Nobel Prize in 2001. And if we click on this thing here, we can see — should see — yeah, there we go. So this is the Nobel website, and you can see that K. Barry Sharpless here got the Nobel Prize in 2001 for his work on chirally catalyzed oxidation reactions; so exactly this reaction. So if anyone can tell us how ethyldiisopropylamine does its trick, who should it be? Barry Sharpless, right? [Laughs]

[Applause]

[Technical adjustments]

Professor Michael McBride: There, we're all set.

Professor Barry Sharpless: Thanks Mike.

Chapter 2. The Reactivity of Allyic Alcohols and Vanadium-Catalyzed Epoxidations [00:02:55]

Well I'd do anything for my friend Mike, because he saved me from publishing errors in JACS [Journal of the American Chemical Society] at least once. And he's saved a lot of us, who were willing to listen to him, because he's got a genius for seeing flaws in the beautiful stories that people tell, and then they realize that their beautiful stories are going to be too slick, usually; and Nature doesn't like things to be too slick. Well I'm going to just use — I think the next slide should be —

[Technical adjustments]

Professor Barry Sharpless: Well this is something which we all possibly manipulate in-utero. It's interesting, because many things in our body, like a part can be on the left or right, and a mirror image — I assume that means they're mirror-image related then. But this is a vein and two arteries that has a sheath on it, and it's the umbilical cord of a mammal. And it's got a nice screw axis; you know how we like to pull our hand down a screw axis. And then I have other reasons why I think I handled it. And then they take it away from you at birth; they cut it and they take it away. It becomes a narcissistically cathected object. And that means you're in trouble because you can't ever heal that wound. You want to get it back, but you can't heal it. You needed that mirror, that umbilical cord, back there when you were born. Anyway, that's why I went crazy trying to do asymmetric synthesis. And now, if you're a chemist, unless you're in deep space — I don't know if they have any carbon atoms out there, but you got to be in a very high vacuum to get them. But this is the way we make organic compounds, right? But we don't really do it this way because we have to live on a condensed phase of the earth. So these are two — but you put them together.

This is a point. Then you get a line. It's polyacetylene really. But let me — give me poetic license of just having it be a line. And I'll take one unit out, and now if I add things across that line, I get into the plane, A and B, or I can add them cis or trans. And essentially you can make anything this way, almost, if you could do it by just surgically adding things. Now we're in the plane. And now the job in the plane — and I've actually drawn a trans-disubstituted 2-butene. And it has a shape like this, right? It has a methyl group here, a methyl group here, and then hydrogens. So there's open space in these quadrants. And if you look at me coming at you this way, and if I turn around, you see it's going to be — well you'll see it's going to be a mirror image situation. But I'm flat. So I have no chirality, as a starting point. I'm in the plane of flat. So then if I add hydrogen to it, I get nothing interesting. But if I add an atom, like oxygen, to it, like here, then I've pyramidalized it. And so I can either add the oxygen from my front, and go like that — epoxide. Or I can add it from my back, and I go like that. And that's the other mirror-image compound. That's a really cool way. And it works. You can make epoxides that way.

And we have reagents today where you can recognize the front and the back of a flat object, that has pro-chirality, like this. And back in the early, in the mid — early seventies, we saw this industrial reaction that was nice. You could take vanadium and molybdenum, and you could epoxidize olefins with these things called — well I don't have the darn structure unfortunately — but it's t-butylhydroperoxide. It's the alcohol group here, with a peroxide bond. And I think you've already learned that peroxide bonds are weak bonds. They have a low lying σ* orbital, and you can attack them on the back side and transfer atoms to make transfer to carbons, like a HOMO of an olefin. So what we noticed — I noticed in one case that allylic alcohols were very much — seemed to be special. It was just noticed by English chemists in the thirties that they seemed to react better than just normal alcohols. And so, but nobody had taken any really — action on that.

Professor Michael McBride: We're not as fast as some people at getting to this stuff. So we should say that an allylic alcohol group is one that has an oxygen on the carbon next to the double bond.

Professor Barry Sharpless: Oh, oh yeah.

Professor Michael McBride: The OH is —

Professor Barry Sharpless: Oh right. I forget. The allylic alcohol — allylic means — allylic, you're alpha, one atom away from an olefin. So you could have anything there; nitrogen and oxygen. So you could have an OH, like in ethanol; that's an allylic alcohol. This is an allylic alcohol right — where is it? I don't — see, I'm sorry, I don't have the right structures. This is the allylic alcohol. It would be, if I go back one; I'll show you, not by going forward.

Professor Michael McBride: There's the allylic alcohol actually.

Professor Barry Sharpless: Yeah, but I got one — like here, this is a carbon, that's a carbon, that's a carbon. By putting an OH group there, that's allylic alcohol. If I have another carbon, it's homoallylic. But it has a handle now, and that handle allows it to grab hold of the vanadium. The vanadium has all these alcohol — these are esters. Imagine that's phosphorous. It is like phosphorous actually. But it's a yellowish liquid, and it exchanges its ligands like the wind, with alcohols. So they come in, they go out. It's a real dance. It's fast. You don't have to worry about that happening. Phosphates, we'd all be dead if that happened; we'd melt. But this is a transition metal. It goes to here. Now, and when I'm the metal, and I have oxygen here, and then I say have an arm. You can imagine my arm being a double bond, with π going like this. And here, on my belly, I'm going to have an atom — and this is an atom like — it should be red. Oxygen's red. But he knows why. I don't. Except for blood. So you come like this and you grab it, and it pops off me — I'm the activated thing on the metal — and it goes on to there. That's an epoxide; a three-membered ring. And if I attack it with a backhand, I would get the opposite result, in terms of attack. And that doesn't happen, because the geometry of the transition states and the projections of the center — you can't get as close to a center shot on this thing, with a backhand.

So that's my way of starting to explain. We're going to explain a little bit more. Mike's going to do some acting later, and you'll see it a little bit more clearly — the topology of this situation. But here it's cool because you get this template. It's like a mandrel. You grab hold of this oxygen atom, which is proto-oxygen. The metal has lots of empty orbitals. It's a d0 metal. It activates the thing this way, and then the olefin lone-pair, the HOMO, attacks the backside, and all this funny business goes on. But there's no mechanism. It's a no-mechanism transfer. It's one of these — like you draw a bunch of arrows and everything happens sort of in a smooth way; no ions involved, unlike SN2 chemistry. Okay, slow step is then this step. And then it comes off and more stuff goes on and it goes around. So that's the trick we played here. The first idea was take advantage of fixing the thing. So once you've fixed this flat system that you want to make go pop up this way, or that way, now you have a chance of doing it, because the thing you're transferring from can recognize this kind of situation, from that one, by — it's part of the same molecule temporarily. In fact, it's a covalent bond, right there, at this point. Too much, I always say too much. Okay, let's keep going.

Chapter 3. Research with Katsuki and the Discovery of Combining Titanium with Tartaric Acid [00:11:45]

Then came a great man from Japan. He's now a head professor at Kyoto, Oshima. He was a fantastic boiler room producer in titanium, osmium; you name it. I think that was where all the work was done for the Nobel Prize; it was done at MIT. But it happened at Stanford; just because I was there for a few years. But you know those early years in your career, when you do your work in the boiler room, are what produce you. Like your personality doesn't change after you're four years old — right? — or three years old; you're pretty much locked in, unfortunately, in some ways. This is vanadium system, and we put on this chiral machinery here, and it gave us 80% ee; that means enantiomeric excess was — it was 90% one isomer — enantiomer, the right hand say, and 10% the other. And this looks good. But it only works for that one substrate, and all the other types with it — you can have like methyls there, or a methyl there, and you get nothing. So you have to get everything customized. This ligand, you got to get the right ligand, the — and that's what you expect with Nature. We're taught by Emil Fisher — the lock and key. Nature is supposed to be very, very touchy about changing her preferences. And that was what we expected. Well I didn't expect that. I expected that this type of olefin and this type, with something, that all these would need a class of catalysts. There are about six shapes of olefins. So that's what my dream was, to get maybe six catalysts. Because we didn't have the paraphernalia that nature has all around.

Well for years we tried. Ten years we didn't get anything. It was always this game of like almost no reliability. And then when we came to titanium — we'd been working with vanadium and molybdenum. We came over to titanium, and we put in tartaric acid ester one day. This wasn't me, it was Katsuki. He should've gotten the Nobel Prize with me, because this means half of this — I picked the right metal, at that time, and the right system, but he's the one who picked the right ligand. And let's go on to the next thing. You'll see, this is tartaric acid. This is so-called natural. It comes from grapes. It's the cream of tartar. When you protonate it, it becomes tartaric acid. Over here is the unnatural. But there's a plant in Africa, Bauhinia reticulata, that has about — it's a huge area, around the Sahara. So there's maybe more of that so-called unnatural than there is of the natural. You know the story about tartaric acid. And it goes way back to Louis Pasteur, with the discovery of tetrahedral carbon; I mean then the precursor to it. Grapes, very sexy looking things. We always like — humans seem to like to drink them, and eat them, and look at them. This is the famous recipe.

So that day, Katsuki took a tartrate and put it together with — oh I think I forgot to say something here. Yeah, just quickly. Maybe I don't have it. No, it comes later. Okay. TBHP was the kind of oxygen we'd always been using. There's the titanium, which comes from titanium dioxide or titanium tetrachloride. Everything that's white in this world, all these — that's titanium dioxide. It's the only white pigment anybody ever uses. And so that's titanium dioxide. It's very insoluble. But you can make it into this soluble derivative, if you do the right chemistry. And this is the wine acid. We put in diisopropyl tartrate as an ester, and we kept getting 90%, 100%, 95 — every olefin we put in that was an allylic alcohol, that had this handle on it, gave Katsuki — about a week of this, and we were just about dying. I was looking at him; he was looking at me. We didn't believe it, you know. So we had to try to kill this. But it was right. And so it meant that here we had some new principle. We could actually just take one catalyst and get them all; almost get them all. But they had to be allylic alcohols.

So this is the famous recipe discovered I think on an October day in 1980 at Stanford. I'd already decided to go back to MIT. I might not have decided to do that, if this had happened sooner. We get depressed when our research isn't working; I mean, whether we don't show it or not. It's the only thing that matters really to a real killer research guy who wants to understand what makes nature tick. And whatever else — if that's going well, life is good. Otherwise nothing's good. So I went back to MIT, which I wouldn't have if I'd gotten this about three months earlier, probably. And here we are. There's Katsuki-san drinking tartaric aqueous — it's about 5% aqueous tartaric acid — in wine and champagne. And this was on the porch at the Mudd Building, which was where our lab was, on the top. And what I'm going to show here is why Katsuki was so important, in a nutshell. This is really a nice little substrate. It has this type of alcohol; it's trisubstituted. But you see the OH here? There's no OH over here. This is a more reactive olefin for the transfer of the oxygen, but it doesn't have the handle. So this one ends up winning by a factor of 200, in this system where the epoxidation requires the thing to bind and go in intramolecularly. And that's nice in its own right, but to get — we got racemate all the time, of course, because we have no face selection. And titanium goes about ten times faster with tartrate. When you add it, this does something to make — every other metal, years before Oshima had taken vanadium and molybdenum, which were the best for isolated double bonds. But you put tartrate in, it kills them. So my instinct would not have been to put tartrate in again.

This is the kind of thing — everything in science, the rules change all the time. You wake up in the morning — you should iterate your favorite desire every morning. It's going to look different to you in the shower. And you can never — the atomic bomb was made by theoretical geniuses. But the guys who did it, actually they had aluminum foil on the source, and they kept it on there for about a year or two. Some guy said, "No, I don't think we need this. What about if we take it off?" That was the answer. I mean, these guys are like us. They just try this, try that, in the lab. Right? That's what gets really things done in this world. The theories are important. Nobody would even think about trying it without the theory. But — okay, "A man in California just won a Nobel Prize for mixing paint and wine!" That's what they said, in the LA Times. It's sort of true. Titanium dioxide and tartaric acid. And there's the — I like my students to know where Mother Earth is. And I like things to be cheap. I don't like to be far away from a river that's strong, or a power source; or in this case, this is tartaric acid. It comes in 100-pound bags. This is Spanish tartar, and this is Victor Martin, who's one of the heroes of this chemistry, from the Canary Islands. This comes to Pfizer. They make things out of it. But it comes in ships in 100-pound bags. And we still have that in our stockroom somewhere.

There's my daughter celebrating the synthesis of the unnatural sugars [laughter] by this method. And oh I got the wrong one, didn't I? I threw out the wrong — yeah, anyway, you're going to have to interpret. This is the mirror-image of the picture that was taken, I guess; because you see the name over here is backwards. But she's looking at the l-sugars. These are the ones we don't make. There's eight of those, the hexoses. And these are the ones Emil Fisher made, in his famous work; about the best chemist who ever lived, except for van't Hoff, who was even better, because he did organic stereochemistry and physical chemistry, and I think he was the greatest chemist who ever lived; van't Hoff was. But anyway, Fisher was better than anybody who lived in this century, in terms of understanding and arguments and rationale of synthesis. And he made a lot of these sugars over here.

But this is Tito Simboli. She's a photographer, my wife's friend. This is taken in 1983. It was on the cover of Chemistry in Britain. And Hannah was seven. And anyway, Tito was a friend. So she took the picture. And this book though, is a huge book, which was gotten by Nancy Schrock, who's in charge of the MIT Archives and is a book restorer. And Nancy — so here we got Nancy Schrock, we got Tito Simboli, whose husband is Dan McFadden; won the Nobel Prize in economics ten years later, and Dick won the Nobel Prize a couple of years ago, and I won it in 2001. So the picture is kind of connected. And there's Lewis Carroll, of course, looking-glass milk. And Emil Fisher. It's a heavy duty picture; for me anyway. [Laughs]

Chapter 4. The Mechanism for Asymmetric Epoxidation of Olefins and the Story of Nexium [00:21:11]

There's Nexium, and maybe we're getting — I went too far into the other stuff. But here's the atom, the red atom, that's going to end up being transferred onto the sulfur. And well Mike showed — told you that the diisopropylamine — what would its role be? I hope I can go to the board now and show you a few things. But this catalyst is exactly the recipe. And it makes billions of dollars a year, this product. Right? But the patent is — the Stanford patent is no longer valid. It's run out; even though it hadn't run out when they started. But they didn't believe that it might. This was an infringing of that patent actually in the beginning. But nobody in universities sues companies. They can't afford to. Stanford couldn't afford to.

Professor Michael McBride: Do you want to go to the board now?

Professor Barry Sharpless: Should I? Yeah.

[Technical adjustments]

Professor Barry Sharpless: Do you remember the vanadium picture? I'll make it titanium now. But there's this alcohol group and —

[Technical adjustments]

Professor Barry Sharpless: And on this we have this peroxide group that's bound, and it's bound datively there. It's got this tertiary-butyl group here. And then the alcohol is bound here, and well it goes sort of down. I'll just try to show some of the stereochemistry. And then underneath it's coming — yeah, it's coming folded like this, underneath. I can't draw very well, but this is supposed to be coming with its lone pairs, it's π bonds, on the backside of this bond. So we're going to do an attack on there, and we'll break this bond, we'll bring the lone pair here in, and we get the epoxide. Now, but the thing about this is patent lawyers can be very creative — like every field has creative people, and Yale is connected to this fellow, Bert Rowland. And he was in Palo Alto at the time, at a company, and he had just gotten famous, or almost infamous, because that patent was really aggressive; this Boyer-Cohen Patent. He wrote the Boyer-Cohen Patent. He's related to Mike, because he got his Ph.D. with Bartlett at Harvard, like Mike did. So that's a long time ago. That's a great-something relationship. And then comes Wiberg, who's here at Yale. He was at Seattle then, and he took a post-doc. He decided he liked chemistry in principle. He was smart; a physical-organic chemist. But he didn't like the lab; he wasn't any good in the lab and he didn't like to be there. So he became a patent attorney, and did very well. And he writes great patents. And so he read our paper, our first paper. We gave him the little communication. And then he got Katsuki and me over to his office, and he interviewed us for about ten minutes. The next day we got a patent. He never changed a word of it.

And I'll tell you why it's creative, I think. We're getting ready for the Nexium. But I could tell you about that. If we don't make it, it won't matter that much. But here's the main — so this is going to be the oxygen; and they should be red but we have pink here. So does this one. But that one doesn't transfer. This is the one that transfers. So what Rowland said is you got a metal, which can grab things. Now the things it grabs, it could be X, it could be sulfur, or an amine group that it could grab. And then he just drew — I think he drew a carbon with no — it could be any carbon. It didn't have to have to have just CH2. And then he drew this, G. And then he drew — yeah, he did draw an oxygen atom here, I guess, but it was activated. And so this oxygen, he didn't have a generality there. That was specific. That was the only thing that was specific, these two things. So this could be any kind of lone pair. An olefin is a lone pair. Sulfur has a lone pair. Phosphorous has a lone pair. Nitrogen has a lone pair. And sure enough, when I read that, I said, "My God, this gives us ideas." So we started doing amino alcohols. All these things work. You know? So this is from Bert Rowland, the patent attorney. He was a co-inventor of those other reactions. We never wrote any more patents. It's a good thing. Don't waste your time on patents. They'll just not use them until they expire. And we never — we made enough to go on a vacation once a year, in a car, not too far away. [Laughter] Katsuki and I, we got — people think we got rich. We got the most — one year we got $20,000 each, or something. That was a really sharp maximum, from the bag-a-bug gypsy moth traps – Disparlure.

But okay, to show the transfer, I think maybe it would be nice for us, Mike and I, to show what features have to be over here in these ligands. You see, the tartrate has this feature. It grabs the titanium, like out front. Let's kind of put it out front here. So the titanium is out here. And down in the back here you have this ester group, and up in the back, over here, you have an ester group. And I can't remember if that's what we planned to do. But out here, on the front, it's coming — the alcohol — this G-group, and there's an oxygen that's hot over here, that can be transferred. So Mike's going to put that — I'm going to put that on and —

Professor Michael McBride: I'm the olefin.

Professor Barry Sharpless: Yeah, you're the — is that okay if you're the ole —

Professor Michael McBride: It's okay.

Professor Barry Sharpless: Maybe it's better if I'm the olefin.

[Laughter]

Professor Barry Sharpless: No, okay. I'll be the titanium.

Professor Michael McBride: Yeah, you're the titanium.

Professor Barry Sharpless: So see, I've got this oxygen that wants to go off, because it's a weak bond. And I've got to get it over to him. He's the allylic alcohol. But we're going to pretend he doesn't have to bind to me. We're just going to do this straight-out attack here. And yeah, it's like that. Right? So I'm blocked here, but I'm open in these quadrants.

Professor Michael McBride: So you're natural tartaric acid.

Professor Barry Sharpless: Is that right? Boy you're quick. I don't — okay, so I'm natural?

[Laughter]

Professor Barry Sharpless: Okay, well wait a minute. And also, but I'm not like this, I'm like that. You see, I'm in a three-dimensional world, because I'm tartaric acid. So I'm sticking out here. So I'm a chiral object now. My mirror-image won't superimpose. But he's, yeah he's that way. He's a trans-olefin. [Laughter] And he looks like an Egyptian, like — anyway.

Professor Michael McBride: I have a double bond here, right?

Professor Barry Sharpless: Yeah.

Professor Michael McBride: The question is whether you're going to come on here or on my back.

Professor Barry Sharpless: Actually the double — we don't have this quite right. Can you get an arm down here? Because the double bonds is here.

Professor Michael McBride: No, no, the double bond is going this way.

Professor Barry Sharpless: Okay. You've got it going that way.

Professor Michael McBride: So here, right now.

Professor Barry Sharpless: Okay, so yeah double bond's going that way. Okay.

Professor Michael McBride: Thank you. So you —

Professor Barry Sharpless: Yeah, that's good.

Professor Michael McBride: Get your tartrate on.

Professor Barry Sharpless: Okay, because I could go like — I may not — it's like this, I'm sorry. Yeah, I got to go like that. I could go like that. That's the mirror-image. Okay, now I'm coming towards him.

Professor Michael McBride: But I want to go on my back.

Professor Barry Sharpless: Okay, you want it on your back? Oh my gosh. Oh I can't do it. See? But that's the other — you need the other mirror-image for that one. So I come this way, and I can — it says "Think Safety". Oh. So now you have to — yeah, if he was a good yoga person, he could pyramidalize it a little better. [Laughter] Yet now when you put him that way, he's one enantiomer, and if you put him on the other way, that's the other enantiomer. Simple. It's real simple, isn't it? Thanks Mike.

Yeah, we just organized that ten seconds before we came over, as you can see. We didn't have any red ping pong balls. It's good with Velcro. You can do it with Velcro. We did it once with Velcro red balls. Okay now.

Professor Michael McBride: Back to the screen.

Professor Barry Sharpless: Yeah, I think we're — oh no, I'm not back to the screen. Let's leave it alone. This is going to be a little —

[Technical adjustments]

Professor Barry Sharpless: Now if you look at the structure of the omeprazole. I won't put all the bells and whistles on it. It's an imidazole structure, benzimidazole, which has a benzene ring, two nitrogens. It's a five-membered aromatic heterocycle. And so you have a double bond and one H. And I can't remember what's on the sulfur. Is it an aromatic ring, or is it a benzyl? It's an aromatic benzyl thing, I think. And you might look at this and say, "This doesn't look anything like Bert Rowland's thing there." Because where are the anchor points that we need, the group that binds the metal? We have the lone pairs all right; you know, two lone pairs, rabbit ears. And there, if you put oxygen on one side, it's the same story as before. This isn't flat to start with. But now we have these pairs. If you put something here, you get one enantiomer. You put something there, you get the other. So the idea was to try to get — and Kagan had done this. He's a fellow that could've won the Nobel Prize, I thought should of. But he's in France. And he had done a lot of great work in asymmetric chemistry. And he took the titanium tartrate catalyst of Katsuki, and he found if you put a little water in it, and you did the right things, and wished — it wasn't as general, but he got very good asymmetric addition.

So titanium was already known to do this. But the feature that's important here is the pKa of this benzimidazole is probably — I looked it up; I think it's below that of an alcohol. So that means it can go on the titanium, reversibly. So you see the titanium has alcohols on it, or ligands like this, or — and this can exchange. So we can have that come off, and this N go on, for this heterocycle, and the benzene ring's down here. And now we have something like that, a covalent bond. And we've got the sulfur here, and we also have the alcohol. So there's an equilibrium where you could invoke that. But this is not as easy to probably do as an alcohol. It's more encumbered, and it's probably — kinetics are slower for this, and off of oxygen, for the hydrogen transfers that are needed. So I hypothesize here that what the diisopropyl — what the Hünig's base, we call it diisopropylethylamine is doing. This is such a hindered amine that it can't react with anything, but it's good for getting protons. So it can't itself get at the oxygen; which you would worry about. It would make an N-oxide. And it can't bind to the titanium anyway. Titanium hates nitrogen. If you have an oxygen around, it'll spit it out so fast it'll make your head spin. I mean, this thing is amazing hating nitrogen — titanium, and silicon too, but especially titanium. So it's never going to get involved with titanium.

So what's it going to do? Well remember when you try to get chemistry going in a system like this, this amine has enough pKa power to pull this proton. And usually you could write a concerted mechanism. Like let me erase some of this and have some room here. Well I shouldn't erase that. I'll go over here. You got this benzimidazole, with a sulfur here, and I'm going to put the hydrogen up here. Whenever you engage compounds in reaction, acid-base reactions, you look for the basic sites and the acidic sites. There's the acidic site. You might think that's where the titanium goes. But usually the mechanism for these reactions involves a simultaneous loss of this proton, and attack here. This is the place you can attack. So you'll be attacking the titanium here, and you're going to get some help from — you're going to get a lot of help from this amine, because you can put it someplace; just take it right off. But you'll have some of that ammonium salt around too, a little bit, transiently. And this will help you, because you need to move things around. That's why this reaction was not good without that amine. You need to scramble the system and keep it rolling, so that the things get on and off.

Catalysis often has this — like there's six, maybe. There's always a loop, and if you look at any catalytic cycle, there'll be usually one step that's really bad news, until you get it fixed. And that step is where it controls the rate. If you get steady state, the slowest step — this is a real democracy, catalysis. There's no step that's more important than any other. The catalyst goes from each step. The titanium is moving through those cycles. If there's a slow step, there's 99.9% of the titaniums that you need, stuck before this one mountain, that goes way up like Mount Everest. You got to drop that one down. If you get that one down, the rate goes way up. And if you get them all the same height, you're really rolling. And that's a sacred rule. The turnover-limiting step is everything, in catalysis. And a lot of times the turnover-ruling step is so high, they don't even know there is catalysis. If you can break that, you'll find a world of catalysis that didn't exist before. So that's what I think is exciting about catalysis; it's alive. As long as you have some energy to dissipate you're — it's life itself.

Then I'll finish this quickly off, without any details. You can see what I'm going to do here. It's just that that hot oxygen atom, that's activated to transfer. This is the handle that would be the allylic alcohol. And so it's a dead ringer for the Katsuki — for the Bert Rowland Patent, right? That's what I think. And the amine, yeah, the amine makes sense here too. And I guess that's —

Professor Michael McBride: So did that patent apply to this process?

Professor Barry Sharpless: Well Stanford made a little — they brought over ten people, from research, and they brought over three lawyers from AstraZeneca; because Stanford thought maybe they should get something out of this. And I told the story like this, and they didn't — they pretended it didn't make any sense at all. You know? And it was pretty unsatisfying. Bert Rowland was there. I don't know if he's still alive. He had cancer then. But so we didn't — but they'd already made six billion dollars; because the catalyst patent was the only extant patent. But it did have this feature in it. And so they were right in the — I think they're right in the face of the patent. But they didn't like that. And we didn't — Stanford wasn't going to pursue it.

Another thing we — I'd like to mention — is they invited me over. This is before this meeting about the patent. I guess I started to notice the old titanium chemistry, and Katsuki and I talked, and we agreed that maybe we should raise our hand. But before that, what happened was I got invited to go to a hotel in Munich, the fanciest hotel, and they put me in this big room upstairs. They wanted me to say something about right- and left-handed medicines. And I had a little 15-minute spot — 1000 or 1200 gastroenterologists, from Germany alone. And they were doing a story about how important drugs have to be optically pure. And that story doesn't wash too well here. And I didn't realize it, when it was going. But I went and I did my part. And it is true that many drugs are toxic in one form and good in the other. And I was able to make those points. But what happens here is — oh, and also they gave me this suite on the top, that had like a spiral staircase to the bed, at the very top; and Madonna had stayed there the week before. [Laughter] That's okay for her. She doesn't have a prostate problem. [Laughter] I hated that place. I had to go — they should've had a fire pole so I could get down to the bathroom.

Anyway. What happens is, as Mike told you, you just look at the facts in that system. And I noticed them just the day I was sitting there in the audience, and the press was interviewing us. And I whispered to the guy next to me, and he looked at me and frowned. It was about the weights. All right? I mean, the racemate 20 mgs, and optically pure. And I think the optically pure is substantially more reactive, in this case, because it gets in better. It's, of course, just a pro-drug, for the thing. But the thing that was missing from the experiments is the forty milligrams of racemate, which would give them equivalence against the one enantiomer of the twenty milligrams of pure. And that was missing, and they doubled the — and it was just a sore thumb there for me. But I didn't say anything; nobody asked me, fortunately. Because if they had, I would've said something about it, I imagine. It's the way we are. Right, Mike? Mike doesn't — he smiles, but he gets the hard evidence. In the middle of your talk sometimes he does this. But for me it's the best thing you can get.

Professor Michael McBride: [Laughs]

Professor Barry Sharpless: If you can get killed by a friend, it's the best way to get killed.

[Laughter]

Professor Barry Sharpless: Anyway, I'm finished. That's it.

Professor Michael McBride: Great.

[Applause]

Professor Michael McBride: Great. Thanks again.

Professor Barry Sharpless: Yeah, you're welcome.

Professor Michael McBride: We maybe have time for one question or so, before we get back to our normal business. Anybody got a question for Professor Sharpless? No. Well thanks again.

Professor Barry Sharpless: You're welcome.

Professor Michael McBride: Have a good trip back.

[Informal discussion between Professor McBride and Professor Sharpless]

Professor Barry Sharpless: Okay thanks. So long. Good luck.

[Applause]

Chapter 5. The Conformation of Rings: Carvone and Cyclohexane [00:41:06]

Professor Michael McBride: Okay, back to our routine now.

[Informal discussion between Professor McBride and Professor Sharpless]

Professor Michael McBride: Oh, he forgot to show this picture. There he is with the royalty of Sweden, and his wife, and the other Nobel Laureates in Chemistry. Okay, and he was going to talk about carvone too. He wanted you to smell it. I told him you'd smelled it. But he had an idea for a novel based on carvone, which he didn't have time to get to. Okay, so we were talking about the conformation of rings. And we talked last time a lot about cyclohexane, and how it distorted — remember, that it's not really ideal, with the axial bonds parallel to the axis. They spread out a little bit. The ring flattens a little bit to minimize the energy, as calculated by molecular mechanics. Now how about in a four-membered ring, instead of a six-membered ring? Well if we look at the, what molecular mechanics says about the source of strain in this system, you can see that the big contributors are bend — that's what Baeyer had already talked about, about the ring strain; having 90° angles is not going to be good; so that's costing 13.5 kcal/mol — and torsion. Because why is there such high torsional energy; 15 kcal/mol, almost? What's the source of that? Anybody see? As you go around the ring, everything is exactly eclipsed. Everybody see that? Every carbon-carbon bond is eclipsed. Right? So that, in molecular mechanics, sums up to almost 15 kcal/mol. Now, if you were — how would you try to minimize the energy? Is this probably the lowest energy form, or can you think of changing it, so that it'll be lower? Any ideas? Kevin?

Student: Make it so the bond, in cyclobutane, aren't exactly parallel to the new bond.

Professor Michael McBride: Yeah, if you make it not exactly eclipsed, if you make it a little bit towards scatter- towards staggered, by twisting about the bonds, that'll lower that 15 kcal/mol. But, can you see what happens, if you begin to twist around the bonds? It sharpens the angles even more, from 90° to smaller than 90°. So here's what happens when you do it. The bending energy goes up, by 2.5 kilocalories per mole. Right? But the torsional energy goes down, by 3.5 kilocalories per mole — right? — if you don't bend it too far. Okay, and the other things stay pretty much the same. So this is a competition between torsion energy and bending energy, and the minimum is reached with a little bit of bending, in order to relieve some of that torsion.

Now how about if you had a five-membered ring? Now you notice you have — notice the bending energy isn't bad, because as Baeyer pointed out, a pentagon, a regular pentagon, has just the right angles — 109.5° about. Right? So there's hardly any bending. But there's still a lot of torsion, from the eclipsing of the carbons. So what you do, if you run the molecular mechanics program, is to cause bending energy to occur; that is, sharpen the angles a little bit, right? But you get rid of a lot of torsional energy by doing that. And this particular conformation — which you can see how it's bent down; the black carbon, eight there — is sometimes referred to as the envelope conformation of cyclopentane. See how it looks like the fold of the envelope bent down? Okay, so there again is a competition between bending and torsion.

Now, like plastic models, molecular mechanics is satisfying because it says not only what the structure should be, but why. What is it that makes the energy the way it is? Okay, so you remember last time we passed around the cyclohexanes. Remember how they clicked, to go from one form to another? So what's happening? The question is, what's the source of the barrier to the cyclohexane ring flip? So here's the chair cyclohexane, and when the ring starts to flip — remember first you go toward a boat, by bending one of those carbons up, like that. Now what do you think the source of the energy is, that makes this hard to do? Why do you go to a maximum energy, as you do this, and then it clicks and it goes down again? Sherwin?

Student: Torsion.

Professor Michael McBride: Torsion. Now, that's a good point. There are, in fact, two bonds, two butanes, that become eclipsed. Right? Both of the ones on the right. This one is eclipsed, but also this one is eclipsed; those four. Right? And indeed, that's worth about 7 kcal/mol of going uphill. But there's an interesting point. You did this with the models — right? — and you felt it click. How do those models, those plastic models, know from torsional energy? Do you see what I mean? If you take ethane in those models, and spin it, it spins completely freely. There's nothing in the model. You could make it — you could make the thing that goes into the tube, and the tube, a little bit triangular, so that as you tried to twist, it would go up in energy and then down again. Right? Do you see how you could do that in a model? But that's not the way those are made. They're cylinders, they rotate freely. So why did the model click? Pardon me?

Student: Bending.

Professor Michael McBride: Bending? Of what?

Student: The bonds.

Professor Michael McBride: Right, exactly. Look at this. Why does the plastic model click? Because when you're halfway across, it's becoming — if it were flat, if the whole ring were flat, what would the carbon-carbon-carbon angles be, if the ring were flat? Sam?

Student: I don't know.

Professor Michael McBride: In a regular hexagon, what's the angle? Dana? You got to be careful holding your hand up, even if it's just to scratch your face. What's the angle in a regular hexagon?

Students: 120.

Professor Michael McBride: 120°. It wants to be 109. So it has to stretch out, if it becomes flat. And even when it becomes only partly flat — right? — this part of it is flat; this one's out of the plane. But that means in order to have this one in there, this bond gets compressed. Right? Because if these angles want to be 109, instead of 120, everything — these are going to move; this one and this one will move closer together. So these are — in this form, when it's half planar, these angles are bent out, and that angle is bent in. So it's the angle bending which in fact does it, does stress those plastic models, which causes it to click. So it's interesting to be able to look at models, or at these calculations of molecules as springs, and see why various things occur. So are molecular mechanics programs useful? Yes, definitely they're useful. Are they true? No. Okay, so we'll go on from there next time.

[end of transcript]