 Okay, we're going to finish up, we just got one last thing to talk about for lecture 13 on anions. So we were talking about a bunch of pKa stuff for anions. And the last thing I'm going to talk about is called the alpha effect, Steve Benner once said that whenever somebody calls something an effect, it means they don't understand what's going on. Otherwise, you'd give it a more descriptive or apt name. I sort of feel like maybe we have a better idea of what's going on with the alpha effect. And it relates to this observation that if you look at the stability of a typical imine where you just have an Nr group where a ketone carbonyl would normally be, you find that if there's a heteroatom attached to that nitrogen, that these tend to be significantly more stable than a regular imine, dramatically more stable. To the point where you can even form them in water and the equilibrium favors the imine over the hydrolyzed ketone. So you already know that carbonyls are stable. Why is it that these should be so stable towards hydrolysis and formation of a ketone? So this is what we refer to as the alpha effect, the fact that there's an extra heteroatom here that tends to make this more stable. That's a thermodynamic interpretation of the alpha effect. It has to do with stability of things. There's also a kinetic side to the alpha effect. What you find in solution is not only do you form more stable adducts, and let me say, let me go ahead and write this down so we're clear, that this is more stable than just a simple either of these two derivatives are more stable than a simple imine. So we call this a hydrazone, and these have now become common fixtures in the field of chemical biology because you can form them in water, and we call this other an oxygen system, an oxyme, and dramatically more stable than an imine. So the kinetic interpretation is that when you have the precursors to these, what you find is these are more reactive in nucleophilic processes, and that's a kinetic effect. You get faster rates of reaction for an amine attacking some alcohol halide in an SN2, you get faster rates of reaction when there's a heteroatom next door. So there's a thermodynamic stability aspect, and then there's a rates of chemical reaction aspect. So you find this issue that it's faster than just a simple amine when there's a heteroatom next door, for example, in an SN2 reaction. Okay, so let's talk about these two aspects of the alpha effect, and I want to start off by talking about what I feel like is the more obvious aspect, and that's the thermodynamic aspect. And so let's come back to the idea of why is it that something like this should react with water? Why should water want to attack that imine the way water attacks a carbonyl? Well, one way to interpret that reactivity is that we can think of any sort of an imine or a carbonyl group kind of like this, that maybe there's some partial positive charge for a ketone. We said there's about a fourth of a positive charge at a ketone carbonyl. So to the extent that charge matters in nucleophilic attack by water at this center to make it hydrolyze, anything that can stabilize that electron deficiency ought to help that system. And there's a resonance structure we can draw for this. And let's draw that resonance structure. If I take that lone pair on nitrogen or the lone pair on oxygen, if it's an oxyne, and I push this over, then what I see is, well, maybe that carbon shouldn't be as electrophilic as I might otherwise think it to be. And that's exactly the case. So that you can see how that, just by drawing a simple resonance structure, you can see how that heteroatom next door helps to make this carbon atom less susceptible to attack by water or hydrolysis. And so you can think of this as a way of explaining the preferential stability of hydrazones and oxynes. Now the tougher thing to explain is the kinetic effect. If there's no carbon atom here, how do you explain why the lone pair on this nitrogen and hydrazine or hydroxylamine, why is it kinetically more reactive? And so let's come back and interpret that. I'll give you a simple kinetic scheme. If you take hydrogen peroxide, let me draw this out. There's hydrogen peroxide. We add a base to that, methoxide anion in the presence of methyltosylate. So what's the product of that reaction? We've got methoxide, that's a good nucleophile. We've got methyltosylate, that's a good electrophile for SN2 reactions. There's no competing E2 elimination reactions here. What you find is you don't get any of this. You don't get any of this methoxide doing SN2 reactions on the methyltosylate. What you see is that the hydroperoxide anion is what attacks the methyltosylate. So what's happening here is there's a very rapid equilibrium. We know that proton transfers are fast to generate the hydroperoxy anion. So there's an equilibrium in here. There's more metho, there's methoxide. There's hydroperoxide and it turns out that peroxides are more acidic than regular alcohols and maybe you could have guessed that just by the fact that there's an extra electronegative atom next door here that can help with that negative charge. And what you don't know is the numbers. It's about pK of 12, so pK of 11.6. So somewhere on the order of about 10,000 times more acidic than a regular alcohol. So the equilibrium just based on pKa's ought to be about 10,000 times more of this than there is of this. There's about 10,000 times more peroxide anion at equilibrium than methoxide. But what you could not have guessed is not only is there more of this is that it's 10 times more nucleophilic in solution. In most cases when you look at simple pKa's if there's 10,000 times more of something it turns out to be 10,000 times less nucleophilic. And then you can't guess what the products are going to be. This is a rare case where not only is there 10,000 times more of this at equilibrium but it's also more nucleophilic. And so the effect you see is that it's the hydroperoxy anion that's more abundant and more nucleophilic that is reacting as the nucleophile in this reaction. Okay, and that's the alpha effect. That's a kinetic effect, this very unexpected fact that the hydroperoxide anion in spite of the fact that that's more favored, more stable also turns out to be more reactive. And people argued about this effect. Why should that be? Like why should that hydroxy group there affect this, affect the reactivity of that O minus? And I'll give you a simple fact. It turns out that that is not poorly solvated. Solvation decreases reactivity. When you solvate an anion with solvent you make it less reactive and solvation of hydroperoxy anion is less effective than solvation of hydroxide anion which is the corresponding nucleophile without an alpha effect. By 21.5 k-cals per mole, that's a huge energy difference. So if you convert that into factors of 10 you get a sense that, wow, hydroxide is really being solvated. You can imagine any respectable solvent forming hydrogen bonds or if there's sodium ions around crowding around that oxygen minus whereas it's just not as effective for hydroperoxide, the hydroperoxide will be more naked and more nucleophilic. Even though it's easy to make a hydroperoxide anion by deprotonating it turns out that you're basically distracting that O minus less with solvation. So it's now generally accepted that the kinetic alpha effect is a solvation issue. It has to do with the poor solvation of alpha effect nucleophiles. Hydroxyl amine is less effectively solvated. Hydrazine is less effectively solvated. These are better nucleophiles than their simple amine counterparts. And so you can take advantage of that in synthesis or chemical biology, the fact that those have extraordinary nucleophilicity. Okay, so that's it for lecture 13 on anions and anions as nucleophiles. And so now we're going to switch over to something closely related, basically acknowledging the fact that I've been scamming you from the beginning. And we're now going to talk about things that are bonded to sodium and lithium and potassium to that first column of the periodic table. If I sketched out a little periodic table here like this, what happens when things are bonded to lithium and sodium and potassium and cesium? Not rubidium, but radioactive. But cesium, things like that. So to get us set up to think about butylithium, sodium alkoxides, we need to first start off by reminding ourselves that we made this claim at the beginning of the course that we were going to explain rates of chemical reaction based on three different factors, three different terms in some sort of hypothetical equation. We said that charge is going to matter. Some things will react faster because minus likes plus or repel because plus hates plus. So like charges repel, opposite charges attract. We said the sterics is kind of obvious. We don't need to discuss that anymore, I think. And then last, we said we were going to use orbital overlap and frontier molecular orbitals to explain things. So when you think about the reaction, for example, of some alkoxide doing an SN2 reaction on methyl tosylate, the reason why this is fast is because there's a large amount of partial positive charge there on that methyl group. That's the effect of the tosylate. And there is an orbital contribution to this, but columnic attraction plays a big role in why that's a fast reaction. If we come over to other reactions, we're not going to be so adept at using that charge explanation. If you look at the reaction, and we'll talk about these more later, of an enamine, we can't use that charge interpretation to explain why SN2 reactions are fast. Here we can't use that charge interpretation to explain why this carbon atom at the end of the pi bond is so good at attacking methyl iodide to make a new C-C bond. There's very little negative charge at the end of this carbon atom in this pi system. This is an issue of orbital overlap, and we'll talk about enamines and enolates and enols more later. So we need to keep these two things separate, this issue of charge and the columnic equation versus orbital overlap, filled orbitals that are nucleophilic, unfilled orbitals that are electrophilic and one to electrons. So I want to come down and talk about sigma bonds as nucleophiles. So here I drew this as a lone pair attacking, but now I want to switch back to just a simple carbon-carbon bond. I think you know that ethane is not a good nucleophile for attacking things, that that's not very good. And so what would you expect if I replaced this carbon atom with a fluorine atom? What would you expect, maybe I want to make this attack things. I'll write E plus for electrophile. So what do you think would happen to the nucleophilicity of this bond if I replaced carbon with a more electronegative atom? Yeah, I would expect that to decrease. I would expect if I replaced carbon with a more electronegative atom, I would expect that to be lower in energy, the sigma bond, the canonical front two orbital for that sigma bond. And so I would expect that to drop. But what happens if I reduce the electronegativity of this other atom? Let me march across to the other end of the periodic table and replace carbon with lithium, another second row atom. Clear on the other side of the periodic table, my prediction is and ought to be that I should expect the nucleophilicity to increase. I should expect that this orbital, this canonical orbital for a carbon-lithium bond ought to be higher in energy, way higher in energy than any other type of regular sigma bond that we've talked about, than any CH bond, any CC bond. So what we're going to see when we talk about alkali organometallics, alkyl-lithiums predominantly, but alkyl-sodium, is that not only do we see an increase in the energies of these canonical sigma bond orbitals, we're also going to see an increase in negative charge on carbon. And there's two factors now at play that make this a vastly superior nucleophile and base to anything else we've talked about. There's that electrostatic calomic charge thing and then there's this MO thing and they're both coming into play. So let's talk about alkali organometallics. So if we make bonds between carbon and lithium or carbon and sodium and carbon and potassium, what kind of predictions should we have? So molecular orbital calculations, ab initial calculations are very powerful. They allow you to predict all kinds of things. Ultimately, at the end of this class, I hope that you can make all kinds of predictions just by knowing the periodic table. That's what we really want. We don't want you to rely on powerful calculations methods to think about chemistry. So I want to start off by looking at the electronegativity of a series of atoms. And I'm not sure why I organize this. I'm going to have three rows here. So I'm going to have one row where I talk about the electronegativity and then I'll have another row where I talk about the covalent character in this bond. And then I'll have another row here where I talk about ionic character in the bond. And really when I talk about ionic character for a bond like that, I'm really thinking about basicity. Proton transfer reactions are really driven by charge. So it's kind of telling me about basicity. Okay, so let me go ahead and give a list of different reagents here as I march across this, my little table here. And I'm going to start off with the middle row. So I'll start off by talking about a proton. And then I'll write magnesium. And then I'll write lithium. And then sodium. And potassium. And rubidium. And cesium. And I don't know why. It seems like it would make more sense if I oriented this vertically because we're kind of marching down this first row of the alkali earth metals from lithium to sodium to potassium to rubidium. Okay, so let's go ahead and write out the electronegativity. The electronegativity of a proton is 2.2. Carbon is 2.5. A proton is slightly more electropositive than carbon. If we compare that to magnesium, there's a huge jump in electropositive character. So when we think about a carbon magnesium bond, so in other words, let me see if I can draw out some sort of scenario for you to help us think about, how do I think about a bond to magnesium or to sodium or to potassium? There's a resonance structure we can draw. If carbon with an electronegativity of 2.5 really is more electronegative than a metal. Then when I draw a resonance structure, I shouldn't give the electrons to the metal. If I draw a resonance structure for this, I should give the electrons to the carbon. And so often as an undergraduate, you use this kind of thinking to help you understand the reactivity of organometallics like organolithiums. Okay, so let's go ahead and draw out some of these organometallics. If you look at the electronegativity, sorry, the covalent character. So if I have a bond between carbon and it's, I have a symbol there, M versus H, I would think of that to be best represented by this, a covalent bond, not ionic. So if I look at the covalent character when this is H, that should have the most covalent character of any of these. But if I replace the M over there with magnesium, that should have less covalent character and more ionic character. And as I drop down the first row from lithium to sodium to potassium, rubidium, cesium, what I expect to see is that things become more and more ionic. This ridiculous looking picture over here starts to be a better and better representation as I drop down from lithium to sodium to potassium, et cetera. So I expect that things will be less covalent. And conversely, I expect things to become more ionic. So conversely, when I draw it some, and it doesn't matter whether it's a bromide or chloride on the green yard reagent or an iodide, when I start to draw these reagents out, that what I should expect is these will start to look more and more ionic. I'm getting really, I shouldn't be uncomfortable, but I'm getting more uncomfortable drawing these sort of covalent bonds to potassium. I'm going to skip rubidium, it's radioactive so nobody would ever do that. But you see what happens in terms of ionic character, now let me use a different pen color because I've crowded things in too much here. So now the ionic character is increasing as we go to these lower and lower in the periodic table. So things start to look more like carbanion minus. There is more partial negative charge as I march over here to alkyl potassium species. Okay, so let's keep that in mind. So this is a general trend and I expect you to know that things become much more reactive as you go in this direction and in a way that's not readily controllable. So I'm going to start off by talking about the simplest kinds of reagents of this class that we can think of and that would be where R is a hydrogen atom. So let's go ahead and talk about hydride reagents. I'm going to go ahead and draw the chemical structure for sodium hydride for you and it looks something like this, you don't have to draw this out. But you get the idea, this is the chemical structure for sodium hydride and I'm using dative bonds here. In other words, I'm not counting, this is not a Lewis structure. If I have three bonds to sodium, that thing ought to have two charges plus two charge on it. So this is, and you have to imagine carrying this out to essentially to infinity. But the idea is that each sodium atom and each hydrogen atom is locked into a cubic lattice. If any of you have ever used sodium hydride or potassium hydride or lithium hydride I don't know why you'd do that. But if you use sodium hydride or potassium hydride, you know it's not soluble in anything. It's not soluble in anything. It just sits there as this gray powder at the, so this is the chemical structure. It doesn't look like this. Not, it doesn't go into solution and dissolve. But we can't draw this every time we do an arrow pushing mechanism. So here's what we're going to do. We're going to do this. We're going to draw it like this. But every time you draw it like that, you're going to remember that it exists as this cubic lattice and that everything is happening on the surface. That none of the internal sodiums and hydrogens are involved in chemical reactivity. So true for sodium hydride, potassium hydride, lithium hydride, just pretend it's a monomer. And we're going to push arrows with this. So if you draw out sodium hydride, so this is a classical, sophomore organic chemistry reaction. I expect you to know that that sodium hydrogen bond is very basic, thermodynamically very basic. Kinetically it's okay for speed as long as you're not trying to pull a proton off of carbon. So you can get this reaction that will give you hydrogen gas and an alkoxide anion. And of course the alkoxide is going to want to coordinate to the sodium in a second step. So it's a good base. And as I implied up above here, things will become more ionic when I switch to potassium hydride. It'll become more reactive, more ionic. This will be much faster. I don't think I have ever worked with potassium hydride without having flames somewhere in my, when I'm trying to quench the reagent. So if you've worked with sodium hydride or potassium hydride, you know that there's a dramatic difference in the activity. I've never been sorry to use potassium hydride. I've been sorry many times to try to use sodium hydride. Let me give you a sense for relative rates for these kinds of reactions as you vary the metal. And we'll look at relative rates for deprotonation and basicity. So if I look at lithium hydride versus sodium hydride versus potassium hydride and you look at the relative rates, I'm going to assign lithium hydride a relative rate of 1. So relative to lithium hydride, sodium hydride is a thousand times faster for deprotonation. That's why you've never used lithium hydride because you don't like reactions that take a thousand days. You like reactions that occur right on the time that's like your whole thesis time is spent doing that reaction. If you get down to potassium hydride, so this is a recipe. If you want your chemistry to work quickly, if you like reactions that work well and quickly and you don't like reactions that are slow and lousy, use potassium hydride. It's screaming white hot aggressive and that's what you want. You can always cool things down and just quench things carefully. So this is what you want to see happening. So if you wanted to, you could, and I'm not saying that this is a great idea, but hypothetically you could use these to deprotonate carbon, but you should expect because this is a charge driven reaction. It's driven by the amount of negative charge on the proton. This might not be very good. Because there's very little, less partial positive charge on the proton when it's attached to carbon. You get less hydrogen bonding there. I'm not going to draw up that whole equation. Okay, so potassium hydride. When you really want reactions to work fast and well, you should switch to potassium for potassium hydride. And it usually, you'll usually have flames when you work it up. So be careful. Wouldn't it be great if every reaction you could just suddenly make a million times faster? And the answer is no. So what I want to do is I want to look at various organometallic reagents, green yard reagents, organomagnetiums versus organosodium versus organopotassium, organolithium. So what happens is we vary the metal when we look at the relative basicity and nucleophilicity of various organometallic species. And I want to start off by drawing out a reaction where let's just imagine, I mean this is one of the few where they've actually measured numbers, so let's just imagine what happens is we take this and we expose it to various different phenyl organometallics. Like you can imagine this being a phenylmagnetium bromide or phenylithium or phenylsodium or phenylpotassium. What happens is we change the metal. What you should expect is as we go through this series towards potassium down the periodic table, you should expect rates of that addition to increase. There's partial positive charge on the carbonyl carbon. So as you go from lithium to sodium to potassium, you should expect this to get faster. And the problem is that's not the only reaction that gets faster. There's a second reaction that speeds up as well as you go, as you change that metal. And that's deprotonation to make emulates. This is just like the SN2 versus Z2. Attack to make carbon-carbon bonds, competes with deprotonation of protons, elimination reactions. So there's two possible products that you can get here. So you can get the product for addition to pi star where you attack the carbonyl. Or you can get the product for deprotonation where you pull a proton off that alpha position. And what I'm going to do is I'm going to show you how that the ratio between these two types of products changes as we change the metal. So if you have phenylmagnetium bromide, then what you see is a pretty good ratio, 97 to 3, for addition to the carbonyl over deprotonation to make the emulate. This is why when you look at sophomore organic chemistry labs, you look at the kinds of reactions they do, they do reactions with Grignard reagents. It's you get relatively good rates of addition to the carbonyl versus deprotonation. The bottom line is that Grignard reagents are very nucleophilic, but they're not that basic. If we switch to phenylithium, it's not quite as good anymore. And this isn't to say that you can't find conditions to get better rates of addition, like switching the solvent to THF or something useful, but you can see already that as we draw, well, let's just stick with the phenylithium. The important point is what we do when we stay in this row. So lithium and magnesium are in different columns, but let me stick with this column here. As we go from lithium to sodium to potassium, because then I think it's clear. We're just dropping down in that single column, the first column of the periodic table. So as we go from phenylithium to phenylsodium, and I'm getting this, sorry, I'm just, so now things start to look are you grubly worse if you're trying to make carbon-carbon bonds? Finally, when we get to phenylpotassium, we don't see any nucleophilic addition to the carbonyl-carbon to make a CC bond with phenylpotassium. If you have to make a choice, and it's not that reactions aren't getting faster. As you go from phenylmagnetium bromide to lithium to sodium to potassium, the reactions are getting way, way faster, but not the reaction you want. What's happening is that this is speeding up. The deprotonation is speeding up at a much more as you switch to these larger metals. The deprotonation is speeding up much more than the addition to the carbonyl is speeding up. That's why you've never made a carbon-carbon bond by adding a metal potassium species, an organo-potassium species to a carbonyl. You've never seen anybody do that because they don't want to get deprotonation. And so what have you seen people use? You've seen people use Grignard reagents and alkylithiums. That's what you've seen people use. Okay, so let's go ahead and look at where we get these organo-lithiums from, or Grignard reagents. I would guess that most of you have probably made a Grignard reagent in some undergraduate organic chemistry lab. Maybe not everybody, but it's a very common lab procedure. I'm guessing that not many of you have made an organo-lithium that way. So let's talk about several ways that you can get your hands on some of these super reactive organometallic reagents. And I want to start off by one of the simplest sort of textbook ways to make things. Where you simply add metal. You just take metal and you mix it with an alkyl halide. So let's just suppose we take sodium and we mix that together with a chloride or a bromide or an iodide, some sort of alkyl halide. And we'll look at how the rates vary for this reaction. So let's start off just by considering this reaction of an alkyl chloride plus sodium metal. And the idea is that we're going to, so sodium chloride is a byproduct of this and the other product that we would get is this alkyl sodium species. So there's a stoichiometry. It takes two sodiums to do that. You can imagine the same thing for lithium or for potassium or for other metals. I'm taking sodium as sort of the prototype example. So the first thing that we have to deal with is what is that, how do you draw the mechanism for this reaction? And unfortunately, we will not be able to draw an arrow pushing mechanism for this reaction. Let me show you how to think about what's going on here. First thing is where are we putting the electrons in this system that involves a radical? Where is the electrons going? What's the nucleophile and what's the electrophile? The alkyl chloride is clearly the electrophile. And so let's just remind ourselves of what is the orbital in the alkyl chloride that's getting attacked? Well, there's no carbocation, so there's no empty P orbitals. I don't see any double bonds. There's no pi star. What we have is a sigma star orbital and that's this orbital right here. There's a sigma star orbital where the electrons are going. So let's draw that empty orbital here, sigma star for the carbon-chlorine bond. That's where the electrons are going and they're going in one at a time. So sodium is going to dump an electron into this orbital, this sigma star orbital. And the sad fact is we have no good way to represent this using Lewis structures. So I'm drawing sodium as an electron plus a sodium cation. So when we do dump one single electron into that antibonding orbital, what happens when we put an electron or two electrons into an antibonding orbital? You weaken the bond. It's like the antibonding orbital. If we put two electrons, we break the bond. If we put one electron, we weaken the bond. I'm not going to draw the whole butyl group here. And we don't have a Lewis structure that allows me to put an extra electron there. I was like, what are we going to do? Put a dash, that doesn't, that's not right. There's not three electrons. We don't want to imply that there's nine electrons on that carbon atom. So we're stuck with this representation of this. All these bonding orbitals down here, sigma orbitals filled. And then one extra electron stuck in an antibonding orbital that makes this bond, and let me erase this and draw this more suggestively, that makes this bond longer and weaker. That extra electron in an antibonding orbital is making that bond ready to break. And so I need a minus charge here. So we don't have a good Lewis representation. So we can't use arrow pushing for this. Sorry, we're not, if I ever ask you, draw a mechanism and it involves making an organometallic using a metal, you won't be able to draw an arrow pushing mechanism. Don't try. You can't do it. Right, that's, you have to start with correct Lewis structures and we don't have that. So now if we lose the chloride anion, so now we've got this long bond, it's weakened, it's got an electron in the antibonding orbital, that chloride anion is going to pop off. If we take these electrons and I'll draw an arrow here for this part of the reaction, if I take this pair of electrons in this bond and give it away here, the chloride will have the minus, what will the R group have? Yeah, it's going to have a radical there. And you'll have your chloride over there. So you can see the effect of weakening the bond because it makes it very easy to homilize that bond to give a chloride anion and a radical. So now we have this radical here and that can pick up another electron. Very fast, very, very fast to pick up another electron. Right, carbon with seven electrons is very unhappy. It wants to have another electron. And so then you end up with carbon minus and sodium plus and the C minus combines with the sodium plus. You could think of it as, in fact, well, if you already drew that, don't worry about it, but if you wanted to think of it as this, just forming a bond there, you could do that. I'm not going to draw an arrow for that. Okay, we don't have an arrow pushing way to draw this out. Don't even try. So what you'll find in this series is that the rates look like this. When you make organometallics from alkyl halide simply by mixing in metals, sodium or magnesium, alkyl iodides are better than alkyl or faster than alkyl bromides, which are faster than alkyl chlorides. That's the general trend that you should expect. Okay, so let's talk about the problems with these kinds of reactions. I don't think any of you have ever made an alkyl sodium species by adding sodium. You may have mixed sodium with an alkyl halide, but you've never isolated the alkyl sodium species and worked with it. I saw this in a review as referred to as chrono's behavior. Why is it that you've never made an alkyl sodium? Here's the problem. Let's suppose you take some sort of alkyl chloride, butyl chloride, and you add some sodium to that, throw in some chunks of sodium metal instantly. As soon as the reaction starts, you've got a problem. You've got this super basic, super nucleophilic thing floating around in the presence of something that is pretty reactive. So the Greek god chrono's was known in mythology for eating his children, and that's exactly what happens here. I don't know anything about mythology, but I'll always remember chrono's because I have no butylithium and whatever. So as soon as you start doing this, you've got a problem. And look at all the problems. Here's problem number one. Oh, SN2. Here's problem number two. Oh, E2. And you get both of those. I'm not going to finish the arrow pushing, but you get the idea is that you suddenly get these dimers in there, and I'm trying to draw eight carbons. I don't know if I've done that. That's one byproduct that you get as you're trying to make it. So your hope is that you'll have this alkyl sodium sitting in the pot at the end, but in fact what you've got is butene gas that floated away and then some octane. And that's just inexpensive. We have a special name for this kind of coupling. Together we call this a Verz coupling. We don't have a name for the E2 process. That's called a Verz coupling. So whenever you see organometallics couple with each other like that, you call that a Verz W, you are Verz coupling. Okay, so here's the general trend for how bad this can be. So how bad is this, let me draw it under here. This process here, how fast is it for the organometallic to come back and react with its precursor? And what you find is that the order of rates is that alkyl potassiums are the worst. So you've never seen anybody make an alkyl potassium because it's so hard to make. You get side reactions while you're trying to make them, alkyl sodium, very fast, these side reactions. So fast that these two are impractical. These side reactions are so bad that you just can't make these without super special conditions. So people have done it. They've managed to squeak out some enough alkyl sodium to make this work. But that's not a practical way. Things do become practical by the time you get to lithium. And of course you can buy tanker, train car tanker loads full of butyl lithium. So this is practical. And of course you know it's practical to make grignard reagents with magnesium, undergraduates do this in undergraduate labs. Okay, so it's very hard with sodium and potassium. Things are so reactive. They come back and they react with the starting material. Okay, let's take a look. So let's focus on alkyl lithiums now. How fast are those side reactions with alkyl lithium as we vary the halide? So in other words, if you're making an alkyl lithium from, well let me draw out that same scheme that I did before because I think it'll make more sense. And then we have to worry about these side reactions here to give you a mixture of E2 plus Virts coupling, the SN2. So how bad is it for alkyl lithium? It depends on the halide. How bad these reactions are. So let's look at what happens as we vary the halide here. And we'll look at the half life for these E2 plus SN2. Side reactions. We want to pay attention to that. Which halide should you start with? When you buy butylithium and you use it, how is that made? When you look at the half lives for these reactions, this is an ether, it's solvent dependent. This is an ether solvent. There's a 40 hour half life for these E2 and SN2 side reactions. That's slow enough to where you can make butylithium before it does this horrible reaction with the starting alkyl halide. If you try to make butylithium from an alkyl bromide, those side reactions are dramatically faster. And then finally, if you try to make butylithium from butyl iodide, those side reactions are so fast now, the E2 elimination, that you'll get just a bunch of octane and butene gas out of that. So the bottom line is end butylithium that you buy that they use in industry is made from alkyl chlorides. From butyl chloride or the chlorides in there. Okay, so how many of you have taken lithium wire and used it to make an alkyl lithium? Nobody, none of you have ever done that, right? That's for industry. When you go into the lab and you make some advanced intermediate that's a vinyl lithium or an aerolythium, you don't use this method. So let's talk about the method you use to make super complex structures with alkyl lithium bonds. This is great for making butylithium, sec butylithium, things like that. So when you're in the lab, you're not adding lithium wire, snippets of lithium wire to alkyl halides. What you're using is an exchange equilibration reaction. There's one member of this class that's not an exchange. So you're using exchange reactions. Let's start off by talking about these types of exchange reactions. This is one of the most common ways to make an alkyl lithium if you've got some advanced intermediate. So let's suppose you want to make an aerolythium and I'll just go ahead and draw this out. It's not interesting. It's just phenylithium because you can buy phenylithium but just imagine that that's a more complex aromatic ring. An indole or a thiophene, benzo-thiophene or something. The way you, wait, sorry, I want to, let me start off with a halogen there. And how is it that you would convert this into some sort of an aerolythium? And the way you do that is simply by adding butylithium. You could use other alkylithium, but butylithium is the cheapest. If you just do a simple stoichiometry analysis, if my other product is phenylithium, the byproduct has to be the alkyl halide. So this is the way you make alkylithium in the lab, typically. And the important point here is it's an equilibrium process. Well, nobody makes T-buly. They always buy T-buly and use that to drive reactions like this. Is that what you're, yeah, you can't, you can't have T-butyl halide here and get this to work. And we'll talk about that in just a moment. There's an equilibrium issue going on here. Okay, so what can X be? So traditionally in the lab, depending on X here, I'm just giving you the identity of X. Iodides are the fastest. So typically when you see this process of lithium halogen exchange, they're using an iodide. It's much, much faster than the next common type, which is tin. It's like tin, that's not a halogen. Trimethyl tin. Very fast metal, tin, lithium exchange. If you look at the relative rates where iodide versus trimethyl staining, it's 80 times faster with the iodide than it is with the tin. And then finally we get down to, and I am not sure why I have this, what I have is trimethyl tin greater than trimethyl tin. So here's what I'm assuming that I meant to have there, which is, so, okay, you may know trimethyl tin is very toxic and nobody wants to work with tricultins. Nobody actually does this. Everybody works with tributyl standings and they're less reactive. So if you want more reactive tin reagents, you have methyls on there, but they're toxic. They want less reactive ones, use tributyl tins and then they're non-toxic, but they're less reactive than trimethyl staining species. Okay, so when you use alkali dides at minus 70, it's essentially instantaneous. It's as fast as you're adding, even at minus 70, as fast as you're adding the butylithium in there, you've got instant equilibration. So just to give you a sense with how fast, so it doesn't really kill you to go down in reactivity like this, to go from instantaneous to seconds or minutes to use the tin. The tin is very powerful. So minus 78, minus 80, those kinds of temperatures, it's super duper, duper, duper fast. Let's go ahead and talk about what, so why would anybody want to do this tin thing? It turns out that you can, you can have tin reagents that have a broader range of substitution than iodides and I'll give you one example. Of something that you can make with tin and isolate and purify that you can't make with the halide. So here's a class of tin reagents. And you couldn't have an iodide here where this tin is. You can't have an iodide next to this methoxy lone pair. If you have a lone pair here and there's an iodide, the lone pair will push out the iodide. You'd never be able to isolate that. But no problem isolating this and running it on a silica gel column. And if you take this and you treat this with butyl lithium, then you can get very fast rates of exchange. And let me go ahead and draw it. You know, I haven't drawn the mechanism for these out. So let me come back over, let me finish drawing the mechanism for this. And it's the same as the mechanism for the iodide and I'll draw that for you as well. Here's the mechanism. Alkalithium is very nucleophilic. It attacks the tin. It doesn't do an SN2 reaction on the tin. It just attacks the tin. And so you end up with tin with five bonds. No problem. Tin could have six bonds and it wouldn't be an issue. We call this a stanate. The suffix to that 8 tells you there's a negative charge on tin. And as soon as you hear stanate, you should be thinking every bond to that tin is now nucleophilic. And so what happens if there's no electrophile poise nearby is that that will collide and you'll now convert that into an alkyl lithium. So you couldn't do this with an iodide. Even though you might like the fact that iodides are faster than tin reagents at this tin lithium exchange. The byproduct that you get in the end is this trimethyl steinol butane species. Okay, so what's powerful about this is that if you have a stereogenic center here, what you can tell is that it's a stereospecific transmetallation. If this tin was sticking, well, I had this center at this configuration, you'll get retention of configuration for where the tin was. In other words, it's not backside attack with the tin and the lithium. Okay, so it's the same idea. If you take butyl lithium and some vinyl iodide. In other words, if I draw the mechanism for this, watch this. There we go. Don't freak out over the fact. It's like, what is that iodide with two bonds to that? I can't deal with that. Sure you can. It's no big deal. Iodine can have seven bonds, no problem. It's not an octet rule thing. Iodine is not in the second row with carbon, nitrogen, and oxygen, and lithium. And so now that's an iodate. Notice that negative charge iodate tells me the bonds are nucleophilic so you shouldn't be surprised that that bond will now attack and give you a vinyl lithium plus the butyl iodide. Okay, so we're going to stop there. When we come back, we're going to talk about the third common way to make alkyl lithium reagents and that's from thio phenyl compounds through reductive lithiation.