 Okay, I want to start off with, so today we're going to talk about anions and really what we're talking about is proton transfers. It's a way to get to anions and think about proton transfers and maybe the nucleophilicity of the species that you can generate. And I want to start off by talking about what ought to be a trivial mechanism. This is the kind of reaction that you see in a sophomore organic chemistry textbook. And let's talk about this mechanism because it's not so simple. It's kind of stunning that this is the kind of thing they ask in introductory organic classes. And proton transfers are important at every single step in this reaction, even starting off at the first step of the mechanism. So you have to think about the things that have to happen here. We've got to get some sort of an oxygen nucleophile to attack water at that carbonyl group. We have to somehow get the phenoxide, the phenoxy group to sudden kind of a state where it's willing to act as a leaving group. And so that relates to proton transfers. So right off the bat we've got a question. What do we do first? Do we protonate the carbonyl first? Like a fissure of stratification? Or do we attack the carbonyl? This is a particularly easy kind of carbonyl to attack, a formate ester, much more reactive than regular esters. And so how do you decide? So notice I've asked you here for the mechanism. I don't do that on the problem sets. I don't say what's the mechanism for this reaction. I've never done that. I always ask you what is a plausible mechanism for this reaction. That's all I'm capable of asking you and that's all I'm capable of evaluating. So how do we decide which of these two starting steps is plausible? If you choose the wrong starting step the rest of your mechanism will be garbage. If you pick the wrong first step. You get to the next stage of the reaction. Your nucleophile adds in and you have to decide how am I going to get a proton off of here and get a proton onto the phenoxy group so that I can now push out that phenoxide leaving group? Should I come along with my acid and protonate the leaving group like this and then use and then come along and deprotonate the water? Or do I deprotonate the water first, right? Of these two steps that we need to transfer a proton, pull one off, put one on, in what order do we do that? How do you decide that kind of stuff? How do you look at those kinds of proton transfers and decide if they're reasonable or not? So let's talk about proton transfer equilibria. That's how you decide. You know something about the equilibria of those proton transfers. So one of the things we had an earlier lecture where we talked about proton transfers and I tried to get across this idea that you should assume that proton transfers are fast. You should assume that all proton transfers are fast and facile. Any proton transfer I've drawn here, you should assume it's fast and facile. And so it boils down to not is it possible for a phenoxy lone pair to pull a proton off of some acid? Yes, that's possible and yes, it's fast. What you should ponder is the equilibrium for that process good. So now let's talk about how do we estimate the equilibria for a protonation of an ester, for a protonation of an ether? And I'll give you two ways to estimate equilibrium ratios and that's what we care about. It's ratios that are intuitive. I have 10 times more money than I had yesterday. I have 100 times more electricity running through this wire. We think about ratios. That's easy to think about. Okay, so I'm going to give you two scenarios where I expect you to know how to use pKas to estimate ratios of protonated versus deprotonated species. And I'm going to start off with buffered conditions. So if you're a chemical biologist, these are the conditions that should matter most to you. You've got some sort of a buffer that either soaks up or delivers protons in order to keep the pH constant. And those extra ions, your buffering ions are present in large concentrations relative to the substrate. So let me go ahead and take an example. This will be a very organic molecule, not very biological. But the point is that when you're under buffered conditions, you have to pay attention to the pH of the solution. So we'll look at deprotonation of this nitro ketone. Very acidic. So to have this nitro group at the alpha position, which is already acidified by the carbonyl, it becomes way more acidic when we have the nitro group on there. It turns out that these protons here have a pKa of 8. That's the pKa of those protons. And so what's the ratio of protonated to deprotonated in solution? Is it reasonable for me to predict that I'm going to deprotonate to make an enolate anion and that that anion is going to attack? Well, it depends on how much is in solution. If there's a lot of it, then yeah, it's probably reasonable that that's a nucleophile floating around. And if there's very little of it, well, then maybe that's not such a reasonable mechanism. So the way you decide what the ratios of these are is you have to know something about the pH. And here's the way that I do this. I'm going to make a little chart here. And we'll just imagine that I know something about the pH of the solution. Let me make a list of pHs here. And I would expect you to be able to estimate the equilibrium ratios at any of these pHs. But there's one pH at which it is easier to estimate the equilibrium than any other pH. And that's when the pH is equal to the pKa for that proton. When I'm looking at a solution that has the same pH as the pKa for my proton, I know for a fact that the ratio of these two species is one to one at pH8. So if I know the pKa and I know the pH of my solution, I can readily have this sort of intuitive sense for what's going on. Half of my substrate in there is deprotonated. That's a lot of deprotonated stuff that's floating around. What happens at pH9 if I make a pH9 buffer? What happens to this ratio? If I go into the lab and prepare a buffer that's pH9 instead of 8, it's one to 10. When I go up in pH by a factor of one, this is logarithmic here, when I go up in pH by a factor of one, my solution is 10 times more basic. So let's just write that down. It should be totally intuitive of you now that there should be 10 times more of the deprotonated substrate in there. Now what happens if I go to pH10? See, this is the kind of math I know how to do. I expect you guys to be able to do this kind of math. Now it's one to 100. That's about 99 percent. I don't know how to work that math exactly. But I would say that's about 99 percent deprotonated. About 99 percent of my ketone is actually enolated pH10. And now I can go backwards. What happens if I pick neutral solution? Let's suppose that I'm interested in physiological conditions closer to pH7. What's the ratio there? So now it's 10 to 1 disfavoring the deprotonated side and favoring the neutral side. Only about 10 percent of my substrate in there exists as the enolate. And notice I don't care at all about the mechanism for deprotonation. We're assuming all of that stuff is fast. We're assuming proton transfers are fast. Don't get caught up in worrying about whether proton transfer, gee, should I draw that arrow? It's always a good arrow to draw a proton transfer, a curved arrow. And the issue is the equilibrium. And as I become more and more acidic by factors of 10, you can see what's happening to the amount of the protonated species that becomes larger and larger. So I expect you, anytime you know the pKa of a lysine side chain, a cysteine thylate, any sort of species, a phosphate group on DNA, I expect you to be able to estimate the ratio of protonated to deprotonated. Most commonly at physiological conditions but in any buffer that you're given. I expect you to have an intuitive sense for, oh, yeah, at physiological pH about 0.1% of lysine side chains exist as the free amine and have a lone pair that can attack things. That's the kind of intuitive sense that I'd like you to have. So there's some equation called the Henderson-Hosselbach equation. I can't even remember the equation. It's somehow or another you can get these numbers out of that but I don't know how to do that. So if you can work that equation and come up with this intuition, then good. Okay, so that's buffered conditions. That's for chemical biology. And there's some organic reactions where you buffer things, not so common. So now let's talk about unbuffered reactions. This is organic synthesis. Hydrolysis of an ester. You're working up some butylithium addition or a Mukayama-Aldal reaction. These are unbuffered reaction conditions that you set up and put into your round bottom flask and stir. So how do we deal with this? It's totally different now. So let me just give you this simple generic equation for a proton transfer like this. I feel like we've talked about this already but I can't remember whether we've talked about it. And there's a simple equation that I can use to estimate the equilibrium constant for any proton transfer. And I just need two numbers to estimate the equilibrium constant. We're not going to argue about whether it's easy for this to pull a proton off. You should assume that's always fast. What you need to worry about is what are the relative concentrations of these species in solution? Okay, so let's go ahead and be more pre, well I gave you some sort of equation here. And I'm even scared to give you the equation because the equation is so much more complex than the actual math you use to solve it. It looks something like this. The equilibrium constant for any proton transfer is equal to 10 to the pKa of the, how do I do this? The product minus the pKa of the reactant. And that looks so complex it scares me. So let's go ahead and take a real example so I don't get too scared by those, that kind of complex equation. So here's a reaction that wouldn't work under buffered conditions very well. Now we're looking at making an enolate by using a strong base, t-butoxide, to pull off a proton. It should be getting t-butanol as a byproduct. And there's a potassium ion floating around. Don't worry about that. That'll be our next lecture, how to deal with that. So the way I would work this equation, the equilibrium constant, how much of each of these species do I have present at solution? In other words, we're not going to worry about whether it's plausible to draw deprotonation. We're going to worry about how much of that deprotonated species is there. There's some there, how much? And so now we're going to work this equation. The pKa for this particular ketone, and I'm giving you numbers for DMSO. It happened to be the solvent for this reaction. In DMSO, I'm giving you the pKa's for this, or 24.7 for the ketone. And we'll talk about the solvent issue later. And 29.4 for the t-butanol. Those are the pKa's in DMSO. And so when I work this equation, I go like this. The equilibrium constant equals 10 to the 30 minus 25, 10 to the 5th. So what's the ratio of deprotonated to protonated? Is it mostly protonated or deprotonated? So a favorable equilibrium constant like this says it's almost all in this, not almost all. But it tells you that the ratio is 10 to the 5th, I'm going to say 100,000, I want it to 1. That means over 99.99% of this exists in the enolate form. And I get that just from these two pKa numbers here. So again, we're not worrying about whether it's plausible to pull protons off with t-butoxide. We're assuming that's always plausible. But we need this sense that, yeah, 99.99% of my ketone exists as the carbanion enolate. There's lots of it in my solution. Okay, so that's the sense that I'd like you to have. Now there's an issue here that I want to point out. And this equation is completely valid under typical conditions. And that is that your equilibrium constant is only going to be as good as the pKa numbers that you throw in there. So I gave you an example where this reaction was in DMSO. But you very rarely know the pKa's in DMSO. Or THF, or hexanes, or dichloromethane. And so what do we do about this? What do you know? And what will you know? You're going to know the pKa's in water. Because any idiot can do titrations in water. There are tables and tables and tables of pKa's in water. Somebody had a burette and they measured in something with phenolphthalein or some indicator. And they did titrations in water. So what I'll tell you is don't worry about the fact that you've mainly only got aqueous pKa's to work with. Use them anyways. You may be a little bit off where we'll talk about where you'll see differences between water and organic solvents. Even though you may not know the pKa and the solvent you're working with, use the water pKa anyways. It'll give you a good estimate and a good starting point. Okay, so I expect you to know ratios of protonated and deprotonated things. At every single stage in your mechanism where you've got a proton transfer step, you've got two species. And I'd like you to have an intuition for the relative amounts of those species and solution. So I'm going to give you a set of pKa's and I want you to memorize them. And there's no simple way around this. You just need to memorize these numbers. And more importantly, I want you to know the trends so that you can predict pKa's for vast numbers of molecules that I'm not going to show you here. So let me give you a set of pKa's. These are all aqueous pKa's relative to aqueous solutions. Some of them are extrapolated. And so let's go ahead and start off. What I'm going to do is I'm going to draw groups of molecules so that we can see trends. I'm going to group these together and I expect you to know these trends. So the first trend I expect you to know is the effect of electronegativity. So let me give you a series of compounds here. I'll start off with ethane and then I'll compare that to, let me draw this the same way, to methylamine and then I'll compare that to methanol. All I'm doing is I'm changing one of the, so forget about the fact there's different numbers of protons on the alcohol versus the methyl. What I really want to focus on is that I'm changing that second row atom from carbon to nitrogen to oxygen. I hope you already have an intuition that this atom becomes more acidic as I replace that with more and more electronegative groups. I hope you already know from your early organic chemistry courses that these are becoming more acidic going down this way and of those this is the most acidic. What you may not know is the numbers and so I want you to know those numbers. The pK for methanol, we're going to use a number 16 for an alcohol so it's actually, that's water, that's most alcohols, it'll range 15 to 17, use 16, it will never hurt you. It won't be a case where you'll be sad you didn't use 15.2 or something. Okay so an amine, the NH on an amine, very hard to pull those off, to pull the, that's LDA, lithium to isopropylamide is pulling the proton off of this. You get an amide anion, nitrogen with a negative charge, very difficult to pull off, 20 orders of magnitude less acidic than an alcohol. Okay finally you get to a simple alkane, there are no reactions that allow you to deprotonate simple alkanes. You should consider that to be the least acidic type of proton that's out there. So pK of 50 relative to water, 34 orders of magnitude harder to deprotonate than an alcohol. Okay so that's electronegativity. So the second type of trend I want you to know about and I want you to remember these numbers would be resonance. How does resonance affect pK? So let me go ahead and start off with that ethane molecule but instead of ethane I'm going to attach a C-C double bond on there. So now if you deprotonate this you'll have this allyl anion. So resonance stabilized allyl anion if you deprotonate that methyl group. And so how would we compare that to an amine versus an alcohol? All right and don't worry about the fact that there's a keto-toddomer that's more stable. You will see exactly the same pK if it's phenol it turns out. You could never have guessed that. Phenol and vinyl alcohol have the same pK. This and aniline same pKa. So the pK is now so I hope you know that this is the most acidic of the bunch the alcohol again there's an electric really we're just changing to more electronegative atoms as we go carbon, nitrogen and oxygen. So now we get over here to this vinyl this enamine is the name of that functional group and the number for that is about 28 so aniline an amino group on a benzene ring will have a pK of about 28. So significantly easier to deprotonate an amino group on a benzene ring than just a simple amino group. Then finally we get to a simple allyl like how easy is it to make an allyl anion? It's a billion times easier to deprotonate allylic CHs than just regular alkane CHs like the ones you have on your hexane solvent. So you're not going to deprotonate your hexane solvent but with the right base you can medilate here on an allyl position. Okay so we can make resonance worth more if I replace this CC double bond with a carbonyl group. That ought to be a better resonance stabilizer for any kind of an anion. I expect you to know the pKa for the alpha positions alpha to a ketone that's what this is. So now we're just talking about the pKa of acetone and I'm going to give you a number of 20 which is a good sort of average value for ketones, alpha to ketones. Okay if we look at an amid, if you're interested in the peptide backbone and understanding how easy is it to deprotonate the NHs on a peptide backbone or a protein backbone it's not easy, it's not impossible but there's a pKa of about 17. So a physiological pH of 7, you know, maybe 1 out of every 10 billion amide backbones is deprotonated. That's not very many. That's a very low concentration. And then finally we come over here and replace this OH or this double bond in this vinyl alcohol with a carbonyl group and we have a carboxylic acid. And that pKa is about 5 for carboxylic acids. It doesn't matter if it's 4.8 or 4.76. Okay so you can see that you see more of an acidity effect with carbonyl groups than with simple Cc double bonds. You should have expected that carbonyl groups are better resonance acceptors for anions. Okay so you need to know those numbers. Those are super common compounds that represent classes of compounds. Let's keep going with our trend. So this is a concept that we've already talked about. So this isn't as many numbers as you might think it is because look I'm giving you the same number. I just gave you this number. What is it for ethane? Fifty, yes, same for hexanes, cyclohexanes, all pretty similar, close enough. So we already talked about 50. But now what happens if we put a double bond in here? It's not the same anymore. Or if we come over and we put a triple bond in between these. So we're changing the hybridization of the atom. And so what happens to the acidity of these? They're going to become more acidic. And it can be quite dramatic here, the effects. Again these are all CHs. There's no resonance, anything going on, atoms with higher P-character. So if you look at the carbanion here, you know the carbanion that you would get by deprotonating this is very nucleophilic, very reactive and very unstable in a way. It's unstable. There's lots of P-character here. But as you reduce the P-character it becomes less reactive, it becomes more stable. So finally when we get down to SP-hybridized carbon, now there's not all that P-character that makes things so unstable and reactive. There's plenty of 50% S-character here in this alkyne lanion. And so that's what's accountable for this and what is the math here? 26 orders of magnitude of acidity. All the same atom, carbon. And we're just changing the hybridization. So I hope by now in this class you respect the effects of hybridization on reactivity, this gives you a sense how that affects acidity. Okay, let's keep going. This won't really make sense unless we walk back and we compare it to the amine way over here. See this methyl amine? PK for methyl amine or diisopropyl amine or any NH on a regular amine is about 36. And so what difference does it make to have a positive charge on the amine? So now I'm talking about a different compound. Not an amine, I'm talking about the PK of an ammonium ion. This is one of the most common mistakes is to confuse an amine for an ammonium. Amine NHs are very hard to pull off. You need butylythium to deprotonate a simple alkylamine. This, anything can deprotonate, well not anything, but just about anything can deprotonate that. And the PK of this is about 11. That's a lysine side chain in a biological context. And so what about a protonated alcohol? Well, protonated alcohol is a little lower, but a protonated ether. You know, it's protonated alcohol is about minus 2, minus 3, protonated ether. Again, the important point is the effect of charge. If we compare this to an alcohol, you know, it's pretty acidic, but when you have a positive charge on the oxygen, it's very acidic. So it's very easy to pull protons off of oxonium ions. That's oxygen with three bonds. We call it an oxonium ion. Okay, so what are some more, let me keep drawing some more species. So I have an ammonium ion. What about an aminium ion? So we talked about aminium ions as reactive intermediates. It's much easier to do nucleophilic attack to aminium ions than carbonyls. The whole field, not the whole field, most of the field of organocatalysis is aminium ions. So these have PKs on the order of about five to, I'll just give you the number of five. And it's not, right, it's not the fact that there's just a simple isolated amine. If you look at the PK for this, for pyridinium, it's five. It's the same. Aminium, pyridinium, it's about five. Okay, protonated oxygen on a carbonyl group. Now it's very, very hard to protonate that. It's extremely hard. If you want to protonate, I'm talking about simple ketones. If you want to put a, protonate a simple ketone, what you see people doing is they add sulfuric acid or hydrochloric acid, serious business acids minus seven. Incredibly hard to protonate the carbonyl of a simple ketone. And let's put these into context. And I, it's like, I don't know why I have this here. I, maybe just you can seem like a no at all. But if, if 16 isn't good enough for you and you somehow need to know exactly the PK for water, you can use 15.7. But I've never had that be any better for me than, than 16. So if you want to use that. The main point here is I want to compare this to hydronium ion. So we know a number for that. And that's minus two, minus 1.7. So remember somewhere here I noted this, oh, this difference. That if you have an, a protonated here with two alcohol groups, it's minus four. Hydronium ion is minus two. And a protonated alcohol where there's just one alcohol is right in between, minus three. But again, I, you know, if you just want to memorize kind of minus three for an oxygen with three bonds and a proton, I'm sure that'll always do you right. Okay, so next trend, the last, no, two more trends to talk about. These are the kinds of numbers where you're going to be in your oral's exam and they'll be at the board drawing something out. And then suddenly you'll hear this question from them in back of you. What do you think is the pKa of that pro, and you'll turn it on and be like, four? And your committee will look at each other, right? They'll wonder why you didn't remember that, that number. So this isn't just for this class. You guys need to know these numbers in perpetuity. Okay, so obviously I'm talking about getting to bigger and bigger atoms as we drop down from second row to third row to fourth row to fifth row atoms. And you'll see a concomitant increase in acidity, aqueous acidity. There's a huge leap when we go from, and this kind of corresponds to leaving groupability. There's a massive leap when we go from fluoride to chloride. So HF is not like hydrochloric acid. It's nothing like hydrochloric acid in terms of acidity. What it has is a nucleophilic counter ion, fluoride. So HBR, and one that's very basic. It's not that different from HCl in terms of acidity. I can't think of in many cases where you get anything in terms of acidity by switching from HCl to HBR, what you get is a more nucleophilic counter ion, bromide. And if you switch to HI, again, if you just want a stronger acid, you're not going to go grab a bottle of hydroiotic acid instead of HCl. You'll go to sulfuric acid. Okay, so finally we're down to, I think, the last important comparison for this series. And that would be thiol versus alcohol. So alcohol has a pK of about 16. What happens when we drop from the second row OH to the third row, SH? So now the pK is about 11. And the pK is actually very considerably depending on what's attached there. So this is getting close to a cysteine side chain. You'll see numbers in the 9s sometimes for the cysteine side chain on protein. pK around 9, so there's considerable variability here. When you start to put polar groups nearby, you'll see that pK drop a little bit more. The important point is it's 100,000 times or more acidic than a regular OH. It is very easy to deprotonate a thiol. If you take any thiol and you put it at physiological pH, there is a substantial concentration of the thiolate anion in solution. And they are super nucleophilic thiolate anions. Yeah, that's a crappy drawn ethyl, badly drawn ethyl. Okay, so now polar effects. You might call these inductive effects would be the term some people would use. Here's a cool number because we have so many tetrahedral intermediates that we're interested in. How easy is it to deprotonate species like this where we have, well, whatever, where we have more than one electronegative atom attached to the same carbon. So when you look at the pK for this, you kind of should expect that it's easier to deprotonate this OH when there's an electronegative atom close by. And what you don't know is how much easier. The nearby electronegative atom should make it easier to deprotonate that. And it turns out the pK is 13.4. So it's the number I keep in mind is 1,000 times more acidic than, but I don't know that the 0.4 really matters. So it's easier. You might have known that. You might not have known how much. Okay, let me redraw a molecule that I've already drawn for you. So what's the pK of this one? Five, yep. I'm trying to understand why I wrote these decimal places because I'm just going to write five. I have 4.8 written in my notes. I don't care whether you memorize 4.8. Okay, now let me give you the number for fluorosidic acid. What happens if we put a fluorine atom here? Right now, if I deprotonate this proton, there's no resonance structure I can draw here in which the fluorine is going to help that. If there's any effect of fluorine, it's kind of like a through space interaction. You know, the nucleus has more protons in the nucleus of fluorine, so it's more positive. That'll help stabilize negative charge way over there on the carboxylate. So that fluorine ought to make it easier to deprotonate here, and it does by this much. So it goes from 4.8 to 2.7. It's 100 times more acidic to have that fluorine close by like that. And as you might guess, to have three fluorines, very electronegative, trifluoroacetic acid is close to zero. It's essentially zero. So 100,000 times more acidic than acetic acid when you see people use TFA, trifluoroacetic acid as a catalyst. Okay, so those are PKAs I expect you to know. You just have to memorize those PKAs. But the trends I expect you, those ought to be much easier. I hope by knowing those PKAs, you'll always be able to rely on some trends that you can extrapolate. Okay, I want to look at some trends here for alkyl lithiums. And let me start off by saying that it is not a bad conceptual tool, especially when you're first starting off. At this point, I don't really expect you to make this jump, but when you first start off in organic chemistry, you take your first year of organic chemistry, this starts to look very bizarre. It's like, what's this lithium carbon bond thing there? And so most books encourage you to think of this as a methyl anion. Trust me, there is no methyl anion in a bottle of methyl lithium. There's no free carbon anion. Every single carbon is bound to at least one lithium and it actually forms aggregates where it's bound to two. So people draw this because they just don't like that. They don't like drawing a nucleophilic sigma bond. But there's nothing you're going to do with this lone pair that you couldn't also do with this carbon lithium. So just think of this alkyl lithium is about the same in terms of nucleophilicity as you might consider that lone pair to be. And the main point I want to make here has to do with the basicity and reactivity of these species. I want to draw a trend here as we increase the size, the number of carbons in each of these species. So I've got a whole set of alkyl lithiums here. And if you think about basicity, what happens if you take a syringe that has methyl lithium in it and you squirt it into the air or into some water? It's not a real show. There's not a lot to see going on there. But if I take a syringe that has T-butylythium in it and I squirt it into the air, it will instantly burst into flames. The second the T-butylythium touches the water vapor in air, it so aggressively deprotonates the water that the heat it generates ignites the solvent and causes a fire. It's instant fire. There's a huge difference in reactivity. So if you look at, and how did I phrase this? This is reactivity, not stability. So if you look at the reactivity of these, the more alkyl groups that you put on a carbon atom increases the solution phase activity. And why am I being so careful about talking about solution phase reactivity? Because if you do calculations, and this is where I'm just pointing out that you have to be careful with calculations, gas phase calculations. If you do a calculation on just a single ethylithium molecule that you've drawn on the computer screen and you compare that, here's what you would find. Or carb anions. I'm going to switch this, instead of talking about this, I'm going to talk about alcohols. Sorry if you've started drawing that. Acidity. So T-butanol acidity, which of these should make a more stable anion as the way you can think about this? Isopropanol versus ethanol, because I think this will be more useful to you, versus methanol. If you look at the gas phase acidity of these, it's the opposite of what you expect in terms of substitution, it turns out that T-butanol is more acidic in the gas phase than isopropanol. And that's more acidic than ethanol and that's more acidic than methanol. It's the reverse of what you see in solution phase. T-butoxide turns out to be more stable than methoxide in the gas phase. In general, I feel like I pointed this out before with cations, in general when you look at the stability of larger molecules in the gas phase by calculations, they seem more stable. And let me give you a way of thinking about why is it that when you do computer calculations, that larger things end up seeming more stable. Let me go ahead and compare for you these two types of molecules here. I'll just start off with methanol or methoxide anion. The thing that will stabilize either of these two molecules is the donation of filled orbitals into unfilled orbitals. So over here, I look at, here's my nucleophilic orbital. What can it donate into? Well, it can donate into this antibonding orbital for the CH bond. So I'm donating now into sigma star for the CH bond. If that were really true, I would expect a little bit of nucleophilicity out of this bond. If I'm really donating into that bond and weakening in it, I would expect it to be slightly nucleophilic. And I'll show you later some reactions where exactly that happens because that donation is real. But now when I come over here and I look at this larger molecule, I can donate with this and that makes this bond more nucleophilic. And if I have another bond here, I can take that, those electrons, and now that nucleophilic orbital can donate into this one. And as I keep going, the more bonds I have, the more anti-bonding empty orbitals I have, and the more empty orbitals I have, the more chance I have to donate filled into unfilled orbitals. So in general, when you do gas phase calculations, you'll see this effect that larger ions, anions or carbocations, look more stable. But when you put them in solution, you don't see, you see almost no effect of that kind. Solvation becomes a more important issue. And we'll talk about salvation in just a moment. Okay, so I'm going to go ahead and, I added some little table here and I'm going to go ahead and skip that. What I want to talk about is the effect of solvents. What effect does solvent have on stuff like this? On chemical reactivity. So I'll give you some rate cons, some relative rates for a Finkelstein reaction. That's just a simple SN2 reaction where you displace a chloride with iodide. And we're just going to look at the effect of solvent on this SN2 reaction. So how does solvent affect this? So I'm going to give you a series of solvents. I'll tell you what the dielectric constant for these are and I'll tell you the relative rates, not the absolute rate constants, but the relative rate constants. As I compare different solvents, I'm going to start off with water, very powerful solvent. Nitromethane, very powerful. And then dimethyl formamide, I'll just write DMF. And then finally acetone. So I'm moving to grease here and I don't have numbers for hexanes. I guess that would be no further down the list here. What I would see is I go to, from polar solvents to less polar solvents is an increase in rate. So I'm going to assign relative rate for SN2 reaction in water as 1. If I simply switch to nitromethane, it's 14,000 times faster. You need to use this kind of thinking in the lab when you set up reactions. What if you could make your PhD 14,000 times faster? Who wouldn't want to do that? I go to DMF. Now it's 800,000 times faster just by switching to a different solvents. I walked over to the shelf, I grabbed a different bottle. And now things are 800,000 times faster. And acetone, that's a 4. So now we're 1.4 million times faster. So instead of talking about why it's faster, let me talk about why these things are getting slower as we go to more polar solvents. It's not all the solvent bulk dielectric constant. So water is the most polar in terms of bulk dielectric constant. There is a trend here. But you get to this stage and there's not an exact trend. There's something else going on here, especially when you look at the nitromethane versus DMF. Let's talk about water. That's the biggest jump. It's like why is water such a sucky reaction? What's such a sucky solvent for this reaction? It's not just polarity for the water. There's another reason why water is not good. So first of all, let me just talk about this. So I have iodide. There's a minus charge. I have a substrate here, chloride. And there's a calomic advantage to having these things get close. There's a partial positive charge on carbon. The closer they get, the more Coulomb's law, and I could calculate this, the more Coulomb's law tells me that it's energetically favorable to bring these closer together. And if you have a dielectric here, the larger the dielectric constant, the bigger this is, the less these charges can see each other. Dielectric, high dielectric hides charges. They can't see each other when the dielectric constant. So that's part of this effect right here. But there's a second major effect here. So dielectric is one origin of this set of differences. But there's a second effect. And I'm going to draw a picture of this iodide anion. And the problem is when you're in a protic solvent, this is, it's the protic solvents that really screw you on this. It's not just this, this ion's not seeing each other effect. The problem is this iodide is already, and I don't know how to draw this. I'll draw this as dashed lines. The iodide is already, in a way, kind of reacting because it hydrogen bonds with the solvent. When you have solvents that are protic that can form hydrogen bonds with nucleophilic species, it's almost like it's already donating into an antibonding orbital. That's what the hydrogen bond is. And with all these distractions around here, you can imagine why it's not so easy for this iodide to get interested in attacking anything. So hydrogen bonds distract nucleophiles. Because most nucleophiles, good ones are anionic or have anionic character. So this is reason number two for the big difference. So there's the charge screening effect of polar solvents and then there's this protic solvent distraction effect. Okay, it's harder to distract bigger atoms. It's harder to get the hydrogen bonds organized around a big iodide, big iodide anion. What you'll find is when you have smaller atoms, it's very easy to form very tight hydrogen bonds. And let's take a look at the practical implication of that as we go to the second row atom, sulfur, ethyl. It's harder to distract bigger atoms and it's easier to distract smaller atoms. So when you put fluoride anion in aqueous solution, this is not very nucleophilic. In organic chemistry, nobody uses aqueous solutions of fluoride anion. Who would do that? The hydrogen bonds between water and fluoride totally shut down their activity of the fluoride. If you're interested in using fluoride as a nucleophile, for example, to attack silicon, what you use is this, tetrabutylammonium fluoride. I'm not going to draw the butyl groups. That can't form hydrogen bonds with the fluoride. If you form any hydrogen bonds with the fluoride, it's over. You've shut down the nucleophilicity. So here's another sort of, I'm not sure what this is. This is just a related factoid. If you're into chemical biology, you should keep this in mind for things like native chemical ligation of proteins. You have cysteine thiols on proteins and amines on proteins. And they're in equilibrium. There's an equilibrium between thiols and thiolate anions, between amines and ammonium. Most lysine side chains are protonated in aqueous solution. Most thiolates are also protonated. But the small amount of thiolate anions that are present in solution, in any aqueous solution, are so nucleophilic that if you have a protein or if you have equal amounts of thiol and amine in a buffered solution, I'll call this pH 7, it doesn't matter what the pH is. You'll find that thiols are more nucleophilic. So you can see what's the problem is if you're under acidic conditions, almost all the amines are protonated. There's no lone pairs on those things to act as nucleophiles. At least there's some thiolate anion. If you try to go to more basic conditions, so at least you have an even playing field, if you go to higher pH to try to get some of your lysine side chains deprotonated, well then you're going to deprotonate even more of the thiols, you'll have even more of the thiols around that can act as a nucleophile. So generally, when you have thiols in aqueous solution, expect them to be way more nucleophilic than lysine side chains. So unless you have 300 times more lysine side chains than thiols, typically it's like 40 to 50, then expect the thiolates to act as a nucleophile, the cysteines. Okay, we're going to stop there. We've got one last little piece that we're going to cover when we get back. I have finished grading about half of those two page proposals and Chris is going to give those back.