Freshman Organic Chemistry II

CHEM 125b - Lecture 5 - Solvation, H-Bonding, and Ionophores

Chapter 1. Puzzle on Alcohol Oxidation Mechanisms [00:00:00]

Professor Michael McBride: So as we move from radical to ionic reactions, HOMO-LUMO reactions, we have to take into account solvation. So they will do that, and H-bonding, talk about ionophores and Brønsted acidity. We might begin nucleophilic substitution.

First, the answer to that puzzle that we had for homework last time, the question is how you get to the aldehyde from the alcohol--by free-radical chain substitution reaction. And do you remember that you would substitute the hydrogen with the chlorine, you know how to do that. And then if you have a base around it--it doesn't have to be a strong base--you can pull off the proton and lose the chlorine--chloride simultaneously in an elimination reaction to generate the carbonyl, and I think a lot of people got that.

And notice that those things happen at the same time. That the departure of the chloride opens up a LUMO that helps take the electrons from the O-H bond. So the O-H bond is weakened by the fact of the chloride leaving. So all those things are happening at once. And we'll talk more when we get to elimination reactions about the timing of that process.

So the base that removes the proton could be a very weak base. It could be just the unshared pair of an alcohol or even of a bisulfate anion.

Now that might not be the way it actually happens in 30% sulfuric acid. So I asked you think of another mechanism, one involving HOMO-LUMO substitution. So the question is what's the HOMO and what's the LUMO. So the HOMO's an unshared pair on the alcohol, and the LUMO is the [sigma]σ* of the nitrogen-chlorine bond, which is especially low because of the positive charge on the nitrogen.

So we can do a substitution reaction. And this will then put the chlorine on to the alcohol. The positive charge you're going to get rid of, which you can just do by the amine that was generated in the same process, it can take off the proton, which you notice is a substitution reaction at hydrogen substituting nitrogen in place of oxygen on the hydrogen.

So then we have the same compound as above, except the hydrogen and the chlorine are exchanged in their position. So you can do the same thing again--bring in a base; pull off the proton and lose the chloride. Now you don't usually think of bases taking protons off from normal carbon atoms. But this isn't a normal carbon atom because the chloride's leaving. So there's something that can take the electrons as the base pulls the proton off. So the same kind of elimination, just backwards, gives the same product. So that was the mechanisms that I was hoping you might be able to see.

Chapter 2. Solvation, Boiling Points, and “Intramolecular Solvation” [00:02:58]

So in the last lecture we were talking about what holds molecules together other than bonding, about non-bonded interactions. Those are visible in the physical properties of the alkyl halides in the boiling point. And alkanes too--any liquid for that matter. So we looked last time and saw that the dipole-dipole interaction could hold alkyl halides together. But that can't be all it is, because that gets large and then gets small again, and the boiling points keep going up all the time. 

So there must be something else, and that something else is polarizability--the fact that if you have a lot of electrons and they're far out in high-n orbitals so that they're easily sloshed around, that is, it's polarizable, then you can have interactions between the dipole and something that gets polarized by the dipole, or indeed between two instantaneous dipoles. And we see that iodide here is much more polarizable than fluoride. So that's the other mechanism that's holding the molecules together. That is, between molecules, the bonding between molecules, the non-bonded interactions.

Then we said last time let's look at hydrocarbons and see what they tell us at butane and isobutane, neopentane, and n-pentane (normal pentane). So here are the boiling points for n-butane and isobutane. And you see that as you make it more branched, it's an isomer, the same atoms. But as you make it more branched, the boiling point goes down. It's easier to pull the molecules apart to make them in the gas phase.

The same thing is true of pentane. Normal pentane you go down eight degrees to get to isopentane. And neopentane goes down by another 18 degrees.

So polarizability does its job well only when things are close together. Remember its 1/RN dependence, so it falls off rapidly, the energy you can get from this. So if you have the same atoms, for example, here in normal pentane, so we got five in a row, we have to get the atoms close together to allow their polarizability to hold the molecules together. So with these you can put them side-by-side, and that's great, and all the atoms in one molecule are near all the atoms in another molecule.

But if it's a branched compound like neopentane here where the four methyl groups are on the same atom, then you can't get them nearly as close to one another for purposes of having the polarizability do the trick. We have a picture here so that you can see it better on the next slide. 

So it's atoms that are near the surface of the molecule that count. Atoms that are deep down inside of molecule don't get close enough to their neighbors to profit from this very much. So there's above n-pentane and below neopentane, and here you can see how much closer you can get the molecules. And if you'd like to see whether you can get them closer than I suggested in holding here, I'll pass these around and you can hold them and have a good time.

Now you can see something a little bit like that in the heat of formation, the stability of the molecules. Notice that the heat of formation of neopentane is larger, more negative, more downhill to form neopentane than to form n-pentane. Part of that must be the same reason that the polarizability could operate within a molecule as well as between molecules. And when you get all the atoms not to be spread out but to be more in a ball, then the atoms within the molecule are near one another.

So what you lose in terms of polarizability from one molecule being adjacent to another one, you gain in interactions within the molecule. So that's not the only reason for this, but that's one reason that branching makes molecules more stable. So it's like solvation, except it happens within the molecule rather than between molecules. And you can see the same principle in a more pronounced way when there's an actual charge to do the polarizing.

So if we look here at the gas phase ionic disassociations, so how much energy does it take to break a chloride away from R to make R+. So we look at the stability then of the R+ cation, and you can see that it's ever so much easier--53 kilocalories easier to make t-butyl than it is to make methyl. Now there are a number of factors that are involved there and we're going to discuss others as well. But one of the factors is certainly that the positive charge gains stability by polarizing the material that's around it. And the more material you have around it, the closer to the positive charge, the more stability you're going to get.

So that's one of the factors that's involved here. It's not the only factor, but this intramolecular solvation--I mean that's sort of a stupid thing to say--but that must be one of the factors that helps explain these trends. 

Now here we look at the boiling points of the normal alkanes with different chain lengths, and halogen on the end. Or in the first case, hydrogen. So, here n-alkanes has hydrogen on the end of the chain, then fluoride, chloride, bromide, iodine. This is a picture from the textbook, the Jones textbook, and here we see them colored so you can see what atom's involved.

So one would infer from this that for some reason--I mean just the impression it makes on me as a viewer is that something suggests that the iodides are low boiling. You see what I'm saying? That curve falls below the others. 

But what does molecular weight have to do with boiling point? Notice what's being plotted here is molecular weight against boiling point. But what relevance does molecular weight have to boiling point? There's sort of an intuitive idea that if something's heavy it's hard to get up in the gas phase. But that has nothing to do with boiling point. That's pure nonsense. 

The question is– is there a better way to plot this to see what's going on? And I think there is, which is instead of plot it against the molecular weight, plot against the length of the chain--how big is the molecule. So if we do that, so we plot n here in H, CH2 taken n times, and then X. And now we'll change X from hydrogen the fluorine to chlorine to bromine to iodine. 

Now fluorine is almost the same as hydrogen. When it's a really short chain, the fluoride is a little bit like--that's methyl fluoride--that's higher boiling than hydrogen. We already explained that--the dipoles that are in the CH3F can interact with one another in a favorable way. But once you get a real long chain, that's not a very big fraction of the total and it comes to look pretty much like the alkanes.

But now look what happens if you go--so the dipole moment is bigger for fluoride than it is for the hydrogen. But the polarizabilities are about the same. But chloride is more polarizable--it has more electrons further out. So although it has essentially the same dipole moment as the fluoride, its polarizability, its greater polarizability is clearly holding things together. And the same thing is true when you go on to bromide, which is still more polarizable. And to iodide, which is the most polarizable of all. 

So in a sort of back hand way it's related to molecular weight, because the more electrons you have in the atom, the further out they are, the more they slosh around, the more polarizable they are. But the electrons have nothing to do with the weight. It's the nucleus that has to do with the weight. But more electrons mean more neutrons and more protons in the nucleus. So it's sort of parallel, but I think it's more informative to look at it this way to see what it is that's really holding the molecules together.

Chapter 3. Solvophobic Forces and Hydrogen-Bonding [00:11:45]

Now you've heard the ancients say, “like dissolves like.” And you've seen things like soup that has globs of the fat floating around and obviously not mixing with water. And you've heard tell of solvophobic forces, probably hydrophobic forces. So the name implies an antagonism between water and the alkane, the grease, that tends to push them apart. So you get islands of grease in a matrix of water. 

Now here’re droplets of mercury on a cover slip, on glass. And you can see that the mercury beads up--it doesn't spread out and wet the glass. So it seems that mercury doesn't wet glass. So there seems to be a phobic force between glass and mercury that tends to push them apart to keep them separate from one another. Here we see droplets of mercury in a polyethylene container, and you can see that it beads up--that mercury doesn't wet hydrocarbon either. So it looks like they are antagonistic to one another. The same thing is true of water. Water doesn't wet a waxed car or a hydrocarbon. That's hydrophobic forces--that's like the grease, they stay apart.

Now, if that's so, can you tell me should mercury repel water? So if we touch a drop of water to a drop of mercury, should they stay apart as beads or should they spread on each other? How many think they should be antagonistic and form beads? How many think they should spread on one another? Well, there's a slight excess there, but pretty close to a tie, and I think there are a number of individuals not voting. So we can do the experiment. See how the water comes over and grabs on to the mercury?

So clearly there's affinity between mercury and water. But, in fact, there's also affinity between mercury and glass, and there's affinity between alkane and water. It's a misleading term to say “phobic,” that these things are avoiding one another. Clearly, mercury attracts water. But alkanes and water or mercury and glass don't repel one another. There's attraction between them as you expect. If nothing else there will be the mutual polarization of one atom by another across this boundary that will tend to hold things together. 

But mercury is particularly good. It has so many electrons out there that could be polarizable. Mercury holds to mercury very, very tightly. And water likes to be near water for hydrogen bonding, which is what we're going to talk about. Water likes to be near mercury, but not as much as mercury likes to be near mercury.

So water with mercury is a good interaction from the point of view of the water. The water is a dipole, the mercury is polarizable, there's good attraction there. But mercury doesn't want to change its shape. It wants to have as many mercury atoms as close together as possible. It's like the alkane, the neopentane--the atoms want to be close to one another to be low in energy. So you don't want to distort mercury because it has so much attraction between the atoms. But water attracts mercury better than it attracts itself. So it spreads out on the mercury.

So these solvophobic forces are not really a question of things being antagonistic to one another. It's a question of something really wanting to be close to itself for some particular reason. Mercury with mercury is very good attraction. But water with water, too. Water will attract grease. But it attracts itself much better. So it wants to stay connected to itself.

This brings up the question of what is it that holds water to itself. And, of course, we heard last time about hydrogen bonding. So let's think about hydrogen bonding. So water has a dipole. So you can imagine that two of them would line up this way--positive to negative, the dipoles.

But actually, this is the calculated dimer structure for water, calculated in 2000. Water gets worked on so much and it's been worked on forever. Anybody that has a theoretical technique tries to do it with water. And the main thing is that although they hold together, they don't hold together very strongly. The dimerization energy's only about three kilocalories per mole. And there's a lot of flexibility in it because it's rather weak. But the very lowest energy, according to this particular calculation, is this one.

Now there's a hydrogen between the two oxygens--the hydrogen of one of the molecules is between the two oxygens. But let's look at the bond distances. There's 95 picometers, that's 0.95 Å; same over here on the left. But the bridging one is 0.9639, 0.96. Well, it's longer, but it's not much longer. So it's lengthened by only 0.5%. So there's not much electron pair from one oxygen donating into [sigma]s* of the other one--that's what you would expect for a bond. It hasn't lengthened the bond to the hydrogen.

Now, how would you expect the waters to line up according to their--we looked at the charge last time, the dipole, but water isn't just a positive charge here and a negative charge here at a certain distance from one another. The electrons are spread all out with different densities. So you can analyze it as multipoles, not only a dipole moment, but also a quadrupole moment, which is plus plus minus minus, and an octupole, and a hexadecapole moment. This is hard to do in your head.

But one thing you can do with Spartan or other graphic programs is draw this electrostatic surface potential that we've talked about to see where on the surface a proton would be happy. That's where it's red, or unhappy where it's blue. And the energies are –47, very favorable in the blue area, to +60, very unfavorable in the red area for a proton at the surface.

Now let's look at the one on the right there and let's rotate it so we sight right along the H-O bond. And if we look we see that it's blue there--a bad place to put a proton, a good place to put something negative. And the color scale here goes instead of –47 to 60, well this is ­–45 to 60. But we can home in on it to get the colors to show what we want by going +55 to +60.

Now you see that there's a very small region there, which is the best place to put something negative, the worst place to put something positive. And if we look at it we see that as you move the proton up and down there or back and forth, you change by half a kilocalorie per mole if you go up and down by half an angstrom, which would correspond then if you went up there--that's where we're looking at on the surface. So that's plus or minus 9 degrees up and down. So that's a fairly tiny target, although it doesn't--half a kilocalorie isn't enormous. You could go further and not lose too much.

But look at the oxygen end. If we rotate that one and look at the water from the oxygen end, we see there's a very big streak of red. That's 10 kilocalories total range on the color scale. If we go here down to 5, which is the same scale being used over here, it's way spread out. And even if we go further to only one and a half kilocalories per mole, for the entire color scale you can see that that red range goes quite far up and down, 1.38 Å, plus or minus 25 degrees. And the energy change is only 0.05. So it's very big red target that you can tolerate. 

So you have a tall target there and a small target on the other one. Now when you put them together this way you see you can bend these down and get these hydrogens far away from that one and still have very good contact between the places that complementary charges are good to sit.

That's a very qualitative sort of analysis of this, but it's a sort of funny minimum energy structure for water. But the important thing is that it's consistent with just how the charges of isolated waters would interact with one another, just charge-charge interaction. There's no special bonding. Quantum mechanics didn't come into this. Once you have the water molecules you'd bring them together and that's a reasonable way for them to fit.

The disassociation energy is not large. It's 3.3 kilocalories per mole, but you can get other waters around and about. So taking a water molecule into the gas phase out of water is much more, of course, than 3 kilocalories per mole.

So from this perspective what makes hydrogen bonding important is that you have these dipoles, but they can get very close to one another, because remember the energy falls off very rapidly with distance. Because hydrogen is so small it can get very close to the oxygen end of a dipole. So it's just dipole-dipole interaction like you have in any of these other things we were looking at, but they can get really, really close to one another. So you get more energy in the case of hydrogen bonds than you do for other dipoles. So nothing special. There's no real bonding to it. It's just that the molecules can come together and be low in energy that way.

Typically, the hydrogen bonding is less than 5%, often much less than 5% as strong as a covalent bond. So a hydrogen bond isn't anything like the bonds we talked about last semester.

Or is it? A hydrogen bonding can be subtle. And it's the subject of continuing investigation. There's a Professor Johnson, a physical chemist in our department, is doing a lot of work on the structure of water. And Professor Jorgensen, a theoretician, has done a lot of work on water as well.

So if you have ions involved, so not neutral waters coming together, but the protonated water coming together with the water, then the question is what kind of hydrogen bonding are you going to get in this case? Now the heat involved in those things coming together, instead of 3.3 kilocalories per mole is 10 times greater. 32 kilocalories per mole. But, of course, there's an entropy price to pay when you pull these things together.

So the actual free energy change in bringing them together isn't nearly as favorable because you're localized in things that were free to move around before. So the free energy change at 300 kelvin is 8 kilocalories per mole in the gas phase.

But here's what's interesting. There's the structure of this H5O2+. What do you notice about it as compared to the water dimer?

Student: I think it's bonded.

Professor Michael McBride: Why do you think so, Lauren? Why do you think it's bonded?

Student: Because the lines are going both ways. It looks like it's sitting right in the middle of it.

Professor Michael McBride: Right. The hydrogen is right in the middle. It's not like the other one where the hydrogen was a normal bond distance from one oxygen and quite far from the other one. It's right in the middle. And in fact, if you do a calculation of how much energy it takes to move that hydrogen back and forth, you see that it's a single minimum. Not a double minimum where it could be part of the time on one, part of the time on the other and jump back and forth. It's a single minimum. The H+ is equally shared. So this case is real bonding. A single minimum in the middle. 

Now, how about if we do the anion example--H3O2–, hydroxide being hydrogen bonded by water. So what's the answer there? We know that the energy involved is rather similar to what it is for the H5O2--it's 27 kilocalories per mole in the gas phase. But the structure is rather different, or is it? Is the hydrogen halfway in between? What do you notice about this one compared to the cation? It a double minimum now, not a single minimum. So it's lower in energy when the hydrogen is on one end or the other. And a maximum, local maximum, in the middle.

However, hydrogen is very light. What does that have to do it? Go back to the beginning of last semester. 

Student: Tunneling?

Professor Michael McBride: Pardon me?

Student: Tunneling? 

Professor Michael McBride: Tunneling. If you look at the very lowest energy level for moving the hydrogen back and forth, you'll see something interesting. So this is a double minimum. The hydrogen is not symmetrically bound along this. But it can't have the minimum energy. And the barrier is very low, only 0.6 kilocalories per mole. So if you look at the lowest-energy vibrational wave function, it looks like that.

So even though it's a double minimum, the hydrogen spends most of its time in the middle. And, of course, the reason it can do that is that--the reason it can spend most of its time in the middle is because there is some bonding that lowers its energy there. So there can be hydrogen bonding. We saw it in H3O+ plus water. We see it in hydroxide plus water. Even though it's a double minimum in potential energy, the thing can be in the middle. So there is real bonding there. 

And in fact, in pure water without ions, if you have high pressure that forces the waters closer together, so the better part of a million atmospheres--and this can be done in a laboratory--the oxygens get closer together and closer together. What was a double minimum becomes a single minimum. And the structure of this Form 10 of ice--water is nothing if not versatile. There are more than 10 different structures of ice that have been observed.

But this particular high pressure one, as you can see from the structure, has hydrogens halfway between. So it's no longer molecular water. It's an entire crystal--the whole thing is one molecule with hydrogen bonds holding it together.

But this is not what we usually are talking about in terms of hydrogen bonds. Normal hydrogen bonds are just this polar interaction, which are favorable because hydrogen's so small and can get close to other things.

Chapter 4. Ionophores and Phase-Transfer Catalysis [00:28:09]

Now we're going to go on from this digression about water to talk about crown ethers, and what are called tailored ionophores. 

So here's a crown ether. You have to be imaginative to see that as a crown. But you can sort of imagine putting it on your head and have all the little things go up and around, and I suppose that's where the name came from. 

So this idea of having an oxygen and then a bridge of two carbons and another oxygen and a bridge, oxygen, bridge in a ring. Those are the crowns. So they have a name, this is 18-crown-6. So it's a crown alternating oxygen and two carbons, oxygen and two carbons. And it's six because there are six oxygens in it. And this was one of the earliest ones studied. And in fact, these crown ethers resulted in the Nobel Prize in Chemistry for the people who discovered and exploited them in 1987. 

Ionophore means an ion carrier--the Greek root phorous means to carry things. So these things are able to carry ions--let's see how that works.

So here's 18-crown-6 viewed from the top and from the side with potassium chloride. And you see the potassium is held inside. I'll pass this around so you can see it. Here's the ring of 6 oxygens connected by pairs of CH2 groups. It turned out that a golf ball is about the right scale to have the radius of a potassium ion on the scale of these models. And you see it fits right in the middle.

If we look at the 6 oxygens, they're not all in a plane, as you can see there. There are three of them up and three of them down. Does that remind you of anything? Megan?

Student: Cyclohexane. 

Professor Michael McBride: Just like cyclohexane--it goes up down, up down, up down, the chair cyclohexane. And the potassium sits right in the middle with three above and three below, and I'll pass that around so you can see it.

So potassium is just the right size to fit in there--fits right in the middle. If we look at the actual distances that are involved you can see these distances are all about 2.8 Å, which is just about exactly the sum of the ionic radii of potassium and the covalent radius of oxygen--2.7 that would be. Those numbers aren't really such great numbers, so that it really fits right. 

Now what if you use a bigger ion like cesium? So cesium is 1.7 Å, so now 3.1 instead of 2.7 for the oxygen-cesium distance.

Now you can see that because cesium's larger it can't get in the middle; it sits on top close to three of the oxygens, and further from the bottom three oxygens. There's the view from the back. Now here's the skeleton with the balls for the atoms of that. And you can see that the closer oxygens are at van der Waals distance, 3.1 Å. But the others are closer to 3.2 or 3.3 Å.

Now that distance is 4.8 Å. What happens if you would make an even smaller ion in the middle? You could try to shrink the ring down. But if you pull those two closer together in order to make all five of these closer to whatever's in the middle, then that one's going to get squeezed out. 

So here's what happens if you go to sodium. And now that distance is 4.0, roughly an angstrom closer in order to get these close to the sodium. Sodium to oxygen is 2.4. So that sixth one got squeezed out, but it came in from the bottom. So if we look at the distances, those are all in the neighborhood of 2.4, 2.5, one is 2.6, so just a little bit bigger than the van der Waals distance. And the one on the bottom is 2.4, so it squeezes underneath and is just happy. But it leaves space on the top, and this particular crystal structure had a water in it. So the water is up on top at 2.32 Å. 

If we get smaller still with lithium, now the lithium can't shrink the ring down anymore. It gets close to just two of them, 2.2 Å--notice that's 2.1 and 2.1, 2.07. But it's further from the others. And notice it sits right in the middle of the mean plane of these things.

But what happens with interactions with the far oxygens? This one has two waters. So the waters bind to the lithium, and then the protons of the water are hydrogen bonded to those far oxygens. So lithium increases its size by incorporating two waters with it, so it can fit in this thing. 

Now look at the equilibrium constant for binding the cations to this 18-crown-6. Notice that potassium is by far the biggest. So these are too big--we looked at cesium. Sodium and lithium were too small, but potassium is just right. So when this is 1 this is 60 times bigger. 

Now the equilibrium constant is, of course, the concentration of the complex divided by the concentrations of the metal ion and the ligand product. So it has units expressed this way, it's per molar. And if you look at the actual value, not just relative to one another, which was what's plotted first but the actual value, you see that it's over a million, this equilibrium constant. Now let's think of what that means. That one's 23,000.

So what this means is that when the ligand, down here, equals--when the ligand times K is 1, then you have 50% that this is 1:1. If you multiply this over here and that comes out to 1, then you have the same amount of metal that's uncomplexed as you have of metal that's complexed. So what that means is that it's half complexed when you have 1 µM of concentration of the ligand. So it holds on very, very tightly.

Now, bear in mind that there's something else that could complex with the potassium ion. The methanol solvent could do so. This is done in methanol solvent at room temperature. Now methanol is 0.79 g/mL. Its molecular weight is 32. So we can calculate what its concentration is. Its concentration is 25 M. So the ligand, at 1 µM, is able to beat out methanol at 25 M. So it's 29 million times stronger than methanol at binding to this by having all these six arranged so they grab it just right, it fits just right.

The energy involved is 13 kilocalories per mole of stabilization. Again, nothing compared to covalent bonds that we talked about. But it's much more favorable in the case of a potassium than it is in the case of sodium, just 5 kilocalories less. But notice what this does when you put this crown around the golf ball. It makes it much bigger. But what holds a salt together is the attraction between the cation and the anion, which depends very much on distance.

So if you make the ions bigger, as we mentioned in the last lecture, they won't interact so strongly. So it's not so stable to be in the solid salt form. It's easier to dissolve. So you can dissolve potassium permanganate in hydrocarbons if you use crown ether to put in there to make the potassium ever so much bigger. 

So this has worked out in lab in a way that's shown here. So you have some organic solvent, and you put something in it that you want to oxidize. So you need an oxidizing agent--ah, potassium permanganate. So we have an organic substance to be oxidized, and we put in potassium permanganate, and it just sits on the bottom of the flask. It doesn't get to where the organic molecules are, so there's no reaction. 

And if you put water in, that will dissolve the potassium permanganate, but it won't get the potassium permanganate where the organic stuff to be oxidized is. It'll just be two phases. A purple phase--a water phase that has the permanganate in it, and then the organic phase just floating above it. But if you put in the crown ether, then the permanganate can dissolve in the organic solvent and can do the trick.

So by making the cation large, 18-crown-6 destabilizes a solid, or aqueous permanganate, allowing the salt to dissolve in hydrocarbons. And people call this purple benzene. I actually have never done this and never seen purple benzene, but Professor Ziegler has done it, and he tells me that purple benzene prepared this way isn't purple at all, it's yellow. So that's just his one experiment. So I don't know that it's never purple, but that's what he told me. So maybe purple benzene looks sort of yellow, according to Professor Ziegler. That would just be because permanganate in an organic solvent has a different spectrum from permanganate in water solvent, which wouldn't be surprising. 

But it works. People are able to do oxidations this way. And it's called phase transfer catalysis, because the idea is to get a reagent from one phase into the other. You use the crown ether as a catalyst. It doesn't get consumed, it just helps the stuff get there and do its reaction in the other phase, so it's catalysis. And you don't have to use crown ethers for this, although they are used for it. You can get a similar effect by having other really big cations that make it possible for an anion to go into the organic phase. So tetraalkylammonium or tetraakylphosphonium salts.

The advantage of these things is that there are solvents that will dissolve both organic things and potassium permanganate, or other salt reagents. But those salts [correction: solvents], like dimethyl sulfoxide that will do that, are dangerous and they're expensive. So this is an example of green chemistry--to do chemistry in a more environmentally favorable way by using these ethers or simple salts as a way of getting the reagents to mix.

Now a further development on this is called cryptands, which you notice is like a crown ether, except that it has a nitrogen in it which is trivalent rather than oxygen being divalent. So you can get another bridge across. So you can get these crypts into which to hide things.

So here's an example of a cryptand. And Jean-Marie Lehn who did this work with Jean-Pierre Sauvage, got the Nobel Prize for this--shared the Nobel Prize for these crown ethers. So now you can put different numbers of these OCH2CH2 groups in here. So you can make different sizes for these rings. 

So here is an example of where l is 1, so you have 1 bridge that has two oxygens, and two bridges that have one oxygen, and I've made a model of that here. So you see there are two oxygens here on the top, but there's one on this way and one here. And two nitrogens also point in with their lone pairs. So that makes a pretty small cavity inside. I'll pass that one around to you.

You won't be surprised to hear then that when you look at the equilibrium constant in methanol for putting various salts in there, that lithium is the one that binds most strongly, the smallest of the ions. It has an ionic radius of a little less than three-quarters of an angstrom. So it fits, but sodium doesn't fit as well, potassium's bad. And these are whited out because, in fact, they were measured but they come way down off the bottom of this graph.

And now if you make one of the other bridges have two oxygens in it, so you increase the size of the hole a little bit, now it binds with sodium the best. And if you put the other bridge bigger it binds potassium the best. And if you put two carbons in one of the bridges, it binds rubidium the best. If you put two carbons in all the bridges it binds cesium best. And I put cesium-137 here because that's a by-product of isotope enrichment in nuclear industry, and it's a common pollutant that they want to get rid of.

So what they're trying to do now is to design cryptands that will readily bind to cesium, so you can take the earth or these million gallon barrels that have this waste stuff in them. Put something in that will specifically get the cesium, which is the worst actor among these radioactive by-products. That'll take them, put it in another phase, and then make it reversible so that you can take the cryptand off and use it again. So there's a lot of work being done on that kind of thing now. 

But this also could be used in biology. For example, there's a bacterial antibiotic called nonactin, which you see is a ring that has a lot of oxygens in it. Now that ring is quite a big one, but you can curl it around like this, as I show here--bend those ends up. And then the chain looks like the seams on a tennis ball. But you have all these oxygens that can then be directed toward the middle.

So here's a picture of that wrapped around a cation. Now it's not easy to see the three dimensions here, but if you let it rotate you can see what the shape is of this thing that's wrapped completely around the potassium ion with 8 oxygens coordinated to the potassium. And there you can see the seams on a tennis ball, which trace this chain. And the equilibrium constant in methanol for holding on to sodium is 512. But for potassium, the one that's shown there, it's 31,000, so it's very selective. It's even tighter with ammonium in the middle.

So how do these things work in biology? I don't know, but the general idea is that they can make ions dissolve in organic things. So the potential in your cells has to do with pumping sodium and potassium ions back and forth and creating a gradient. And if you suddenly put something in that'll say to potassium, you can go wherever you want. That destroys this and it's bad for living things. So anyhow, biology has a number of things that can do this kind of trick with ions as well.

Chapter 5. Energetics of Gas-Phase Heterolysis [00:45:05]

Now the next topic is the importance of solvent for ionic reactions. And I'm going to start this since we started a little bit late. But we'll complete it next time. 

So we're interested in the reaction, a water molecule surrounded by water, in water solvent, dissociating to give H+ and OH–. And we want to know what energy's going to be involved in that. What will the equilibrium constant be? And I think you people know the equilibrium constant from your AP chemistry. But we're going to look at it a different way as to what's involved, what role the solvent plays in this. 

So the way we're going to do it is to look at it without solvent by going into the gas phase. So if we pump in 6.3 kilocalories per mole, we can make the water that's in water become water in the gas phase. And now we know we can look up in Ellison's table a bond dissociation energy for water and find that it's going to take 120 kilocalories per mole to get it apart to H atom and OH radical. 

And then if we transfer an electron from one to the other, that's then going to get us the species we're interested in, but in the gas phase. That's 392 kilocalories per mole in toto to do those three things. And remember the Coulomb energy is the gorilla in this particular scheme. It's that 332 kilocalories divided by angstroms. So that's the big one is that electron transfer and getting the ions apart. 

So the equilibrium constant in the gas phase is 10–290--not a very favorable process. But we know that it can happen in water. So the solvent has a lot to do with it, and we'll talk about that next time. 

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