About this Author
College chemistry, 1983
The 2002 Model
After 10 years of blogging. . .
Derek Lowe, an Arkansan by birth, got his BA from Hendrix College and his PhD in organic chemistry from Duke before spending time in Germany on a Humboldt Fellowship on his post-doc. He's worked for several major pharmaceutical companies since 1989 on drug discovery projects against schizophrenia, Alzheimer's, diabetes, osteoporosis and other diseases.
To contact Derek email him directly: email@example.com
May 27, 2009
There are a lot of ways to think about the chemical reagents that we have stirring around in our flasks. But one of the basic ones, and one of the most useful, divides them into classes according to whether they’re in solution or not.
When things are in solution, they may act funny, but at least everything’s starting out on the same footing. If all the components are dissolved (and if everything’s stirring the way it should), then they all have the chance to find each other and do their respective things. But if some reagent is still a solid in there (powder, chips, what have you), that takes you into the nonintuitive world of surface chemistry.
This actually happens quite a bit. Plenty of standard organic reactions involve insoluble things where the chemistry takes place on the surface. There’s formation of a Grignard reagent from magnesium turnings, deprotonation with powdered sodium hydride, hydrogenation over palladium-on-charcoal – these are all classics. And I'm not even mentioning the surface-driven industrial scale catalyst systems today, which is unfair of me, since the economies of the entire industrialized world depend on them. But in all cases, the real details at the molecular level of these reactions are not easy to work out.
People are still arguing, for example, over just how catalytic hydrogenation works on the metal surface, although the general details of the mechanism are known. That one’s complicated by not just being the plain metal, but a weirdo solution of hydrogen in the metal lattice. There’s no dispute, though, that the reaction is taking place on the surface of the metal, and that the higher the surface area the better off you are.
That’s one big variable right there: surface area. Finely divided substances are very different players in these systems, and many chemists find (early in their lab careers) that they’ve unwittingly bought front-row seats for a demonstration of just how different they can be. Finely divided powders have a lot of surface area in them, and if that’s a rate-limiting factor, you can find yourself with something that’s easily a hundred times more reactive just by picking up a different bottle of what appears (at first glance) to be the same substance. I once saw someone substitute lithium powder for lithium sand in a prep without thinking about this issue, and not so much later, I got to see the same guy clean the inside of his fume hood out with a scrub brush.
But there’s more than just surface area affecting some of these reactions. Grignard formation, for example, seems to take place (at least initially) in fresh breaks or cracks on the magnesium surface. That exposes metal that hasn’t had a chance to become coated with anything (like a layer of magnesium hydroxide), and (zooming in) it also may reveal individual reactive magnesium atoms, left out on the edge and insufficiently surrounded by their teammates. Once these react and fly off into solution, the ones around them become exposed, and so on, and the oxidized layers become undermined and flake off. The standard Grignard-initiation tricks are all designed to speed this process along. A drop of iodine will react quickly with any magnesium points or edges, exposing still more fresh rough surface, as will reaching down under the solution and breaking the turnings with a spatula (or, alternately, grinding them with a heavy stir bar).
These days, what’s really complicating things is the ability to generate (and characterize) nano-sized particles. At some point, these things can stop behaving like tiny bits of the bulk substance (which can be enough of a difference in itself, as mentioned above), and start acting like completely new beasts. And the really nano-sized stuff has a better chance of actually being in solution – but that brings on various headache-inducing thoughts about what “being in solution” means on this scale. If you have clumps of (say) palladium a few dozen atoms wide, which manage to be solvated enough to float around, is that a heterogeneous reaction or a homogeneous one? At that size, is that a "surface", or not (and is the reaction really taking place on it?) What if the nanoparticles are immobilized on a solid support - do they stay and react there, or is the reaction driven by the few that escape? (That effect has been noted in the Heck reaction, among others).
We need to understand these things better than we do - there are surely a lot of very useful things that could be done if we had better control over catalysis and surface chemistry. It's going to keep a lot of people occupied for a very long time.
+ TrackBacks (0) | Category: Inorganic Chemistry | Life in the Drug Labs
September 16, 2008
I’ve neglected to note the death of Neil Bartlett, famous for showing that the noble gases would in fact form chemical bonds. This work was a real triumph, since the great majority of scientific opinion at the time was that such compounds were impossible. Bartlett, though, formed a rather startling compound while working on the platinum fluorides, which he realized was actually a salt of dioxygen. The idea that oxygen would be oxidized to a cation in an isolable salt was weird enough at the time, and Bartlett realized that if this could happen, then the same system should be able to oxidize xenon.
And so it did. It’s difficult to convey how much nerve it takes to do experiments like this. I don’t mean the dangers of working with such reactive fluorine compounds, although that’s certainly not to be ignored. (Bartlett spent much of his career working in this area, and only a skilled experimentalist could do that and remain in one piece). No, it’s actually very hard to get out there on the edge of what’s known and do things as crazy as making salts of oxygen and fluorides of noble gases, Consider that if you’d lined up a hundred high-ranking chemists to vet these experiments beforehand, most of them would have pursed their lips and said “Are you sure that you’re not just wasting your time on this stuff?” It takes nerve, and not everyone has it – but Bartlett did, and he had the brains and the skills to go along with it. You need all three.
There’s a good appreciation of him in Nature, which points out – to my mind, absolutely correctly – that he should have won the Nobel Prize for this work. In fact, I thought he had for a long time, and only a few years ago realized that I had that wrong. (I may have been reinforced in my opinion by a statement in Primo Levi’s The Periodic Table). I think that if you polled chemists as a group, you’d find that a majority would be under the same impression – and if that’s not a sign of the highest-level work, having everyone surprised that you never got a Nobel, then I don’t know what is.
+ TrackBacks (0) | Category: Inorganic Chemistry | Who Discovers and Why
May 15, 2008
I was running a copper-catalyzed coupling reaction the other day when my summer intern asked me how it worked. I showed her the mechanism that the authors of the paper had proposed, but pointed out that it was mostly hand-waving. The general features are probably more or less right: the copper iodide presumably does form some kind of soluble complex with the amino acid that’s needed in the reaction mix, and that may well form some sort of complex with the aryl halide, which opens up the ring to nucleophilic substitution, etc. If this were an exam, I’d give full points for that one.
But a lot of these couplings are, as I pointed out to her, very hazily worked out. The Ullman reaction, in various forms, has been with us for many decades, and there are more variations on it than you can count. If it always worked reasonably well, or if people had any strong ideas about how it did so, the literature on it wouldn’t be in the shaggy shape it is. Copper chemistry in particular has been (simultaneously) a very useful area for people to discover new reactions, and a horrible trackless swamp for people trying to explain how they work.
All you have to do is look at the vicious exchanges between Bruce Lipschutz and Steve Bertz during the 1990s about whether such as thing as a “higher-order cuprate” exists. I have absolutely no intention of reconstructing this argument; I would have to be paid at a spectacular hourly rate to even attempt it. It's enough to say that the arguments raged, in an increasingly personal manner, about what state the copper metal was in, what ligands coordinated to it, and what the active form of these reagents might be (as opposed to what the bulk of the mixture was at any given time). It culminated in what must be one of the most direct titles for a scientific paper I've ever seen: It's on lithium! An answer to the recent communication which asked the question: 'if the cyano ligand is not on copper, then where is it?'. That's in Chemical Communications 7, 815 (1996), if you're interested (here's the PDF for subscribers). Bertz continued to shell Lipshutz's position past the time when any fire was being returned, as far as I can tell, and continues to work in the area. Lipshutz, for his part, hasn't published on the higher-order cuprates in some time (being no doubt heartily sick of the whole topic), but has kept up a steady stream of work on new reactions involving copper, nickel, and other metals.
So if well-qualified researchers, brimming with grad students, postdocs, and grant money, can argue for years about copper mechanisms, I'm going to stay out of it. As time goes on, I'm increasingly indifferent to reaction mechanisms, anyway. I want to get product out the other end of the reaction. And while there are times when knowing the mechanism can help reach that goal, those times do not occur as frequently as you might hope.
+ TrackBacks (0) | Category: Chemical News | Inorganic Chemistry | Life in the Drug Labs
April 28, 2008
Dr. Warfield Teague is retiring this year, which makes me feel old. He was one of the professors who helped make me what I am today – in his case, partly by keeping me out of his chosen field of inorganic chemistry. It was a good move on his part; I’d surely have blown something up good and thoroughly when I got to grad school, such are the opportunities in that area.
Unfortunately for both him and for me, his Advanced Inorganic course ended up scheduled for 7:40 AM back in early 1983. I started out my college career with a barrage of classes at that hour, and made every one of them. My sophomore year, I only skipped one class, and I waited for the lightning bolt to descend even for that one. But my junior year I had a professor or two whose lectures could be safely (even profitably) missed, and I began to get in the habit.
Teague wasn’t in that category, though. His lectures were fine; it’s just that they took place so early in the morning. My roommate David and I, both chemistry majors, found it harder and harder to summon the activation energy needed to make it out of the thermodynamic sinks of our beds. Dr. Teague’s threat to come over and teach the class in our dorm room didn’t quite do the trick (while lying there in bed, actually, the idea had a certain appeal). But his threat to start giving top-of-the-morning quizzes did. I showed up, and kept showing up. First year of grad school, now that’s where I started slacking off in my classes in earnest. But not all of the professors I had that year could communicate the facts of their specialty as well as Dr. Teague could for his.
The lab part of the course, that I would have shown up at 6 AM for. I don’t know how he’s done it in recent years, but 25 years ago (not possible, that), we could do pretty much any lab procedure that Dr. Teague would sign off on. There was a requirement that we do at least one low-temperature one, one high-temperature one, one metal complex, and so on. So the dozen or so of us in the class would root around through Inorganic Syntheses or the like, looking for interesting stuff. And there’s plenty of it in there, let me tell you.
In my case, the most memorable included the preparation of fluorosulfonic acid from scratch. Scratch means you start from concentrated hydrofluoric acid, a fine substance for the spirited undergraduate chemist to become familiar with. I can still hear the peculiar whine that solid KOH pellets make when you toss them into a plastic dish of the acid – they’ve a pretty short half-life in there, I can tell you. And I also made the magnesium analog of ferrocene – magnecene, I guess you’d call it – by one of those don’t-be-afraid-of-the-obvious routes: heat some magnesium turnings to about 600 C in a tube furnace, and pass fresh cyclopentadiene monomer vapors over them. Works great. And while you shouldn’t be afraid of the paper synthesis, red-hot magnesium metal is something else again.
While I was thus engaged, my classmates were setting off thermite reactions, making phosgene from carbon tetrachloride (chromium trioxide, five hundred degrees, nothing to it), and preparing titanium tetrachloride from the ground up. (I can’t recommend that particular prep – the liquid “tickle-four” comes out bright green from being around 1 molar in dissolved chlorine gas, so you’re going to want to redistill it, most likely). We learned a fair amount of inorganic chemistry, and more than a fair amount of lab technique. As evidence for that, we all survived.
Whether the latest generation of undergrads will get these kinds of experiences, I don't know. But I'm glad I did, and I'd like to thank Warfield Teague for providing them.
+ TrackBacks (0) | Category: Inorganic Chemistry