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DBL%20Hendrix%20small.png College chemistry, 1983

Derek Lowe The 2002 Model

Dbl%20new%20portrait%20B%26W.png 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: Twitter: Dereklowe

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November 25, 2002

A Chemical Wish List

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Posted by Derek

After some years in the drug discovery business, I have a few modest requests to make. To start with, I think that the existing periodic table is too limited, with a number of elements that clearly were left out. I have a shopping list of what's needed to complete the set, and here's number one:

We need something electron-donating that's no bigger than a fluorine. Medicinal chemists, when they hear that, usually say "Hmmm. . .yeah!" That is, when they're not telling me I'm nuts. For those outside the field, I'll explain: when you want to substitute something for a hydrogen on a carbon atom, fluorine's usually your boy, because it's not that much larger, and you get a lot of other effects as well.

Untold numbers of useful things have been made this way. The carbon-fluorine bond is a lot stronger than the C-H bond it replaced, and the fluorine has some odd properties all its own. (You go from polyethylene to teflon when you make that substitution, to pick one example - I wouldn't recommend buying a polyethylene-coated frying pan unless you want your next batch of scrambled eggs to take on a completely new dimension.) The lower reactivity means that fluorocarbons generally don't burn worth a hoot, and it lets med-chem folks like me stick fluorine atoms all over our structures to keep the liver from ripping them to shreds. Not even the liver seems to be able to tear up a fluorocarbon.

(That low reactivity is also why freon-type compounds were assumed to be completely inert for so long. As it turns out, much of the ozone-depetion reaction they set off is caused by some other carbon-chlorine bonds breaking apart in the hard UV light of the upper atmosphere. You really have to beat the heck out of a compound to break C-F bonds, but C-Cl will turn on you.)

Of course, all these interesting properties come at a cost. Fluorination reagents tend to be a bit exotic and expensive, for one thing. That's because they're mostly a way to put a leash on elemental fluorine itself, and a strong leash it had better be, because it's about as nasty as it gets. Plain fluorine will start in on basically every compound there is, and plenty more: it'll make steel wool, for example, burst into flame, so you don't even want to think about what it'll do to your hand. Whatever it does, it does with a huge release of energy; it's just too happy to pitch in. Ask Berthelot, the French chemist who tried to make fluorocarbon compounds in the 1800s by using the same sorts of reactions that worked for chlorine. He blew himself up several times in a row; how he survived is a real mystery. Most of the early fluorine investigators poisoned themselves with the stuff to some degree.

You also pay with some physical properties you might not want. Fluorine owes its weirdness to the way that it crazily pulls electronic charge off of any atom it's attached to. (You can practically hear the electrons being sucked over across the bonds - well I can, anyway, but maybe I've been doing this stuff too long.) That's all fine, but what about the times when you don't want all the electron density piling up on one end of your molecule? Wouldn't it be good to have something that pours charge back into the system, but keeps that handy fluorine size?

Well, I can tell you that if we had this element available, there would soon be some pretty puzzled enzymes and receptor binding sites out there. Medicinal chemists would be putting it on everything and making compounds these poor proteins never even dreamed of. We spend a lot of time adjusting polar groups on our molecules to make them tuck into binding pockets more tightly - something like this would give us work to do for many years. For starters, I'd take all the compounds that got worse when I fluorinated them, and put this stuff on instead. Couldn't hurt.

And there's nothing but the laws of physics to keep us from having it. Maybe some other universe is built so that they have my element in it. If so, there are probably some fellow chemists sitting around wishing that they had something like fluorine. . .

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