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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: derekb.lowe@gmail.com Twitter: Dereklowe

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August 31, 2007

Here It Goes

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Posted by Derek

Stability is a relative concept in chemistry. In the lab, we tend to use the term a bit loosely, and we mix up with “reactivity”. But those are two axes of an x-y graph, and there are chemicals in all four quadrants. Stable and non-reactive? Sure, for whatever value of “non-reactive” you choose. Stable and reactive: how about acid chlorides? You can keep many of them happily for years away from water, amines, etc., but open the flask and they’ll be ready to party. Unstable and non-reactive? An odd category, but I’d say that something like a polyazide or polynitro compound would fit. It doesn’t do much with other chemicals; it just falls apart on its own, and how. And unstable and reactive? Oh, yeah, we have those, all right.

In the lab, there’s a large middle ground of things that sort of gradually deteriorate on you, but not so quickly as to be a nuisance. Solutions that used to be clear pick up a yellowish cast, crystals get cloudy. This is the sort of stability that people are used to seeing with newsprint paper and household chemicals like bleach – they’re good for a while, but you can’t expect them to hang around forever. In research, you deal with this by either buying new stuff (the industrial way!) or re-purifying the old bottle by distilling or recrystallizing it (the academic way, by necessity).

After these compounds, though, you come to the ones that can give you trouble. There are a lot of compounds that are only stable on a time scale of days, hours, or minutes, and you’ve got to keep an eye on these guys. Often the rate of decomposition is very dependent on how pure the stuff was at the beginning. Trace amounts of water, oxygen, or other such rare substances can start one of these down the slope.

The dangerous ones are the compounds whose decay begets their own decay. These will run away on you, and if there’s enough compound in the flask where heat transfer is a problem, the process can turn violent. At this point, we’re shading over from “troublesome decomposition” to “explosive hazard”. Things like this are best kept as cold as possible, and in dilute solution. Concentrating them or warming them is a deliberately provocative act for which payment will be due.

Even without explosions, this sort of thing can be alarming. I’ve heard of intermediates that were so lively that initial clearish substances in a round bottom flask turned brown and began to fume as the person walked down the hall holding the stuff. Generally, that only happens once, the first time you make one of these beasts. After that, you take appropriate precautions (like having the next reaction step set up right next to this one, ready to go). Or, of course, you just decide that you can live without that one, and never make the darn stuff again.

Comments (9) + TrackBacks (0) | Category: Life in the Drug Labs


COMMENTS

1. Curious Wavefunction on August 31, 2007 9:44 AM writes...

You have touched upon a very important point; a lot of times people (including myself) make the mistake of thinking of stability in absolute terms by saying things like "Oh, that looks pretty stable". For starters, we always have to distinguish between thermodynamic and kinetic stability. A thermodynamically highly stable compound may not be easily formed because the energy barrier needed for it to form is too high.

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2. walkerma on August 31, 2007 11:33 AM writes...

Another interesting case are chemicals that become more reactive the more pure they are. A bottle of commercial benzaldehyde will last for years on the shelf. But if your distill it, a large portion of it will oxidize in a few days to benzoic acid. I am not aware of any stabalizers being included in commercial benzaldehyde so there mus be something else going on.

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3. MolecularGeek on August 31, 2007 11:53 AM writes...

Walkerma,
Maybe the commercial benzaldehyde is in a container that has had the oxygen displaced and is stored under dry nitrogen? If there's nothing to reduce, it's hard to oxidize anything. But once you open it and expose it to oxygen...

MG

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4. milkshake on August 31, 2007 6:57 PM writes...

Benzaldehyde oxidation is autocatalytic radical process. Once you get it going it will keep going as long as some oxygen is around.

I was making a triflate from acetonide of a partially protected polyol. The triflate was reasonably stable to be columned (quickly, in DCM) but working with the neat purified stuff was fraught with danger - as soon as a little triflic acid was generated it promptly turned into a black tar. The first time I made it, as soon as I put the clear oil on highvac, I saw brown-black plague spreading from few points all over the flask until the whole amount turned into foamy asphalt. I solved the decomposition problem by adding a drop of lutidine to all my column fractions and skipping the evaporation, I used the DCM solution straight as it came off the column, for the next step

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5. Michelle on September 1, 2007 8:14 AM writes...

A delightful vivid view from the bench! I'll share this one with my general chem students in the spring, who wonder why we make them think so much about stability and reactivity. (We use the ACS text, which seriously focuses on this as an underlying unifying concept.)

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6. Great Molecular Crapshoot on September 1, 2007 3:28 PM writes...

And then there's the strange case of cyclopropane. Thermolysis of this highly strained cycloalkane is more favorable than the less strained cyclohexane. But it's more difficult to yank a hydrogen atom off cyclopropane. Counterintuitive... maybe? But think what CYPs need to do in order to metabolise alkyl groups.

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7. Liquidcarbon on September 2, 2007 3:51 PM writes...

Derek, is there really so many non-chemists in your audience so that you refrain from mentioning organometallics that tear hydrogens from solvents (even from carbon tet, yep) and oxygen from the glass?

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8. milkshake on September 3, 2007 7:26 AM writes...

In Soviet Russia, carbon tet tears hydrogens from you!

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9. Jonadab the Unsightly One on September 26, 2007 7:52 AM writes...

Liquidcarbon: yes, there are non-chemists in the audience. I'm a network administrator who majored in math, for instance, and In the Pipeline is one of my favorite blogs to read.

With that said, I believe if I were a chemist, reacting with glass would be a really good way for a substance to get on my "avoid" list. I think that's a good deal scarrier than merely reacting with air. How would you reliably contain such things? Quartz bottles and flasks? Yow. Why not just go the whole way and work with antimatter.

Reacting with solvents, on the other hand, seems relatively normal, although I'm not sure precisely how something would take hydrogen out of carbon tetrachloride, given that unless I'm very confused about its chemical structure it doesn't normally have any hydrogen to take. Perhaps someone who works in or at least majored in chemistry could explain that one to me.

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