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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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In the Pipeline: Don't miss Derek Lowe's excellent commentary on drug discovery and the pharma industry in general at In the Pipeline

In the Pipeline

« Costs and Benefits, Risks and Rewards | Main | How Long Can This Go On? »

July 26, 2004

How Not to Do It: Sulfenyl Chlorides

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Posted by Derek

Someone came to my lab today to borrow some thiophenol, a request that made me think of something that happened in my first summer of undergraduate research - twenty-two, gulp, years ago. Now, thiophenol is not known as a great inducer of nostalgia. Like the other small-molecule sulfur compounds, it reeks without letup. It's a major part of the smell of burning rubber, so if you can imagine that concentrated and put into a bottle, you've got a pretty good idea. It's distinctive.

I was using this cologne as a starting material, reacting it with sulfuryl chloride, which is another reagent that no one is going to dab behind their ears. It's a reactive chlorinating agent and a fairly strong oxidizer, and it'll make you shake your head and snort if you come across its fumes, for sure. The two of them together make an eyebrow-raising mixture - I was exiled to a lab at the other end of the building while I ran this one, just because of the potential smells.

Heating this brew gives you phenylsulfenyl chloride, a red oil which combines the foul properties of its parent compounds. You distill the stuff out of the reaction, cap it up, and store it in the cold. I think it's too reactive to be an article of commerce; you have to make it fresh. And make it fresh I did, even though everything around me smelled as it had been dead for weeks. In the freezer with the stuff for the weekend (we didn't work grad-school hours at my all-undergraduate school, not even in these summer research programs.)

Monday morning I went down and picked up the flask. Hmmm. . .no longer a red oil. Odd. The stuff had changed to a pale yellow solid, which didn't seem right. I wasn't sure of the compound's freezing point, but the color change alone made me wonder. I stood there puzzled for a minute or so with the tightly stoppered flask in my hand, holding it up to see what I could make of the stuff. And then, with a loud gunshot bang, the top of the flask exploded in my hand.

I jumped straight up in the air, flinging away the lower part of the flask that I was still holding. I came down on the balls of my feet, in a fight-or-flight stance, looking around wildly. I didn't feel as if I'd been injured, but I'd never had anything blow up while I was holding it, either, so who kenw? After a second or so I looked down to see if I was OK. And weirdly enough, I was. I can't imagine how I managed not to pick up some glass shrapnel, at least - perhaps even at that early point in my career I had enough sense not to point the neck of a round-bottom flask toward me. My hand was fine - I kept flexing my fingers in wonderment. After a minute of two of stalking around the room, shaking and gibbering, I started looking around to see what had become of the chemical.

I found it about ten feet away, a lens-shaped piece of light yellow stuff, molded smooth by the inside of the flask. Whatever it was, it wasn't melting again, and it sure wasn't phenylsulfenyl chloride. We figured out what it was pretty quickly, but I think I'll leave its identity as an exercise for the technical portion of my readership - guesses to go in the comments below. If you get it right, you'll know why it blew up, too!

Comments (5) + TrackBacks (0) | Category: How Not to Do It


COMMENTS

1. brian on July 27, 2004 7:49 AM writes...

I'll guess diphenylsulphane(?) [two phenyl rings with a pair of sulphurs in between] and chlorine gas.

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2. Drew on July 27, 2004 1:03 PM writes...


Perhaps diphenyl(thio?)ketene?

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3. Stewart on July 27, 2004 9:13 PM writes...

I think Brian's close, but I'll guess it's the polymer [-C6H4-S-]n and chlorine gas.

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4. Derek Lowe on July 27, 2004 9:27 PM writes...

Brian nailed it. It turned out to be diphenyldisulfide (nice NMR, melting point and all), and the chlorine gas is what blew the flask.

Now, why the stuff decided to disproportionate on storage, under nitrogen, in the dark, in a freezer, is another question. Maybe it's an autocatalytic process - could be why no one sells the stuff commercially (at least, there's nothing in ACD or the like.)

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5. hound on August 1, 2004 12:21 PM writes...

Sounds like one of the reagents you could buy in dilute ether solution under nitrogen, perhaps you had traces of water or some metals in your synthesis, despite the nitrogen. Well it would take a lot of LC/MS work to find exactly what is going on, probably not worth it.
just found your blog and it looks really good! i'll be back

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