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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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November 16, 2004

Things I Won't Work With: Ozonides

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Posted by Derek

I've never done an ozone reaction myself. In fact, I haven't seen anyone else do an ozonolysis in years now, and I wonder if this reaction is passing into chemical history. These guys are hoping not.) Many chemistry departments have an electric gizmo to produce ozone in small quantities, and I get the impression that they're mostly gathering dust.

Ozone attacks a carbon-carbon double bond, initially making an ozonide, a hair-raising five-membered ring that has three oxygens in a row. That rearranges to a still-alarming one with two on one side, separated by carbons from the other. That falls apart on workup to two carbonyl compounds (or other things, depending on what you add to the reaction.) It's a very clean way to oxidize a double bond and make reactive handles out of its two ends.

But it tends to be something that's done on a small scale, because those ozonides are packed with energy and ready to hit the town. In general, we chemists shy away from compounds with lots of single bonds between the elements on the right-hand side of the periodic table. Those guys tend to have a lot of electron density on them, and bonding between them is a careful, arm's-length affair, sort of like porcupines mating. Two oxygens single-bonded make a peroxide, and those generally blow up. A small ring with more oxygens in it than carbons will almost invariably blow up if you try to concentrate it or handle it too briskly.

I'd do an ozonolysis if I needed to (although first I'd have to find our machine and see if it even works.) But you couldn't pay me to try to isolate the intermediate ozonides. But you can pay some people, like Prof. Pat Dussault, who was a post-doc down the hall from me when I was in graduate school. He's made a career out of oxygen-oxygen bonds, no small feat.

Comments (13) + TrackBacks (0) | Category: Things I Won't Work With


COMMENTS

1. Daniel Newby on November 17, 2004 4:27 AM writes...

"A small ring with more oxygens in it than carbons will almost invariably blow up if you try to concentrate it or handle it too briskly."

TATP—triacetone triperoxide—the National Molecule of Palestine. ;-) It's actually surprisingly stable for something that wants to be gas.

My organic chem prof in high school told us about her graduate work on an ozonide (if I'm remembering right—anyway it was something exuberantly reactive). She cooked up a nice sample and brought down the window on the hood to go do something else. That was enough. Fortunately the hood was explosion proof, but whoever put it in didn't think to make the wall explosion proof too. It's not every day you see people in the hallway looking out of your hood. The glassware was left in very small pieces.

Permalink to Comment

2. otey on November 17, 2004 9:34 AM writes...

Derek -

Anyone else having problems with the column of links on the right spilling into your text? I can only read your posts in the "print" format.

Thanks.

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3. The Novice Chemist on November 17, 2004 11:44 AM writes...


Au contraire, mon frere. The ozonizer at this midwestern private university is well and truly used throughout the department. That may have something to do with the lax (that's stretching the word) safety standards of academic chemistry.

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4. Derek Lowe on November 17, 2004 2:04 PM writes...

Probably also has to do with me being in industry - I shouldn't have generalized quite so broadly. We don't tend to use ozone because we know we're not going to scale it up, so it would be a major error to use it in a key step of your synthesis. It's not like all the things we use can be run on scale, but we try to stay away from things that will obviously have to be abandoned first thing.

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5. Industrial Chemist on November 17, 2004 2:29 PM writes...

Ozone actually scales quite nicely and there are many different ways to use it safely industrially. When used properly, I think it can be much safer environmentally than disposing of osmium, since one tends to not generate any excess O3 than one actually needs. It is however, before scaling up, very necessary to do the calorimetry to make sure that one understands the thermodynamics of your actual system and the resulting intermediates first. To otherwise is sheer foolishness, but in our case the calorimetry is part of our standard scale up practices for all reactions. It is very possible to design very well-behaved reactions that scale beautifully.

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6. silyld on November 18, 2004 8:04 AM writes...

Was visiting an academic teaching lab in the UK only yesterday where they are showing 3rd UG students how to do an ozonolysis reaction and they've been doing it there every term for the last 15 yrs i can remember !

They also then let them lose to play with it on their own too.

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7. Derek Lowe on November 18, 2004 8:58 AM writes...

On a small scale, ozone isn't a problem, and there's always a workup step in the same flask. The "thing I wouldn't work with" is the isolated ozonide, which is only done by big-time peroxide fans like Pat Dussault, etc.




Industrial Chemist (above) is right that some places run large-scale reactions, but not every company is set up to do this. At the places I've worked, the process people would have come after me with pitchforks if I'd sent them a synthesis with a key ozonolysis step.

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8. Mike on November 18, 2004 11:32 AM writes...

Here's a tip for ozonizer users, learned from a graduate school labmate: When the oxygen tank runs out, make sure that the tank you replace it with is not a hydrogen tank.

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9. Derek Lowe on November 18, 2004 2:50 PM writes...

Ay. That makes me want to run for cover just thinking about it. Gegen den Dummheit kaempfen Goetter selbst vergebens. . .

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10. Harry on November 18, 2004 6:50 PM writes...

As a corollary to Mike's comment- always make sure to disconnect the cylinders at the cylinder valve,and not at the first Swagelok fitting. This makes it all too easy to mix up gases - sometimes with ummmm interesting results.

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11. Giagan on November 19, 2004 10:27 AM writes...

In graduate school, I several times carried out a reductive ozonolysis on a carvone-derived cyclohexenone, following a literature procedure for a similar substrate. Once or twice, I did the reaction on 15 grams (90 mmol) of material. You will love the reductive part of this ozonolysis--I warmed the -78 degrees mixture to 0 degrees and bubbled nitrogen gas through, then added Pd/C and stirred vigorously under hydrogen gas! The reaction gave large quantities of an aldehyde carboxylic acid, and I am still alive to tell the story. You'll see the reaction in a soon-to-be-published synthesis of guanacastepenes A and E.

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12. Chris Hoess on November 19, 2004 10:45 AM writes...

Oddly enough, there was an article in Nature this August describing an ozonide that's in Phase I trials as an antimalarial.

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13. Derek Lowe on November 19, 2004 12:49 PM writes...

You know, you're right. I'd forgotten about that one. There's a well-studied natural product, artemisinin, with a trioxolane (secondary ozonide) ring that's responsible for its antimalarial activity. Simpler trioxolanes were studied, and it looks like one was indeed selected for clinical development. (The Nature paper is here.)




I'm amazed that that works in vivo, and I wish them luck. I wouldn't recommend banging on a jar of that drug with a hammer, though.

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