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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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March 28, 2004

Thing I Won't Work WIth (2): Nickel Carbonyl

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Posted by Derek

Synthetic organic chemists rely a lot on inorganic chemistry. We let metals do a lot of work for us, particularly when it's time to do the real arc-welding of carbon-carbon bond formation. I have a pretty typical synthetic background, and over the years I've used palladium, platinum, sodium, iron, copper, rhodium, aluminum, mercury, silver, manganese, lithium, titanium, chromium, cobalt, zinc, ruthenium, vanadium, tin, magnesium, cerium, potassium, and probably a few more that escape me right now. Never sit near a chemist and give him any excuse to rattle off a list of elements.

I've never used elemental nickle metal, but I have broken out some of its salts from time to time. I especially enjoy the vivid green of nickel chloride, whose solutions look for all the world like lime jello. Not that you'd want to substitute that in your favorite recipe: nickle salts are rather toxic, and are suspected carcinogens to boot. But I'd work with them all day long to avoid dealing with another nickel compound, its tetracarbonyl.

That's a complex of nickel with carbon monoxide. CO has a good amount of electron density left on its carbon, and it'll line up on a metal atom, slotting into its electron orbitals and making itself at home. You can find carbonyl complexes of all the transition metals, as far as I know. Many of them are liquids, which is rather disconcerting when you consider their metal heritage.

Nickel carbonyl is a liquid, but it can barely restrain itself from being a gas. It boils at 43 C, so it has a pretty substantial vapor pressure, and that's a real problem. Said vapor, as you'd imagine, is rather weighty. It's not one of your wafting-away-on-the-summer-zephyr sort of vapors; it's more like a sort of ghostly molasses. It's so heavy that you really can't rely on a standard laboratory fume hood to contain it, because that's not the sort of hazard they're built for. Depending on the air flow and the sash, the stuff can just ooze right out the front of the hood and pour out into the lab.

You don't want it there. Breathing it is most unwise, because those CO ligands are not stapled on very well. If they find another metal that appreciates them more, they'll bail out, and an excellent candidate is the iron in your hemoglobin. There go four equivalents of carbon monoxide into your blood cells, and there's only so long you can keep that up. And there's the nickle, too - alone, bereft, with only your proteins to complex to. Wonderful. Recall that the metal is toxic all on its own, and you've now dosed in the most bioavailable manner possible. If you make it through the carbon monoxide spike, you have long-term metal poisoning to deal with.

Even if the vapor doesn't get the chance to wander around poisoning you, it can amuse itself right in your fume hood. If it rolls across a hot surface, of which there are no shortage in most working hoods, then it can explode, leaving behind a vile haze of carcinogenic nickle soot. An exploding toxin with a high vapor pressure - I just don't know what else you could ask for in a laboratory reagent. No doubt it does many interesting and useful reactions. They can save 'em for me, because I'm not that desperate yet.

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