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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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October 29, 2006

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Posted by Derek

In his book The Periodic Table, Primo Levi mentioned in passing that "chlorides are rabble". That struck me as very well put, and is proof enough (should anyone need one) that Levi had a real feeling for his chemistry. The reason that comes off so well is that when you look through a chemical catalog, if there's only one salt of a given element available, it's almost always the chloride. They're common, in every sense of the word.

The kinship of the positive ions, the elements themselves, are well known. That's how Mendeleev worked out the periodic table, and generations of chemistry students are taught about the similarities among its columns. It's all true, of course, but there are subtle kinships of the counterions, too, a faint Y axis to the strong X of the elements.

Most of the chlorides are quite boring - white powders, almost invariably. The more chromatic elements still manage to do something for you: nickel chloride, for example, is a vivid green (copper less so), and chromium (III) chloride is a striking metallic-flake purple. But if you can't get colorful with elements like those, your counterion is a total loss, anyway.

Fluorides are almost never colorful, but they have a tough nature about them, reflecting their ultimate-hard-anion character. Iodides are the other end of the scale - that's such a big, fluffy ion that it hardly seems bound at all sometimes. Even light is enough to mess with it, and it's a rare iodide that doesn't have a warning on its label to keep it out of the sun, for fear of it turning brown in a death-tan of oxidation to free iodine.

Sulfates are nearly as boring as chlorides, but with a bit more character to them. Nitrates (similar salts from a strong acid) have a much different feel to them, since when you're working with them you can never quite get the thought of explosions out of your mind. It's not completely accurate, but it's still true that you could mix potassium sulfate with sulfur and charcoal forever and never discover gunpowder. The word "nitrate" itself has a menacing sound that it'll never lose.

If you want real problems, though you have to turn to even more loosely bound, oxygen-rich things like bromates, iodates, and (above all) perchlorates.
That's about as bad as it gets inside the confines of inorganic chemistry - to get crazier, you have to trespass into organicky things like azides. Most of the organic counterions, though, are carboxylate salts, which are relentlessly similar to each other. No explosions here - if there's one salt of a element that's guaranteed to be more yawn-inducing than its chloride, it must be the acetate.

These are all classical ions, known for centuries. The fluorides are probably the most nouveau of the lot, since even though some of them occur naturally, most of them had to wait until the industrial development of the element later on in the 19th century. But that led in the 20th to all sorts of odd creatures that (so far as I know) are never found in natural minerals at all. The higher fluorides, things like tetrafluroborate and hexafluorophosphate, have only human fingerprints on them. When you work with those salts, you've thrown your lot in with the synthetic, the man-made, the new and improved. Even weirder ones are surely on the way.


Comments (14) + TrackBacks (0) | Category: Life in the Drug Labs


COMMENTS

1. steve s on October 29, 2006 10:54 PM writes...

Fluorides are almost never colorful, but they have a tough nature about them, reflecting their ultimate-hard-anion character. Iodides are the other end of the scale - that's such a big, fluffy ion that it hardly seems bound at all sometimes. Even light is enough to mess with it, and it's a rare iodide that doesn't have a warning on its label to keep it out of the sun, for fear of it turning brown in a death-tan of oxidation to free iodine.

I have a BA in physics, and I know little about chemistry. Nevertheless, you've just illuminated a little part of the world for me. Thanks.

Permalink to Comment

2. kiwi on October 29, 2006 11:45 PM writes...

triflate salts are another nice example of manmade fluoro anions. useful too.

Permalink to Comment

3. Jeff Bonwick on October 30, 2006 5:13 AM writes...

So, Derek... if the cure for cancer turns out to be a heavy-metal salt of octaazocubane, may we safely assume that you won't be the guy flying to Sweden?

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4. Tom Womack on October 30, 2006 5:25 AM writes...

Nature's firier synthesis vats turn out to produce the most peculiar things:

http://www.minsocam.org/ammin/AM79/AM79_381.pdf

reports the discovery of natural NH4BF4 deposited around the hotter fumaroles of Fossa on the island of Vulcano; Na2SiF6, too.

avogadrite (K,Cs)BF4 and ferruccite (NaBF4) are both found around fumaroles on Vesuvius.

A lot of group-1 haxafluoroaluminates show up in Greenland among cryolite (Na3AlF6) deposits, of course. Towards the end of http://www.mindat.org/strunz.php?a=3&b=B&c=03 you find some truly peculiar polyfluoroaluminates like Bøgvadite and Jørgensenite, and the 'related compounds' even suggests (though unconfirmed) a polychloroaluminate.

www.mindat.org has a remarkably comprehensive catalogue of naturally-occurring substances (though sadly lacking a search-by-element option); is there some obvious chemical reason why hexafluorophosphates don't show up while tetrafluoroborates do?

Permalink to Comment

5. Supercritical on October 30, 2006 6:34 AM writes...

The interest in ionic liquids (one of your favourite areas) is also driving forward the design of new and improved ions.

Permalink to Comment

6. Supercritical on October 30, 2006 8:36 AM writes...

The scarcity of PF6 anions could be because they are not stable in the presence of moisture, they release HF. This is also the case with BF4 anions, although i think they might be slightly more stable.

Permalink to Comment

7. Derek Lowe on October 30, 2006 9:17 AM writes...

Well, raise my rent. I had no idea that tetrafluoroborates and the like were found in nature - those must have been some mighty ugly conditions.

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8. Tom Womack on October 30, 2006 11:59 AM writes...

I'm always surprised at the stuff that comes out of volcanoes; Kudriary Volcano, Iturup Island, Russia had a fumarole around which crystallised pretty pure rhenium disulphide.

This is highish pressure and 700 centigrade, nastier than man industrial processes; I've no good idea where either the boron or the fluorine's coming from, they're hardly common elements in rock.

Permalink to Comment

9. Derek Lowe on October 30, 2006 1:02 PM writes...

True, but they're a heck of a lot more common than rhenium. Makes you wonder what kinds of compounds will show up in places like Olympus Mons on Mars. . .

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10. Amos Langely on October 30, 2006 2:11 PM writes...

What a bad rap for common salt, NaCl...

Common salt is considered by most authorities as an essential ingredient of our food. Salt was used not only as a food, but as an antiseptic in medicine. Newborn babes were bathed and salted, a custom still prevailing. The Arabs of the desert consider it so necessary, that in the absence of salt they bath their infants in camels' urine. Elisha is said to have healed the waters of Jericho by casting a cruse of salt into the spring. Abimelech sowed the ruins of Shechem with salt to prevent a new city from arising in its place.

Salt is emblematic of loyalty and friendship. A custom of pledging friendship or confirming a compact by eating food containing salt is still retained among Arabic-speaking people. The Arabic word for “salt” and for a “compact” or “treaty” is the same. Once an Arab has received in his tent even his worst enemy and has eaten salt (food) with him, he is bound to protect his guest as long as he remains.

Permalink to Comment

11. Chrispy on October 30, 2006 3:43 PM writes...


I recall sometime ago a colleage was asked to make a gram of a compound. It was a multistep synthesis and he came out with 3/4 of a gram as the chloride salt. So he converted it to the tossic acid salt and was good to go!

Permalink to Comment

12. Matt on October 30, 2006 3:56 PM writes...

Relativily speaking, perchlorate is quite stable. Its tetrahedrally symmetric, has no lone pairs and Cl is maximally oxidized... there is not much room for shenanigans. The real fun starts when you shift down to chlorate.. you lose the symmetry, there is a lone pair and Cl can reodx both ways. Hence perchloric acid is readily available and chloric acid is a curiosity.

Permalink to Comment

13. Monte Davis on October 31, 2006 10:25 AM writes...

DL: "I had no idea that tetrafluoroborates and the like were found in nature - those must have been some mighty ugly conditions."

And of course, our ambient conditions -- oxygen and water and UV -- are pretty ugly from a disinterested PoV. Per your remark on iodides, lots of compounds that aren't especially hard to form are rare in [this patch of] nature because they can't survive double-teaming by a killer solvent and killer oxidizer.

Permalink to Comment

14. eugene on November 4, 2006 5:48 PM writes...

Well, nature can't make the B(ArF)4 anion though. That one's gotta be synthetic.

Permalink to Comment

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