About this Author
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
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July 14, 2008
Posted by Derek
Cyanogen bromide is not a nice reagent. It’s not quite on my list of things that I refuse to use, but it’s definitely well up on the list of the ones I’d rather find an alternative to. The stuff is very toxic and very volatile, and reactive as can be.
But it’s not the worst thing in its family. A good candidate for that would be cyanogen azide, which you get by reacting the bromide with good old sodium azide. Good old sodium azide, which is no mean poison itself, will do that with just about any bromide that’s capable of being displaced at all. Azide is one of the Nucleophiles of the Gods, like thiolate anions – if your leaving group doesn’t leave when those things barge in, you need to adjust your thoughts about it. Cyanogen bromide (or chloride) doesn't stand a chance.
Cyanogen azide is trouble right from its empirical formula: CN4, not one hydrogen atom to its name. A molecular weight of 68 means that you’re dealing with a small, lively compound, but when the stuff is 82 per cent nitrogen, you can be sure that it’s yearning to be smaller and livelier still. That’s a common theme in explosives, this longing to return to the gaseous state, and nitrogen-nitrogen bonds are especially known for that spiritual tendency.
There were scattered reports of the compound in the older German and French literature, but since these referred to the isolation of crystalline compounds which did not necessarily blow the lab windows out, they were clearly mistaken. F. D. Marsh at DuPont made the real thing in the 1960s (first report here, follow-up after eight no-doubt-exciting years here). It's a clear oil, not that many people have seen it that state, or at least not for long. Marsh's papers are, most appropriately, well marbled with warnings about how to handle the stuff. It's described as "a colorless oil which detonates with great violence when subjected to mild mechanical, thermal, or electrical shock", and apologies are made for the fact that most of its properties have been determined in dilute solution. For example, its boiling point, the 1972 paper notes dryly, has not been determined. (The person who determined it would have to communicate the data from the afterworld, for one thing).
The experimental section notes several things that the careless researcher might not have thought about. For one thing, you don't want to make more than a 5% solution in nonpolar solvents. Anything higher and you run the risk of having the pure stuff suddenly come out of solution and oil out on the bottom of the flask, and you certainly don't want that. You also don't want to make a solution in anything that's significantly more volatile than the azide, because then the solvent can evaporate on you, making a more concentrated stock below, and you don't want that, either. Finally, you don't want to put any of these solutions in the freezer - a particularly timely warning, since that's one of the first things many people might be tempted to do - because that'll also concentrate the azide as the solvent freezes. And you don't want that. Do you?
Actually, the careless researcher shouldn't even work with cyanogen azide, or anything like it, but you never can tell what fools will get up to. The compound has around a hundred references in the literature, a good percentage of which are theoretical and computational. Most of the others are physical chemistry, studying its decomposition and reactive properties. You do run into a few papers that actually use it as a reagent in synthesis, but I believe that those can be counted on the fingers, which is a good opportunity to remind oneself why they're all still attached.
In fact, the reason I got to thinking about this wonderful little reagent was a recent paper in Angewandte Chemie, which details the preparation of horrible compounds like the one shown. But what does the experimental section spend the most time warning you about? The cyanogen azide used to make them. Enough said.
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February 26, 2008
Posted by Derek
In a comment to my post on putting out fires last week, one commenter mentioned the utility of the good old sand bucket, and wondered if there was anything that would go on to set the sand on fire. Thanks to a note from reader Robert L., I can report that there is indeed such a reagent: chlorine trifluoride.
I have not encountered this fine substance myself, but reading up on its properties immediately gives it a spot on my “no way, no how” list. Let's put it this way: during World War II, the Germans were very interested in using it in self-igniting flamethrowers, but found it too nasty to work with. It is apparently about the most vigorous fluorinating agent known, and is much more difficult to handle than fluorine gas. That’s one of those statements you don’t get to hear very often, and it should be enough to make any sensible chemist turn around smartly and head down the hall in the other direction.
The compound also a stronger oxidizing agent than oxygen itself, which also puts it into rare territory. That means that it can potentially go on to “burn” things that you would normally consider already burnt to hell and gone, and a practical consequence of that is that it’ll start roaring reactions with things like bricks and asbestos tile. It’s been used in the semiconductor industry to clean oxides off of surfaces, at which activity it no doubt excels.
There’s a report from the early 1950s (in this PDF) of a one-ton spill of the stuff. It burned its way through a foot of concrete floor and chewed up another meter of sand and gravel beneath, completing a day that I'm sure no one involved ever forgot. That process, I should add, would necessarily have been accompanied by copious amounts of horribly toxic and corrosive by-products: it’s bad enough when your reagent ignites wet sand, but the clouds of hot hydrofluoric acid are your special door prize if you’re foolhardy enough to hang around and watch the fireworks.
I’ll let the late John Clark describe the stuff, since he had first-hand experience in attempts to use it as rocket fuel. From his out-of-print classic Ignition! we have:
”It is, of course, extremely toxic, but that's the least of the problem. It is hypergolic with every known fuel, and so rapidly hypergolic that no ignition delay has ever been measured. It is also hypergolic with such things as cloth, wood, and test engineers, not to mention asbestos, sand, and water-with which it reacts explosively. It can be kept in some of the ordinary structural metals-steel, copper, aluminium, etc.-because of the formation of a thin film of insoluble metal fluoride which protects the bulk of the metal, just as the invisible coat of oxide on aluminium keeps it from burning up in the atmosphere. If, however, this coat is melted or scrubbed off, and has no chance to reform, the operator is confronted with the problem of coping with a metal-fluorine fire. For dealing with this situation, I have always recommended a good pair of running shoes.”
Sound advice, indeed. I'll be lacing mine up if anyone tries to bring the stuff into my lab.
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May 30, 2006
Posted by Derek
Perchloric acid almost makes my list by itself, although technically I can't quite include it, since I've already used it. I used the commercial grade, which is 70% strength in water, and it's pretty nasty stuff. It'll chew through your lab coat and give you burns you'll regret, as you'd expect from something that's rather stronger than nitric or sulfuric acid.
But it has other properties. The perchlorate anion is in a high oxidation state, and what goes up, must come down. A rapid drop in oxidation state, as chemists know, is often accompanied by loud noises and flying debris, particularly when the products formed are gaseous and have that pesky urge to expand. If you take the acid up to water-free concentrations, which is most highly not recommended, you'll probably want to wear chain mail, because it's tricky stuff. You can even go further and distill out the perchloric anhydride (dichlorine heptoxide) if you have no sense whatsoever. It's a liquid with a boiling point of around 80 C, and I'd like to shake the hand of whoever determined that property, assuming he has one left.
Perchlorate salts show similar tendencies. The safety literature is just full of alarming stories about old lab benches that had had perchlorates soaked into them years before and exploded when someone banged on them. They're a common component of solid rocket fuels and fireworks, as you'd figure. As with other lively counterions, the alkali metal salts (lithium, sodium, etc.) are comparatively well-behaved, with things heading downhill as you go to larger and fluffier cations. I've used things like zinc and magnesium perchlorate, but I would refuse, for example, to share a room with any visible samples of the lead or mercury salts.
People have made organic perchlorate esters, too, which doesn't strike me as a very good idea - unless, of course, you're actively searching for a way to blow up your rota-vap. Which is exactly what happened in the paper I saw on the synthesis of ethyl perchlorate, as I recall. If you'd like to make your mark, this seems to be a relatively unexplored field. The problem is, the mark you're most likely to make is in the nature of a nasty stain on the far wall.
Perhaps the most unnerving derivative I know of is fluorine perchlorate. That one was reported in 1947 (JACS 69, 677) by Rohrback and Cady. It's easily synthesized, if you're tired of this earthly existence, by passing fluorine gas over concentrated perchloric acid. You get a volatile liquid that boils at about -16 C and freezes at -167.3, which exact value I note because the authors nearly blew themselves up trying to determine it. The liquid detonated each time it began to crystallize, which is certainly the mark of a compound with a spirited nature.
The gas, meanwhile, blows up given any chance at all - contact with a rough surface, with tiny specks of any type of organic matter, that sort of thing. The paper notes that it has "a sharp acid-like odor, and irritates the throat and lungs, producing prolonged coughing". My sympathies go out to whichever one of them discovered that. No, if it's all the same to science, I think I'll let others explore the hidden byways of perchlorate chemistry. . .
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May 23, 2006
Posted by Derek
There are a number of reagents that you used to be able to buy which are no longer around. Some of these have just fallen out of favor, but a compound has to go pretty far down the list before no one sees any profit in selling it. The more common reasons for the disappearance are a bit more dramatic.
A notorious example is "Magic Methyl" (methyl fluorosulfonate). Flurosulfonate is about as good a covalent leaving group as nature provides, and Magic Methyl was accordingly one heck of a way to methylate anions that turned up their noses at anything difficult. Problem was, though, that it also tended to methylate the user. There was at least one fatality in the 1970s from a not-very-large spill of the stuff, and by the time I got to grad school it had been pulled from commercial supply. It's never coming back, either. You can still make the stuff and use it yourself, and people do once in a while (not to mention things that are even more reactive, although that one's not volatile, at least). But there are research organizations that forbid even that.
There are substitutes, but nothing's quite in the same league. Methyl triflate is the closest thing going, as far as I know. It's an open question as to how much less nasty that one is - you can still buy it by the gallon. No one's been killed by it, but if someone dropped a bottle near me I'd still hold my breath and dive out the door.
Dess-Martin reagent is one that's appeared and disappeared over the years. It's a useful oxidizing reagent, which tends to react very cleanly and on some substrates that are hard to work with otherwise. Making it has always been a nerve-wracking process, though. The reagent itself shouldn't be heated, but is reasonably well-behaved. But the intermediate compound in the synthesis (IBX) has been known for some time to be erratically explosive, especially if it's allowed to dry out. It's sensitive to impact, which always made for a good time when it was time to get it out of the funnel after filtering it.
The fun didn't stop there. The last step in the synthesis, right after the IBX formation, was famously wonky, and has only been ironed out in recent years. Or so I'm told - I made a couple of hundred-gram batchs of the stuff, fifteen years ago, going two for three in attempts on the last step, and do not plan to do so again. You can buy the reagent at the moment, but it's been dropped from catalogs before (as Aldrich did during the 1990s).
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March 3, 2005
Posted by Derek
Column VI of the periodic table doesn't start out smelly, but that's probably just because we run on its first element, oxygen. Animal ancestors of ours who felt woozy all the time from the stench of oxygen didn't leave much of a legacy, so we're all pretty positive about it. But when you start moving down into the next rows, everything changes.
Sulfur's next, and its fame as a reeking element is well deserved. Skunks, rotten eggs, burning tires - they all have delightful sulfurous tang, and we have sulfur compounds in the lab that are worse yet. But most people don't think about the elements to come.
The next heavier element in the series is selenium, which most people will have heard of primarily from its presence in health food stores. It is indeed an essential trace element, although I'd think that if your cuisine includes a reasonable amount of garlic (as it should!) then you're getting all the selenium you need. You don't want to overdo it, because this essential dietary factor is also a pretty efficient poison, which is a useful First Lesson in Toxicology right there. (And no, I don't think it's possible to get selenium poisoning from eating too much garlic; I think many other effects would kick in before you noticed any selenium-related problems.)
Selenium compounds are, if anything, more intrinsically noxious than sulfur ones. Imagine a sort of hyperskunk, scattering its enemies before it and making them carom off trees and dive into ponds. The heavier selenium atoms tend to make the compounds less volatile, though, so you don't always get their full bouquet. The smaller compounds get in their licks, though. One of the simpler selenium-rich compounds, for example, is carbon diselenide, an exact homolog of the carbon dioxide in your breath and in your glass of soda. Instead of a gas, the selenide is an oily liquid with a higher boiling point than water. Most of us organic chemists have never seen it.
Which is just fine. The first report of the compound in the chemical literature is from a German university group from 1936, and it was a memorable debut. A colleague of mine had a copy of this paper in his files, and he treasured a footnote from the experimental section which related how the vapors had unfortunately escaped the laboratory and forced the evacuation of a nearby village. The authors stressed the point that its aroma was like nothing that they'd ever encountered.
The compound made a few appearances over the next couple of decades, but one of the next synthetic papers dates from 1963. (That's Journal of Organic Chemistry 28, 1642, for you curious chemists.) The authors are forthright:
"It has been our experience that redistilled carbon diselenide has an odor very similar to that of carbon disulfide. However, when (it is) mixed with air, extremely repulsive stenches are gradually formed. Many of the reaction residues gave foul odors which were rather persistent (and) it should be noted that some of the volatile selenium compounds produced may be extremely toxic as well as foul."
Something for everyone! At least it lets you know when it's coming. Interestingly, in recent years, the compound has actually made a comeback, with more references to it in the past twenty years than in the fifty before. It's been used to prepare a number of odd compounds that have shown promise as organic semi- and superconductors, and there's actually a commercial source for the disgusting stuff (which may be a first.) I'd like to see what they ship it in.
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November 16, 2004
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.
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August 26, 2004
Posted by Derek
The azide group (three nitrogens bonded together in a row, for the non-chemists in the crowd) has several personalities. Unfortunately, most of them are hostile. Azide anion, as you find in sodium azide, is pretty toxic. It shuts down several important enzymes, and it's often used in biology labs as a general metabolic poison.
Covalent azides are a different sort of beast. Having something directly bonded to the group stops it from being an enzyme-killer, for the most part, but you have a new problem to worry about: explosiveness. In general, reasonably high molecular weight azides are OK to handle (e.g., the early anti-HIV drug azidothymidine). I've made some of that sort, since azide displacement is a classic (and useful) way to get a nitrogen into your molecule. But the smaller ones aren't worth the risk.
That's because the higher the percentage of nitrogens in the formula, the more you have to worry. Thermodynamically, nitrogens bonded to each other are always regarded as guilty until proven innocent - there's always the fear that they're going to find a way to throw off their civilized clothes and revert to wild nitrogen gas. That's a hugely stable compound. If your structure goes that route, all that extra bonding energy it used to have ends up diverted into flying shrapnel and loud noises.
A few years ago I saw some Israeli escape artists has prepared triazidomethane, which I wouldn't touch with somebody else's ten-foot titanium pole. One carbon, one hydrogen, and nine nitrogens - look at the time! Gotta run! But there's always worse. Just today I was reading a soon-to-be-published paper in Angewandte Chemiefrom some daredevils at USC. They've prepared titanium tetraazide, of all things. One titanium and twelve nitrogens: whoa! Podiatrist appointment! See you later!
You can isolate the stuff, it seems, as long as you handle it properly. It turns out that brutal treatments like, say, touching it with a spatula, or cooling down a vial of it in liquid nitrogen - you know, rough handling - make it detonate violently. I think that staring hard at it is OK, though. The authors recommend using everything you have for protection if you're zany enough to follow their lead: goggles, blast shield, face shield, leather suit (!) and ear plugs. Those last two suggestions are unique in my experience, and quite. . .evocative of what you have to look foward to with these compounds. (We don't have any leather suits around where I work, although I'm sure I'd look dashing in one.)
Some of the folks on the paper have a joint appointment with an Air Force missile propulsion research lab. They've found a home. Me, I'll be way over here.
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August 4, 2004
Posted by Derek
If you cool things down enough, you can turn almost anything into a liquid (or into a solid, if you're really insane about it.) Chemists use liquid ammonia fairly often, for example, though it's been some years now since I've needed any. People outside the field think of the aqueous solution of ammonia gas (household ammonia) when you say "liquid ammonia", but I'm talking about the pure stuff. Cool the gas down below about -33 C, and you'll condense it out to a clear liquid that's sort of like a thinner version of water.
It's easy enough to do, with an ammonia tank and a condenser full of dry ice. But once, over twenty years ago, I had a chance to see someone use one of those rigs to condense something a bit more exotic: pure hydrogen cyanide. That's another one that people confuse with the aqueous solution. But pure HCN has a fairly high boiling point, for such a small molecule, and condensing out is no problem - as long as you have more nerve than you have sense.
The fellow doing it was down the hall from me in graduate school, and he was doing an obscure reaction which forms a geminal dinitrile, which themselves are rather obscure compounds. (That's probably because this bug-eyed route is the best way to make 'em.) He was dressed in full suit and respirator gear, for which he'd had to get trained. Everyone else had cleared out of the lab, but someone was watching him at all times from the hallway, just in case.
I thought to myself, "When am I going to get the chance to see pure liquid HCN again?", and went down to see, ready to bail out if anything started going wrong. It looked just like ammonia, clear drops rolling down the cold condenser and dripping into the round-bottom flask below. But there was enough HCN in there to kill off the lot of us, if (im)properly handled.
I've worked with plenty of cyanide since then, and even plenty of reactions that have produced small whiffs of HCN vapor. (As I think I've mentioned, it doesn't smell as much like almonds as it's said to, in my opinion.) But I doubt very much if I've worked with enough of it to match the amount that I saw in that flask, that day - there must have been a couple of moles of it in there. A lifetime supply that was, in many sense of the word. . .
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March 28, 2004
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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March 3, 2004
Posted by Hylton Jolliffe
I'm still working on my reply to the Matthew Holt article I mentioned yesterday, so I thought I'd do one of the awful reagents that I spoke of. I'll kick things off with hydrogen fluoride.
The chemically inclined members of my audience might be saying "Hold it! You said yesterday that you'd used hydrofluoric acid!" And that's true, and that stuff is certainly bad enough on its own merits. It gives terribly painful burns, and it eats through glass, to pick two of its fine qualities. But if you're going to be precise, hydrofluoric acid is a water solution of hydrogen fluoride, HF. That's a gas, and it's a lot worse.
Actually, it's just barely a gas. In a cool room it'll condense out as a liquid (it boils at about 20 degrees C, which is 68 F.) The straight liquid must really be a treat, but I've never seen it in that form, and would only wish to through binoculars. It's sold compressed in metal cylinders, like other gases, but it needs some care in packaging. The stuff is so corrosive that special alloys need to be used, usually ones high in nickel. If you stick an ordinary gas regulator on top of an HF cylinder, you're in for trouble, and the complete destruction of the regulator is the least of your worries.
HF has actually been used right out of the cylinder for a long time in Merrifield peptide synthesizers. It's the traditional way to cleave the peptide off the resin at the final step, so there are actually a lot of people who've used the stuff. But it's in a dedicated apparatus that is (that had better be) well sealed, and people treat it with due respect. At a former employer of mine, there was an accident with one of these machines right before I joined the company. The shout "HF LEAK!" went out into the halls, and I'm told that the whole area set a never-to-be-equaled evacuation record. This was one of those drop-things-right-where-you-stand type evacuations, a real sauve qui peut moment.
I've caught some whiffs of HCl, like any chemist has, and it'll wake you up for sure. And I was wrestling with a lecture bottle of HBr gas in grad school, only to have it start to hiss onto my shirt - which was never the same afterwards. But I've never smelled HF, and I hope I never will. As bad as it is on metals and glass, it's much worse on living tissue, although (as I mentioned) a lot of synthetic peptides can stand up to it.
Oddly enough, it's not that strong an acid in the traditional sense. The fluorine doesn't want to let go of the proton enough. It's strong enough to burn, but the big problem is how penetrating it is. As soon as it hits anything moist - like your lungs - it dissolves in the water and turns into hydrofluoric acid again. And that soaks into tissue very readily, with the acid part doing its damage along the way, and the fluoride merrily poisoning enzymes and wreaking havoc. The damage isn't immediately apparent, and there are terrible cases of people who've been exposed and didn't realize it for hours - by which time a lot of irreversible damage had been done.
Fortunately, I have very little cause to even think about using HF. I don't do Merrifield peptide synthesis, and the only times I even use the solution forms of the reagent are on a very small scale and in weakened form (like its complex with pyridine.) Should some lunatic discover a wonderful reaction that requires the gas, I will respectfully pass. As will everyone else.
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