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: email@example.com
Cadmium is bad news. Lead and mercury get all the press, but cadmium is just as foul, even if far fewer people encounter it. Never in my career have I had any occasion to use any, and I like it that way. There was an organocadmium reaction in my textbook when I took sophomore organic chemistry, but it was already becoming obsolete, and good riddance, because this one of those metals that's best avoided for life. It has acute toxic effects, chronic toxic effects, and if there are any effects in between those it probably has them, too.
Fortunately, cadmium is not well absorbed from the gut, and even more fortunately, no one eats it. But breathing it, now that's another matter, and if you're a nonchemist wondering how someone can breath metallic elements, then read on. One rather direct way is if someone is careless enough to floof fine powders of them around you. That's how cadmium's toxicity was discovered in the first place, from miners dealing with the dust. But that's only the start. There's a bottom of the list for breathable cadmium, too, which is quite a thought. The general rule is, if you're looking for the worst organic derivatives of any metal, you should hop right on down to the methyl compounds. That's where the most choking vapors, the brightest flames, and the most panicked shouts and heartfelt curses are to be found. Methyl organometallics tend to be small, reactive, volatile, and ready to party.
Dimethyl cadmium, then, represents the demon plunked in the middle of the lowest circle as far as this element is concerned. I'll say only one thing in its favor: it's not quite as reactive as dimethyl zinc, its cousin one row up in the periodic table. No one ever has to worry about inhaling dimethyl zinc; since it bursts into ravenous flames as soon as it hits the air, the topic just never comes up. Then again, when organozincs burn, they turn into zinc oxide, which is inert enough to be used in cosmetics. But slathering your nose with cadmium oxide is not recommended.
Even though dimethylcadmium does not instantly turn into a wall of flame, it can still liven the place up. If you just leave the liquid standing around, hoping it'll go away, there are two outcomes. If you have a nice wide spill of it, with a lot of surface area, you fool, it'll probably still ignite on its own, giving off plenty of poisonous cadmium oxide smoke. If for some reason it doesn't do that, you will still regret your decision: the compound will react with oxygen anyway and form a crust of dimethyl cadmium peroxide, a poorly characterized compound (go figure) which is a friction-sensitive explosive. I've no idea how you get out of that tight spot; any attempts are likely to suddenly distribute the rest of the dimethylcadmium as a fine mist. Water is not the answer. One old literature report says that "When thrown into water, (dimethylcadmium) sinks to the bottom in large drops, which decompose in a series of sudden explosive jerks, with crackling sounds", and you could not ask for a clearer picture of the devil finding work for idle hands. Or idle heads.
Even without all this excitement, the liquid has an alarmingly high vapor pressure, and that vapor is alarmingly well absorbed on inhalation. a few micrograms (yep, millionths of a gram) of it per cubic meter of air hits the legal limits, and I'd prefer to be surrounded by far less. It's toxic to the lungs, naturally, but since it gets into the blood stream so well, it's also toxic to the liver, and to the kidneys (basically, the organs that are on the front lines when it's time to excrete the stuff), and to the brain and nervous system. Cadmium compounds in general have also been confirmed as carcinogenic, should you survive the initial exposure.
After all this, if you still feel the urge to experience dimethylcadmium - stay out of my lab - you can make this fine compound quite easily from cadmium chloride, which I've no particular urge to handle, either, and methyllithium or methyl Grignard reagent. Purifying it away from the ethereal solvents after that route, though, looks like extremely tedious work, which allows you the rare experience of being bored silly by something that's trying to kill you. It is safe to assume that the compound will swiftly penetrate latex gloves, just like deadly and hideous dimethylmercury, so you'll want to devote some time to thinking about how you'll handle the fruits of your labor.
I'm saddened to report that the chemical literature contains descriptions of dimethylcadmium's smell. Whoever provided these reports was surely exposed to far more of the vapor than common sense would allow, because common sense would tell you to stay about a half mile upwind at all times. At any rate, its odor is variously described as "foul", "unpleasant", "metallic", "disagreeable", and (wait for it) "characteristic", which is an adjective that shows up often in the literature with regard to smells, and almost always makes a person want to punch whoever thought it was useful. We can assume that dimethylcadmium is not easily confused with beaujolais in the blindfolded sniff test, but not much more. So if you're working with organocadmium derivatives and smell something nasty, but nasty in a new, exciting way that you've never quite smelled before, then you can probably assume the worst.
Now, as opposed to some of the compounds on my list, you can find people who've handled dimethylcadmium, or even prepared it, worse luck, although it is an (expensive) article of commerce. As mentioned above, it used to be in all the textbooks as a reliable way to form methyl ketones from acid chlorides, but there are far less evil reagents that can do that for you now. It's still used (on a research scale) to make exotic photosensitive and semiconducting materials, but even those hardy folk would love to find an alternative. No, this compound appears to have no fan club whatsoever. Start one at your own risk.
Over the years, I've probably had more hits on my "Sand Won't Save You This Time" post than on any other single one on the site. That details the fun you can have with chloride trifluoride, and believe me, it continues (along with its neighbor, bromine trifluoride) to be on the "Things I Won't Work With" list. The only time I see either of them in the synthetic chemistry literature is when a paper by Shlomo Rozen pops up (for example), but despite his efforts on its behalf, I still won't touch the stuff.
And if anyone needs any more proof as to why, I present this video, made at some point by some French lunatics. You may observe the mild reactivity of this gentle substance as it encounters various common laboratory materials, and draw your own conclusions. We have Plexiglas, a rubber glove, clean leather, not-so-clean leather, a gas mask, a piece of wood, and a wet glove. Some of this, under ordinary circumstances, might be considered protective equipment. But not here.
When we last checked in with the Klapötke lab at Munich, it was to highlight their accomplishments in the field of nitrotetrazole oxides. Never forget, the biggest accomplishment in such work is not blowing out the lab windows. We're talking high-nitrogen compounds here (a specialty of Klapötke's group), and the question is not whether such things are going to be explosive hazards. (That's been settled by their empirical formulas, which generally look like typographical errors). The question is whether you're going to be able to get a long enough look at the material before it realizes its dream of turning into an expanding cloud of hot nitrogen gas.
It's time for another dispatch from the land of spiderweb-cracked blast shields and "Oh well, I never liked that fume hood, anyway". Today we have a fine compound from this line of work, part of a series derived from N-amino azidotetrazole. The reasonable response to that statement is "Now hold it right there", because most chemists will take one look at that name and start making get-it-away-from-me gestures. I'm one of them. To me, that structure is a flashing red warning sign on a dead-end road, but then, I suffer from a lack of vision in these matters.
But remember, N-amino azidotetrazole (I can't even type that name without wincing) is the starting material for the work I'm talking about today. It's a base camp, familiar territory, merely a jumping-off point in the quest for still more energetic compounds. The most alarming of them has two carbons, fourteen nitrogens, and no hydrogens at all, a formula that even Klapötke himself, who clearly has refined sensibilities when it comes to hellishly unstable chemicals, calls "exciting". Trust me, you don't want to be around when someone who works with azidotetrazoles comes across something "exciting".
It's a beast, all right. The compound is wildly, ridiculously endothermic, with a heat of formation of 357 kcal/mole, all of which energy is ready to come right back out at the first provocation (see below). To add to the fun, the X-ray crystal structure shows some rather strange bond distances, which indicate that there's a lot of charge separation - the azides are somewhat positive, and the tetrazole ring somewhat negative, which is a further sign that the whole thing is trembling on the verge of not existing at all.
And if you are minded to make some yourself, then you are on the verge of not existing at all, either. Both the initial communication and the follow-up publication go out of their way to emphasize that the compound just cannot be handled:
Due to their behavior during the process of synthesis, it was obvious that the sensitivities (of these compounds) will be not less than extreme. . .
The sensitivity of C2N14 is beyond our capabilities of measurement. The smallest possible loadings in shock and friction tests led to explosive decomposition. . .
Yep, below the detection limits of a lab that specializes in the nastiest, most energetic stuff they can think up. When you read through both papers, you find that the group was lucky to get whatever data they could - the X-ray crystal structure, for example, must have come as a huge relief, because it meant that they didn't have to ever see a crystal again. The compound exploded in solution, it exploded on any attempts to touch or move the solid, and (most interestingly) it exploded when they were trying to get an infrared spectrum of it. The papers mention several detonations inside the Raman spectrometer as soon as the laser source was turned on, which must have helped the time pass more quickly. This shows a really commendable level of persistence, when you think about it - I don't know about you, but one exploding spectrometer is generally enough to make recognize a motion to adjourn for the day. But these folks are a different breed. They ended up having to use a much weaker light source, and consequently got a rather ugly Raman spectrum even after a lot of scanning, but if you think you can get better data, then step right up.
No, only tiny amounts of this stuff have ever been made, or ever will be. If this is its last appearance in the chemical literature, I won't be surprised. There are no conceivable uses for it - well, other than blowing up Raman spectrometers, which is a small market - and the number of research groups who would even contemplate a resynthesis can probably be counted on one well-armored hand.
This fine reagent was mentioned here (disparagingly) in the comments the other day, and I knew that it was time to add it to the list. I've had some other selenium entries before, and they're all here for the same reason: their unsupportable stenches. Everyone, even people who've never had a chemistry class in their lives, knows that sulfur compounds are stinky, of course, but it's a problem that continues as you move down Group XVI of the periodic table.
And it's not like plain phenol itself has no odor. It's strong, penetrating, and completely unmistakable. As soon as I get a whiff of the stuff, I'm immediately transported back to the Verser Clinic, the small hospital in the town I grew up in back in Arkansas. Phenol smells like an old-fashioned medical office; it was used for many years as a disinfectant (and was, in fact, introduced as such by Joseph Lister himself). If you move it down a notch to sulfur, you get thiophenol, which is easy to describe: burning rubber - the pure, potent, platonic ideal of burning rubber, bottled up and daring you to open the cap. I can't say that I won't work with thiophenol, since I have (very much to my regret, at times), but I've used it most reluctantly, and probably haven't touched it in at least fifteen years.
Ah, but move down one more element and you have selenophenol, and that's a more exotic reagent. I've never seen any, and after reading the descriptions, I never want to. Actually, let me take that back: I'd look at some from the other end of the lab. What I never want to do is open any of it up. The chemical literature has numerous examples of people who are at a loss for words when it comes to describing its smell, but their attempts are eloquent all the same. A few years ago, Gaussling at the Lamentations on Chemistry blog referred to it as "The biggest stinker I have run across. . .Imagine 6 skunks wrapped in rubber innertubes and the whole thing is set ablaze. That might approach the metaphysical stench of this material." So we'll start with that.
I believe that this lovely compound is commercially available (if you're anywhere close to anyone making it, you'll know about it). But should you wish to prepare it with your own hands, do violence to your own schnozz, and drape yourself out of your own window while you throw up into your own rhododendrons, feel free to use this reliable preparation from Organic Syntheses, circa 1944. This features the note that "it is frequently advisable to work with [selenium compounds] on alternate days", which I suppose is to give them time to work their way out of your nasal passages.
I'm not so sure. When I was a teaching assistant in grad school, I taught three labs a week one semester, and one of those labs, damn it all, was the phenyl Grignard reagent. We had them making it in diethyl ether, outside of the small and inadequate fume hoods, and the solvent fumes were fit to strip paint. By the end of the Monday lab, I was well saturated with ether and had a terrible headache, which returned as soon as I caught my first whiff of the stuff on Tuesday afternoon. I barely made it through that lab, mostly by holding my breath and using a lot of hand gestures, and I took the opportunity on Wednesday to get as much fresh air as I could. But when I came back for the Thursday session, the first first wave of ether vapor washed over me and nearly stretched me out on the tiles. I taught the entire lab from the hallway, shouting and waving like Monty Python's "Semaphore Version of Wuthering Heights". So in my mind, the choice between getting these things over with and stretching them out is still not settled.
That Org Syn prep also notes that it can produce small amounts of hydrogen selenide, which is very toxic indeed (and will give you a sore throat, too, apparently, before it kills you). This luckless graduate student from the 1920s got to experience both of these bracing selenium room fresheners in the course of his work:
Berzelius described the poisonous effect of hydrogen selenide quite impressively; "In order ta get acquainted with the smell of this gas I allowed a bubble not larger than a pea to pass into my nostril ; in consequence of its smell I so completely loss my sense of smell for several hours that I could not distinguish the odor of strong ammonia even when held under my nose. My sense of smell returned after five or six hours, but severe irritation of the mucous membrane set in and persisted for a fortnight' The writer has been working on the gas for some time and was also quite seriously affected once, the injury persisting for many days. That it is more poisonous than the hydrogen sulphide is well known."
So you have that to look forward to on your way to selenophenol. And at your destination? Assuming your nose is still attached to your face, you'll experience what few chemists ever have. I'll let this 1908 report from Wisconsin take over:
When benzeneselenonic acid in solution is treated with reducing agents such as hydrogen sulphide, sulphur dioxide, or, best, with zinc and hydrochloric, acid selenophenol is obtained as a yellow oil with an overpowering and most nauseating odor. . .The odor of diphenyl diselenide is extremely disagreeable but is not nearly so bad as that of selenophenol.
. . .The effect of selenophenol on the skin is very similar to that of thiophenol, forming blisters which itch intensely. After a time, these dry up, the skin scales off, and there appears to be a deposit of red selenium beneath it. The odor of selenophenol is very penetrating, and is nauseating beyond description.
Gloves, man, gloves. Unless, of course, you wish to be tattooed with elemental selenium while being nauseated beyond description. Should this be your idea of a fun Saturday night, I will not stand in your way.
Let's start with the name. Quite a mouthful, isn't it? Believe me, that one's pretty chewy even for experienced organic chemists. We see lots of more complicated nomenclature, of course, but this one some features some speed bumps, that make you go back to make sure that you're reading it correctly. I'll take you through my own thoughts as an example.
You skip to the end in chemical names - Mark Twain would have felt about them the same way he felt about the German language. But this brings me up short, because very few chemists could walk up to the board and draw an isowurtzitane. And I am not among their number. I have a vague picture of these "wurtz" compounds being funky three-dimensional cage structures, and that much only from having probably read too many photochemistry papers over the years. So the only thing that "isowurtzitane" calls to mind is some complicated framework of fused rings, looking like one of those wire sculptures that unexpectedly fold up flat when you pull on them.
Moving on out, as you do in a systematic name, I see that this is a hexaaza variation, which makes the picture a bit fuzzier. That's a lot of nitrogens substituted for carbons, and the first thought is that this must be some weirdo condensation product of ammonia, some aldehyde, and who knows what. You can get some pretty funny-looking structures that way, like hexamethylenetetramine (which I've actually used a couple of times). I don't know where those nitrogens are, I think to myself, but I'll bet that's how they got there, because any other pattern would be a synthetic nightmare. So far, so good. But now comes the unexpected habanero.
Hexanitro? Say what? I'd call for all the chemists who've ever worked with a hexanitro compound to raise their hands, but that might be assuming too much about the limb-to-chemist ratio. Nitro groups, as even people who've never taken a chemistry class know, can lead to firey booms, and putting six of them on one molecule can only lead to such. And since there are six nitrogens and six nitro groups, the first assumption must be that these are all bonded to each other. I mean, come on, leaving the nitro groups attached to the carbons is for wimps. So that means that someone, somewhere, has perversely made a poly-N-nitro cage compound, as if they'd been dared to cram the most bond energy into the smallest space.
That, as it happens, is exactly the case. Hexanitrohexaazaisowurtzitane, or CL-20, was developed as a highly energetic, compact, and efficient explosive. What makes it unusual is not that it blows up - go find me a small hexa-N-nitro compound that doesn't - but that it doesn't actually blow up immediately, early, and often. No, making things that go off when someone down the hall curses at the coffee machine, that's no problem. Making something like this that can actually be handled and stored is a real accomplishment.
Not that it's what you'd call a perfect compound in that regard - despite a lot of effort, it's still not quite ready to be hauled around in trucks. There's a recent report of a method to make a more stable form of it, by mixing it with TNT. Yes, this is an example of something that becomes less explosive as a one-to-one cocrystal with TNT. Although, as the authors point out, if you heat those crystals up the two components separate out, and you're left with crystals of pure CL-20 soaking in liquid TNT, a situation that will heighten your awareness of the fleeting nature of life.
Stabilized or not, I still won't get near it. For one thing, I'm a drug discovery chemist, and if you think a structure like this is going to be a drug, then you must be on some strong ones yourself. No, the thing about these compounds is that they can be handled as long as they're very pure and formulated just right. The side products from their synthesis, well, those might not be so nice. And if a batch gets contaminated, or doesn't come out so clean, well, that might not be so nice, either. Synthesizing polynitro compounds is no chocolate fondue party, either: if you picture a bunch of guys wheeling around drums of fuming nitric acid while singing the Anvil Chorus from Il Trovatore, you're not that far off the mark. You really have to beat the crap out of a molecule to get that many nitro groups on it, which means prolonged heating of things that you'd really rather not heat up at all.
No, I'll leave the can-you-top-this nitration chemistry to those that love it. You guys just go ahead and stuff as many energetic bonds as you can into the smallest tangles; I'll be over here in the bunker cheering you on, and jumping a foot in the air every time someone sneezes. I'm not cut out for hexanitro anything.
Addendum: it's an odd thing, but when you search for information on this compound, a significant number of the Google hits are for its environmental effects. This is an explosive, meant for munitions and destruction, but there are all kinds of studies on its effects on earthworms, fish, soil microorganisms, and so on. Steven Pinker must be right when he says that violence is getting tamer all the time.
Well, it's been a bit too serious around here this week. So I thought today I'd step back to a period when men were men and chlorine azide was a reactive, toxic, and unstable compound that was only good for finding out what sort of explosion it would set off next. What's that? You say that that's still about all it's good for? Staying power, that's what I call it. If you work with the halogen azides, you work with things whose essential character time does not alter.
"Until they blow up", you say. Ah, but that is their essential character. It's the things around them that alter. Make sure you don't put anything next to them that you're not comfortable seeing altered - you know, all sudden-like.
A reader forwarded this 1943 JACS article, the first comprehensive study of chlorine azide, and it's a joy to read. Part of the fun is, of course, watching these folks set off the fireworks. (The challenge with a substance like chlorine azide is finding something that it won't react with violently):
Owing to the extreme instablity of the compound accurate determinations of the bioing and melting points have not been made as yet. Numerous explosions, often without assignable cause, have occurred during the experiments. . ."
Another thing I always enjoy in these papers is the list of recommended protective gear. No leather suits this time and (interestingly) no earplugs. Nope, it's straight to the Iron Man look. These azidonauts endorse:
". . .masks and breast-plates of sheet iron worn by observers during times of danger. Each mask is provided with a rectangular pane (7 x 3 inches) of shatter-proof glass. Although scores of violent detonations have occurred, with resultant demolition of much apparatus, no personal injury has been suffered."
That last part is sort of a "no graduate students were maimed during the course of this research" statement, which really is good to know. But another nice thing about this paper is the way some parts of it are written, in a style which was a bit formal and archaic even for 1943. A sample:
"If small pieces of yellow phosphorus be added, with stirring, to a solution of chlorine azide in carbon tetrachloride at 0C, the solution gradually becomes turbid, and a succession of slight explosions takes place beneath the liquid. If stirring be omitted until the maximum turbidity is attained, the slightest agitation results in a detonation that demolishes the apparatus. . ."
Do not be omitting the stirring, then. I have to say, not being used to this sort of chemistry, that if I saw these events going on in my fume hood that a series of slight explosions might well take place beneath my iron breastplate. What else doth chlorine azide detonate with? Well, in case you had any doubt, the gaseous reagent "reacts violently" with sodium metal. They had four explosions at -78C, while the fifth run (persistence!) yielded a mixture of sodium chloride and sodium azide. (Actually, the other runs probably yielded that, too, albeit as a fine haze). I really have to salute the dedication involved in finding that out, though - after two or three violent explosions, you or I might be tempted to just say that we couldn't determine the products of the reaction. But they were made of sterner stuff back in 1943.
The date does make me wonder if there was war research money involved; I wouldn't be surprised. But chlorine azide has not been weaponized, nor will it be. It remains, with its chemical relatives, off in a part of chemical science that's safe from human exploitation. It's a spacious game preserve, that territory, and over the gate is the ornate motto Noli me tangere. Take heed.
Tetrazole derivatives have featured several times here in "Things I Won't Work With", which might give you the impression that they're invariably explosive. Not so - most of them are perfectly reasonable things. A tetrazole-for-carboxyl switch is one of the standard med-chem tricks, standard enough to have appeared in several marketed drugs. And that should be recommendation enough, since the FDA takes a dim view of exploding pharmaceuticals (nitroglycerine notwithstanding; that one was grandfathered in). No, tetrazoles are good citizens. Most of the time.
It's when they get put in with the wrong sort of company that they turn delinquent. What with four nitrogens in the ring and only one carbon, they do have a family history of possible trouble - several sections of this blog category could just as accurately be called Things That Suddenly Want To Turn Back Into Elemental Nitrogen. And thermodynamically, there aren't many gently sloping paths down to nitrogen gas, unfortunately. Both enthalpy and entropy tilt things pretty sharply. A molecule may be tamed because it just can't find a way down the big slide, but if it can, well, it's time to put on the armor, insert the earplugs, and get ready to watch the free energy equation do its thing right in front of your eyes. Your heavily shielded eyes, that is, if you have any sense at all.
Nitro groups are just the kind of bad company I mean, since they both bring their own oxygens to the party and pull electrons around in delightfully destabilizing ways. So nitrotetrazole is already not something I'd feel good about handling (its metal salts are primary explosives), but today's paper goes a step further and makes an N-oxide out of a nitrogen on a nitrotetrazole ring. This both adds more oxygen and tends to make the crystal packing tighter, which raises the all-important kapow/gram ratio. (There is, of course, little reason to do this unless you feel that life is empty without sudden loud noises). The paper mentions that "Introducing N-oxides onto the tetrazole ring may . . . push the limits of well-explored tetrazole chemistry into a new, unexplored, dimension.", but (of more immediate importance) it may also push pieces of your lab equipment into unexplored parts of the far wall.
Turns out that you can make the N-oxides through pretty mild chemistry (oxone at room temp), which is surprising. Until now, only a handful of the things had been made, most using an indirect route using hydrazoic acid (next!). The only direct oxidation of a tetrazole had been done with the relentlessly foul hypofluorous acid(next! keep moving!), which itself has to be made fresh from fluorine gas (next! thank you! next!). These recipes pretty much excluded most reasonable people from skipping through this green and sunny field of knowledge. Still, if you're a reasonable person, you're probably not yearning to make nitrotetrazole oxides in the first place. These things have a way of evening out.
So what are these fine new heterocycles like, anyway? Well, the authors prepared a whole series of salts of the parent compound, using the bangiest counterions they could think of. And it makes for quite a party tray: the resulting compounds range from the merely explosive (the guanidinium salts) to the very explosive indeed (the ammonium and hydroxyammonium ones). They checked out the thermal stabilities with a differential scanning calorimeter (DSC), and the that latter two blew up so violently that they ruptured the pans on the apparatus while testing 1.5 milligrams of sample. No, I'm going to have to reluctantly give this class of compounds a miss, as appealing as they do sound.
Several of the new compounds show similar detonation properties to RDX, albeit with less thermal stability. This is one reason we're reading about them in the open literature; the ideal explosive acts incredibly stable under a wide range of conditions, then loses its composure all at once at just the specified moment. We don't seem to be quite there yet. What I expect is that the authors are probably trying to work this same N-oxide magic on the azidotetrazolates instead of the nitro compounds. Now, that'll be a hoppin' bunch of compounds - for all I know, the research groups involved have already tried this and just haven't been able to get anything out past the lip of the flask yet. I'll be monitoring the literature for signs. On the other hand, if you live in Münich or College Park, you can probably monitor the progress of this work by listening for distant booming noises and the tinkle of glass.
That would be my reaction if asked to take a look at the structures in this new paper in JACS. As the authors, who tiptoe gingerly every morning into the State Key Laboratory of Explosion Science and Technology in Beijing put it:
"The larger the number of directly linked nitrogen atoms, the more difficult the compound is to synthesize. The difficulties in synthesizing and handling polynitrogen compounds are a direct consequence of their high endothermicities; a further complication is the almost complete absence of methodology for preparing such compounds."
Every word of that is true, and doesn't it just sound appealing? Here, go invent some completely new chemistry in order to make some compounds that are just trembling with the desire to explode. And if the reactions don't work? No prob: all the side products will probably be horribly explosive, too. Good luck determining just which one of them it was that demolished your hood!
But hold on. At first glance, this structure is terrifically unappealing, unless your chemical sensibilities are bent the right way, in which case, there's not much hope for you. The beast has eight nitrogens in a row, which I believe ties the current record. What's startling about the compound is that it's weirdly stable: it doesn't decompose until nearly 194 degrees C, which is quite bizarre. You'd think, by looking at it, that it would hop up and do its big death scene at about one-tenth that temperature. I mean, I've made potential drug candidates that fell apart at lower temperatures than that. (The amount of electron delocalization this compound has probably keeps its personality from coming through).
The other odd thing about this one is that it changes color on exposure to light. That central double bond will flip around to cis instead of trans, which changes the color of the crystals from yellow to blue. (I remember making a photochromic compound of this sort in an undergrad experiment, which I believe was some sort of Chichibabin pyridine thingie; it sure as heck wasn't this!) Exposing this sort of structure to UV light also isn't the first thing I'd want to do, either - the fact that it'll reversibly go through a transition like that also points to its mild, friendly nature.
But heck, we can fix that: hang some nitro groups off of it, guys! Put some more nitrogens in those rings! Go for the record! Surely the State Key Laboratory of Explosion Science can make some compounds that, you know, explode. Look, guys, I've had Chinese colleagues that seemed to have no problem making things that blew up. (To be sure, I've known people from a number of different backgrounds who had that talent; it springs up everywhere) So I know that you can do it.
I can't even decide whether to put this in the "Things I Won't Work With" category at all, since it looks like I could not only work with it, but beat on it with a ball-peen hammer. What kind of polyaza compounds are people turning out these days, anyway?
For once, I'm going to farm out a "Things I Won't Work With" post to someone else. For those who missed it in the comments, here's the link to the PDF of Max Gergel's extraordinary memoir "Excuse Me Sir, Would You Like to Buy a Kilo of Isopropyl Bromide?" Gergel founded Columbia Organic Chemicals, and if you want to see how it was done in the Old Days, this is the place to go. A sample:
". . .As we chatted, as if the thought had struck him for the first time, the old rogue said, "You know Gergel, I have a prep you could run for us which would make you a lot of money." Now this was the con working on the con. When my mother told me that a gentleman had called from town asking to visit Dr. Gergel there was no one at the plant except the two of us; when Parry, whom I already knew by reputation, sauntered in disguised as a simple country bumpkin I knew he was the director of research for Naval Research Labs, and his mission was to find someone foolhardy enough to make pentaborane. News travels. I met him at the door and told him that I was simply a lab flunky but would fetch Mr. Gergel, that my boss was extremely smart but had been prevented by the war effort (in which he had served valiantly and with distinction) from getting a PhD; that right now Mr. Gergel was extremely busy with priority reaction but would be able to see him in ten minutes—which gave me time to change my clothes and wash my face. He never realized that we were the same person. Parry chatted with me in the breezy, confidential voice that has been used by every con man since Judas Iscariot and told me that the only reason that the Navy was willing to farm out this fascinating project was simply luck of qualified personnel. That my splendid contribution to Manhattan District was well known by the military, that people spoke of me as a true Southern prodigy. (The old devil was so good that I listened with gradually increasing preparedness to make pentaborane, although I had been forewarned that it was dog with a capital "D". . .
I came across the book in Duke's chemistry library in 1984, a few years after its publication, and read it straight through with my hair gradually rising upwards. Book 2 is especially full of alarming chemical stories. I suspect that some of the anecdotes have been polished up a bit over the years, but as Samuel Johnson once said, a man is not under oath in such matters. But when Gergel says that he made methyl iodide in an un-air-conditioned building in the summertime in South Carolina, and describes in vivid detail the symptoms of being poisoned by it, I believe every word. He must have added a pound to his weight in sheer methyl groups.
By modern standards, another shocking feature of the book is the treatment of chemical waste. Readers will not be surprised to learn that several former Columbia Organic sites feature prominently in the EPA's Superfund cleanup list, but they certainly aren't alone from that era.
Everyone's heard of cyanide, whether they've spent any time in a chemistry classroom or not. And if you form a covalent bond to the carbon of that CN group, you've got a nitrile, and those are familiar compounds to any organic chemist. But what if you flip the group around and bond it via the nitrogen? That gives you a weird situation, where the nitrogen has a formal positive charge and the carbon is left with a formal negative one, which looks somehow unnatural. But that's an isonitrile (isocyanide) for you.
They're actually quite useful, although I'd guess that the majority of chemists have never encountered one. But if they have, they've remembered it, because isonitriles are not shy about announcing their alien character. Our noses can immediately tell the difference between garden variety nitriles and their evil twins. The former often have no smell at all, or run to a faint spicyness. The latter smell like. . . like. . .well, I've never actually been downwind of the Abominable Snowman's armpit or been had my eyeglasses fogged up by a Komodo dragon with stomach trouble, but those are the examples that come to mind.
Fragrance expert Luca Turin has described isonitriles as "the Godzilla of scent", and that's accurate, if you also try to imagine Godzilla's gym socks. "Penetrating" and "repulsive" are good words to describe your typical isocyanide. It feels like the odor is aggressively storming your nasal passages, and it really makes you want to be somewhere else very quickly. This abstract gets the point across well. Problem is, it can be one of those smells that stays with you ("I like it in here. Stop bothering me."), so going somewhere else - although a recommended first step - is not always enough to do the trick. As a paper on the synthesis of these fine compounds puts it:
It should also be noticed that due to the extremely distressing odour of isocyanides the application of the usual techniques of purification is especially difficult since the exposure to isocyanide vapours, even at very low levels, must be rigorously avoided.
Good advice! But hard to put into practice if you use the things at all, since it doesn't take much. Even the not-so-volatile isonitriles get up on the table and shout at you, but the low-molecular-weight ones are truly hard to take. And the pride of that bunch seems to be the n-butyl, which should come as no surprise. Straight-chain butyl compounds are well known to be just a poor match for human sensibilities. Butyl alcohol is stinky, butylamine foul, butyraldehyde reeks, butyric acid is famously disgusting, and butyl mercaptan is a standout even in the vile crowd of thiols.
So butyl isocyanide is, well, something to experience. I've never had the pleasure, and will take pains not to. I can do no better than to quote the 1937 observations of one of the first groups to figure out how to prepare this noble reagent in quantity:
Butyl isocyanide proved to be so disagreeable to manipulate that none of its physical constants except boiling point were determined. Even when a hood with an extra forced draft was used, the odor pervaded the laboratory and adjoining rooms, deadening the sense of smell and producing in the operator, and in others, severe headaches and nausea which usually persisted for several days.
Sounds great. Many of you may have had similar experiences at some point, but it usually takes more than just spilling a drop of stuff on the floor to bring them on. Here's the whole thing, in a bottle! Available wherever fine chemicals are sold, actually. Don't just live an average, boring life: go wild - go isocyanide!
The latest addition to the long list of chemicals that I never hope to encounter takes us back to the wonderful world of fluorine chemistry. I'm always struck by how much work has taken place in that field, how long ago some of it was first done, and how many violently hideous compounds have been carefully studied. Here's how the experimental prep of today's fragrant breath of spring starts:
The heater was warmed to approximately 700C. The heater block glowed a dull red color, observable with room lights turned off. The ballast tank was filled to 300 torr with oxygen, and fluorine was added until the total pressure was 901 torr. . .
And yes, what happens next is just what you think happens: you run a mixture of oxygen and fluorine through a 700-degree-heating block. "Oh, no you don't," is the common reaction of most chemists to that proposal, ". . .not unless I'm at least a mile away, two miles if I'm downwind." This, folks, is the bracingly direct route to preparing dioxygen difluoride, often referred to in the literature by its evocative formula of FOOF.
Well, "often" is sort of a relative term. Most of the references to this stuff are clearly from groups who've just been thinking about it, not making it. Rarely does an abstract that mentions density function theory ever lead to a paper featuring machine-shop diagrams, and so it is here. Once you strip away all the "calculated geometry of. . ." underbrush from the reference list, you're left with a much smaller core of experimental papers.
And a hard core it is! This stuff was first prepared in Germany in 1932 by Ruff and Menzel, who must have been likely lads indeed, because it's not like people didn't respect fluorine back then. No, elemental fluorine has commanded respect since well before anyone managed to isolate it, a process that took a good fifty years to work out in the 1800s. (The list of people who were blown up or poisoned while trying to do so is impressive). And that's at room temperature. At seven hundred freaking degrees, fluorine starts to dissociate into monoatomic radicals, thereby losing its gentle and forgiving nature. But that's how you get it to react with oxygen to make a product that's worse in pretty much every way.
FOOF is only stable at low temperatures; you'll never get close to RT with the stuff without it tearing itself to pieces. I've seen one reference to storing it as a solid at 90 Kelvin for later use, but that paper, a 1962 effort from A. G. Streng of Temple University, is deeply alarming in several ways. Not only did Streng prepare multiple batches of dioxygen difluoride and keep it around, he was apparently charged with finding out what it did to things. All sorts of things. One damn thing after another, actually:
"Being a high energy oxidizer, dioxygen difluoride reacted vigorously with organic compounds, even at temperatures close to its melting point. It reacted instantaneously with solid ethyl alcohol, producing a blue flame and an explosion. When a drop of liquid 02F2 was added to liquid methane, cooled at 90°K., a white flame was produced instantaneously, which turned green upon further burning. When 0.2 (mL) of liquid 02F2 was added to 0.5 (mL) of liquid CH4 at 90°K., a violent explosion occurred."
And he's just getting warmed up, if that's the right phrase to use for something that detonates things at -180C (that's -300 Fahrenheit, if you only have a kitchen thermometer). The great majority of Streng's reactions have surely never been run again. The paper goes on to react FOOF with everything else you wouldn't react it with: ammonia ("vigorous", this at 100K), water ice (explosion, natch), chlorine ("violent explosion", so he added it more slowly the second time), red phosphorus (not good), bromine fluoride, chlorine trifluoride (say what?), perchloryl fluoride (!), tetrafluorohydrazine (how on Earth. . .), and on, and on. If the paper weren't laid out in complete grammatical sentences and published in JACS, you'd swear it was the work of a violent lunatic. I ran out of vulgar expletives after the second page. A. G. Streng, folks, absolutely takes the corrosive exploding cake, and I have to tip my asbestos-lined titanium hat to him.
Even Streng had to give up on some of the planned experiments, though (bonus dormitat Strengus?). Sulfur compounds defeated him, because the thermodynamics were just too titanic. Hydrogen sulfide, for example, reacts with four molecules of FOOF to give sulfur hexafluoride, 2 molecules of HF and four oxygens. . .and 433 kcal, which is the kind of every-man-for-himself exotherm that you want to avoid at all cost. The sulfur chemistry of FOOF remains unexplored, so if you feel like whipping up a batch of Satan's kimchi, go right ahead.
Update: note that this is 433 kcal per mole, not per molecule (which would be impossible for even nuclear fission and fusion reaction (see here for the figures). Chemists almost always thing in energetics in terms of moles, thus the confusion. It's still a ridiculous amount of energy to shed, and you don't want to be around when it happens.
So does anyone use dioxygen difluoride for anything? Not as far as I can see. Most of the recent work with the stuff has come from groups at Los Alamos, where it's been used to prepare national-security substances such as plutonium and neptunium hexafluoride. But I do note that if you run the structure through SciFinder, it comes out with a most unexpected icon that indicates a commercial supplier. That would be the Hangzhou Sage Chemical Company. They offer it in 100g, 500g, and 1 kilo amounts, which is interesting, because I don't think a kilo of dioxygen difluoride has ever existed. Someone should call them on this - ask for the free shipping, and if they object, tell them Amazon offers it on this item. Serves 'em right. Morons.
Organometallic reagentss come from large tribes, and there are always wild cousins up in the hills. A good place to look for the livelier ones is in the simplest alkyl derivatives, and you should go all the way down to the methyls if you want to know their real character. Ignore the halides. Methylmagnesium bromide you can get in multiliter kegs; they might as well sell it in Pottery Barn.
Dimethylmagnesium, though, is not an article of commerce. I've made it myself. So although it's definitely something you want to keep an eye on, I can't very well say that I won't work with it. And the other metals? Dimethyl mercury I will not get within yards of, for very well-founded reasons. Trimethylaluminum is a flamethrower extraordinaire, with a solid reputation among pyromaniacs. I've used the stuff, although I wasn't whistling while I was syringing it out. Handling it in solution, as I did, is less stressful than using the pure stuff - I'd definitely want to sit down and think about that one.
But neat dimethyl zinc. . .no, I don't think so. A colleague of mine made some in graduate school, and came down the hall to us looking rather pale. He'd disconnected a length of rubber tubing from his distillation apparatus and seen it go up in immediate, vigorous flames. "This stuff makes t-butyllithium look like dishwater" is the statement I remember from that evening. You can buy the pure stuff from Alfa, if you're inclined to run a head-to-head comparison. Do make sure to post the video on YouTube; that's as close as I want to get.
One problem is that it's a pretty volatile compound, boiling at 46C, so there's plenty of vapor around to start a party. The diethyl analog is a bit better, but it's nearly as pyrophoric. The Library of Congress discovered this in the 1980s and 1990s, during a long-running project to deacidify old documents. The diethyl zinc reacts with the acid in aged wood-pulp papers, neutralizing it, lightening the color, and stiffening the paper, so you'd think it would be ideal. Well, except for the instant-bursting-into-ravenous-flames part. Making sure that all the reagent was gone before opening the hatch, that was rather important. The pilot plant for this process suffered from some regrettable explosive bonfires before the whole idea was abandoned. Interestingly, one of the biggest problems seems to have been that the treated books were (at least at first) rather odorous, and some colored book covers were initially affected. You can sense a certain testiness about these issues in the Library's final report on the subject:
It has also been established that tight or loose packing of books; the amount of alkaline reserve; reactions of DEZ with degradation products, unknown paper chemicals and adhesives; phases of the moon and the positions of various planets and constellations do not have any influence on the observed adverse effects of DEZ treatment.
You'll notice that the LOC didn't even bother with the dimethyl compound, and I think I'll take a tip from them.
My recent entries in this category have, for the most part, been hazardous in a direct (not to say crude, or even vulgar) manner. These are compounds that explode with bizarre violence even in laughably small amounts, leaving ruined equipment and shattered nerves in their wake. No, I will not work with such.
But today's compound makes no noise and leaves no wreckage. It merely stinks. But it does so relentlessly and unbearably. It makes innocent downwind pedestrians stagger, clutch their stomachs, and flee in terror. It reeks to a degree that makes people suspect evil supernatural forces. It is thioacetone.
Or something close to it, anyway. All we know for sure is that thioacetone doesn't like to exist as a free compound - it's usually tied up in a cyclic thioketal trimer, when it's around at all. Attempts to crack this to thioacetone monomer itself have been made - ah, but that's when people start diving out of windows and vomiting into wastebaskets, so the quality of the data starts to deteriorate. No one's quite sure what the actual odorant is (perhaps the gem-dimercaptan?) And no one seems to have much desire to find out, either.
There are sound historical reasons for this reluctance. The canonical example (Chemische Berichte 1889, 2593) is the early work in the German city of Freiburg in 1889 (see page 4 of this textbook), which quotes the first-hand report. This reaction produced"an offensive smell which spread rapidly over a great area of the town causing fainting, vomiting and a panic evacuation.". An 1890 report from the Whitehall Soap Works in Leeds refers to the odor as "fearful", and if you could smell anything through the ambient conditions in a Leeds soap factory in 1890, it must have been.
The compound shows up sporadically in the literature until the mid-1960s, when several groups looked into thioketones as sources of new polymers. The most in-depth analysis took place at the Esso Research Station in Abingdon, UK, where Victor Burnop and Kenneth Latham got to experience the Freiburg Horror for themselves:
"Recently we found ourselves with an odour problem beyond our worst expectations. During early experiments, a stopper jumped from a bottle of residues, and, although replaced at once, resulted in an immediate complaint of nausea and sickness from colleagues working in a building two hundred yards away. Two of our chemists who had done no more than investigate the cracking of minute amounts of trithioacetone found themselves the object of hostile stares in a restaurant and suffered the humiliation of having a waitress spray the area around them with a deodorant. The odours defied the expected effects of dilution since workers in the laboratory did not find the odours intolerable ... and genuinely denied responsibility since they were working in closed systems. To convince them otherwise, they were dispersed with other observers around the laboratory, at distances up to a quarter of a mile, and one drop of either acetone gem-dithiol or the mother liquors from crude trithioacetone crystallisations were placed on a watch glass in a fume cupboard. The odour was detected downwind in seconds."
Now that's a compound to be taken seriously. How do you work with something that smells like hell's dumpster? Like this:
"The offensive odors released by cracking trithioacetone to prepare linear poly(thioacetone) are confined and eliminated by working in a large glove box with an alkaline permanganate seal, decontaminating all apparatus with alkaline permanganate, eliminating obnoxious vapors with nitrous fumes generated by a few grams of Cu in HNO3, and destroying all residues by running them into the center of a wood fire in a brazier."
So there you have it - just install a fireplace next to your hood (what every lab needs, for sure) and remember that, in a thioacetone situation, fogging the area with brown nitrogen oxide fumes will actually improve the air. (This is from Chemistry and Industry, 1967, p. 1430, if you need more details, and I hope you don't).
The Klapötke group at Munich are some of the masters of alarming chemical structures, and they basically seem to own the field of chalcogen azides. Perhaps the competition for this class of compounds is not as intense as it might be - the other labs doing this sort of thing are collaborations between USC and various military research wings. But they're still interesting beasts.
A few years ago, both groups reported the synthesis of tellurium azides, with the Munich group sending in their paper a few days before the USC/Air Force team sent in theirs. The parent tetra-azide was explosive, to be sure, but could be kept at room temperature without necessarily blowing up. Klapötke's group and the USC group (led by Karl Christe) then teamed up to tackle the corresponding selenium analogs, which were reported in 2007.
And they're a livelier bunch. The selenium tetra-azide is another yellow solid, like the tellurium compound, but it's rather harder to keep it down on the farm. Taking some selenium tetrafluoride (see below) and condensing it with trimethylsilyl azide at -196 °C did the trick. After warming things up (you'll note the relative use of that term "warming"), they saw that:
"Within minutes, the mixture turned yellow, the color intensified, and a lemon-yellow solid precipitated while the reaction proceeded. Keeping the reaction mixture for about 15 min at -64 °C resulted in a violent explosion that destroyed the sample container and the surrounding stainless-steel Dewar flask."
Did I mention that this prep was performed on less than one millimole? Spirited stuff, that tetra-azide. The experimental section of the paper enjoins the reader to wear a face shield, leather suit, and ear plugs, to work behind all sorts of blast shields, and to use Teflon and stainless steel apparatus so as to minimize shrapnel. Hmm. Ranking my equipment in terms of its shrapneliferousness is not something that's ever occurred to me, I have to say. It's safe to assume that any procedure which involves considering which parts of the apparatus I'd prefer to have flying past me will not get much business in my lab, no matter how dashing I might look in a leather suit.
That procedure deserves a closer look, though. You can't just crack open a can of selenium tetrafluoride whenever you feel the urge, you know. That stuff has to be made fresh, as far as I can see, and the way these hearty sons of toil make it is by reacting selenium dioxide with chlorine trifluoride. Yep, that stuff, the delightful compound that sets sand on fire and eats through asbestos firebrick.
So if you're going to make selenium polyazides, your day starts with chlorine trifluoride and I'm sure that it just rolls along from there. Before you know it, you've gone from viciously reactive halogens, paused to prepare some disgusting selenium fluorides, made some violently unstable azides that explode if you stick your tongue out at them and hey, it's dinnertime already. . .
An early favorite has appeared in my “most alarming chemical papers” file for this year. Thomas Klapoetke and Joerg Stierstorfer from Munich have published one with a simple title that might not sound unusual to people outside the field, but has made every chemist I’ve shown it to point like a bird dog: “The CN7 Anion”. The reason that one gets our attention is that compounds with lots of nitrogens in them – more specifically, compounds with a high percentage of nitrogen by weight – are a spirited bunch. They hear the distant call of the wild, and they know that with just one leap of the fence they can fly free as molecules of nitrogen gas. And that’s never an orderly process. If my presumably distant cousin Nick Lowe does indeed love the sound of breaking glass, then these are his kinds of compounds. A more accurate song title for these latest creations would be “I Love the Sound Of Shrapnel Bouncing Off My Welder’s Mask”, but that sort of breaks up the rhythm.
These Bavarian rowdies have prepared a series of salts of the unnerving azidotetrazolate anion. As they point out, the anion was described back in 1939 (in what I hope was a coincidental association with the outbreak of the Second World War), but its salts are “rarely described in the literature”. Yes indeed! People rarely spray hungry mountain lions with Worcestershire sauce, either, come to think of it.
After reading this paper, I’m considering taking my chances with the mountain lions. The authors report a whole series of salts, X-ray structures and all, which range from the “relatively stable” lithium and sodium derivatives all the way to things that couldn’t even be isolated. In the latter category is the rubidium salt, which they tried to prepare several times. In every case, the solution detonated spontaneously on standing. And by “spontaneously”, they mean “while standing undisturbed in the dark”, so there’s really just no way to deal with this stuff. It’s probably a good thing they didn’t get crystals, because someone would have tried to isolate the hideous things. The cesium salt actually did give a few crystals, which they managed to pluck from the top of the solution and get X-ray data on. A few hours later the remaining batch suddenly exploded, though, which certainly must have been food for thought.
The authors went on to investigate the thermal behavior of these wonderful compounds, another risky move. As it turns out, they have calorimetry data on only five of the salts, because when they got to the sodium derivative, “a violent explosion destroyed the setup”. They also did sensitivity tests, using a standard drophammer rig from the Bundesanstalt fuer Materialforschung, evocatively abbreviated as BAM. These, along with the friction and spark tests, put these compounds well into the “primary explosive” category. Well, the ones that they could get data on, that is: the potassium and cesium compounds blew up as they tried to get them into the testing apparatus. So it’s safe to assume that they’re a bit touchy, too.
One of my favorite parts of the paper is the mention (found in much of the recent high-energy materials literature) that high-nitrogen compounds are worth investigated as “green” explosives, which makes me think that the whole environmental-rationale business must be reaching its end points. The notion of a more environmentally friendly way to blow things up aside, I have to salute the paper’s authors. They’ve made compounds that no one will have to make again, and survived the experience. Read the paper and be glad that this wasn’t your PhD project. . .
Now this is a fine substance. Also known in the older literature as fluorine azide, you make it by combining two other things that have already made my “Things I Won’t Work With” list. Just allow fluorine (ay!) to react with neat hydrazoic acid (yikes), and behold!
Well, what you’re most likely to behold is a fuming crater, unless you’re quite careful indeed. Both of those starting materials deserve serious respect, since they're able to remove you from this plane of existence with alacrity, and their reaction product is nothing to putz around with, either. The first person to prepare the compound (John F. Haller back in 1942) survived the experience, and made it (rightfully) the centerpiece of his PhD dissertation. But relatively few buckaroos had the fortitude to follow his trail over the years, and it’s not hard to understand why. Haller himself wrote on the subject in 1966 from an industrial position at Olin Mathieson, and got right to the point:
”(Fluorine azide) is described as a greenish-yellow gas at room temperature, liquefying at −82°C when diluted with nitrogen and freezing to a yellow solid at −143°C. Evaporation of this solid generally results in violent explosion.”
Yes, it does, and that does tend to slow down the march of science a bit. Not until 1987 was an improved procedure published, from Helge Willner and group in Hannover. (We'll see him again - most of his publication list falls into the "Things I Won't Work With" category, and I really have to salute the guy). Basically, it was the same reaction, but done slowly and Teutonically. You start off by making absolutely pure anhydrous hydrogen azide, which is a proposal that you don't hear very often around the lab, and is the sort of thing that leads to thoughts of career changes. (Maybe I could go into the insurance business and sell policies to whoever took over the prep). The next step is introduction of the fluorine, and when elemental fluorine is the most easily handled reagent in your scheme, let me tell you, you're in pretty deep. After the reaction, attention to painstaking fractional evaporation at very cold temperatures, in the best traditions of German experimental chemistry, is needed to clear out the reactants along with some silicon tetrafluoride, difluorodiazene, and other gorp. Willner's group managed to make about 20 milligrams of the pure stuff, but strongly recommend that no one ever make more than that. As far as I can tell, no more than a few drops of the compound have ever existed at any one time. This is not really a loss:
”The synthesis of pure N3F by the method described above was repeated more than 30 times without explosion. But if N3F is cooled to -196 C or N3F is vaporized faster than described, very violent explosions may occur. One drop of N3F will pulverize any glass within a 5-cm distance.”
They managed to get pretty full spectroscopic data on the compound while they had it, which was good of them, and even explored its chemistry a bit. Life must have a peculiar vividness when your job is to come in and see if triazadienyl fluoride does anything when you expose it to fluorine monoxide. (Oddly, they report that that reaction is OK – go figure). Still, most of the literature on this compound remains computational, rather than experimental (other than Willner's lab), and unless it turns out to be the secret to faster-than-light travel or something, that situation will continue to obtain. It's already good for accelerating Pyrex fragments past the speed of sound, but there are easier ways.
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.
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.
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. . .
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).
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 Chemistry28, 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.
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.
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.
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. . .
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.
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.