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
Here's an NMR imaging blog with details of a recent problem in an Indian facility. Two people ended up stuck to the machine, pinned by an oxygen cylinder (!) that one of them brought into the room. Both sustained injuries.
There are two questions here: one is how anyone is allowed to wheel a ferromagnetic metal cylinder anywhere near an NMR magnet, and the other is how it took so long to quench the magnet once the accident had occurred. That latter point is addressed by the blog link above - the hospital is saying that the emergency quench circuit malfunctioned, and that it took four hours for a GE technician to arrive and get things shut down. I'm no NMR hardware expert, but I wonder about that one myself. As that blog post concludes:
Whether or not GE was really at fault in Mumbai we shall learn eventually, I hope. (I have heard rumors that some sites like to bypass their quench circuit in order to avoid having the cost of recharging the magnet should the quench button get activated. Insert your own exclamations of disbelief here because I'm incredulous.) In the mean time, this sorry saga is an opportunity for all of us to review our own procedures and take the extra moments to ensure that we've done everything humanely possible to eliminate risks. There really is no excuse.
Note: this was a post on my old blog site, and never made the migration over to the current "In the Pipeline". I was reminded of it this morning, and thought I'd bring it more out into the light.
There are reports (updated here - DBL) that Mars may have hexavalent chromium compounds in its surface dust, which is already being brought up as a concern for future human exploration. I agree with comments I've seen that this is putting the cart in front of the horse a bit, but it also means that I probably wouldn't be a good candidate for the expedition. I've already had my lifetime's exposure to Cr(VI).
Back in grad school, I had an undergraduate assistant one summer, a guy who was pretty green. I'll refer to him by an altered form of his nickname, henceforth as Toxic Jim. I shouldn't be too hard on him, I guess: I was a summer undergrad in my time, too, and I wasn't a lot of help to anyone, either. But TJ did manage to furnish me with some of my more vivid lab stories in his brief time in my fume hood.
One morning I showed him how to make PCC. That's pyridinium chlorochromate for the non-organic chemists out there, an oxidizing agent that doesn't seem to be used as much as it was 15 or 20 years ago. Even in '85, you could buy it, but the freshly-made stuff was often better. It certainly looked nicer. Like all the Cr(VI) salts, it has a vivid color, in this case a flaming orange. I shouldn't say "flaming;" that's getting ahead of the story. . .
It's not hard to make. You take chromium trioxide, a vicious oxidant in itself which comes as clumpy fine purple crystals, and dissolve it in 6N hydrochloric acid. That's an easy solution to whip up, since it's just concentrated HCl out of the jug cut 1:1 with water. I had Toxic Jim do all this - weighing out the chromium compound, making the HCl. During that part I couldn't resist quoting the ancient adage, which works well in the East Arkansas accent of my youth: "Do like you oughter, add acid to water." Most chemists either remember that one, or they remember the syrupy conc. acids splattering all over their arm when they did it (once!) the other way around.
We set up a three-neck flask with an overhead stirrer to run this in. That's just a motor mounted above the flask, turning a shaft with a paddle on the end of it. Works well for really thick mixtures, which this was supposed to turn into. As things turned out, it was even thicker than planned, for a brief exciting interlude.
In went the HCl, out of a big Erlenmeyer flask, and in went the chromium trioxide. Here's where the wheels began to come off. Instead of a vivid red-orange solution, the stuff got dark and began to thicken. I could tell it was getting hot, too, since you could see the clear wavery solvent vapors coming out of the open necks of the flask. And that was wrong, too - you don't get that so much with water vapor. It's the mark of organic solvent fumes, with their different density and refractive index.
And so it was. TJ had indeed grabbed the wrong Erlenmeyer. Not the one he'd just mixed up the HCl in, but one from another part of the bench that contained ethyl acetate from a big chromatography run the night before. Ethyl acetate is a pretty poor substitute for hydrochloric acid, most of the time, when you stop to think about it.
Then the overhead stirrer began to bog down, which takes a mighty thick mixture to achieve. I hadn't added up what had happened at this point, but I knew that things were going wrong in all directions at once. I pulled the glass hood sash down some more, saying "I think you better stand back -" WHOOOOMPH!
And there it went! The whole reaction went up in a big fireball, which filled a good part of the hood and came roaring out of the gap in the front sash. I felt the heat roll over me, yelled something incoherent, and bolted for the safety shower. I didn't have to run up Toxic Jim's back, either: he was making for the door in championship time. Pulling the chain of the shower dumped a hundred gallons of ice water on me immediately, not that I needed any more waking up.
When I opened my eyes and took inventory, things weren't as bad as I thought. Limbs and appendages all present, head and facial hair still attached - though lightly singed and frizzed - skin not even sunburnt, although it (along with my lab coat) was generously splattered with green. That was what remained of the chromium trioxide. It was now the Cr(III) oxide, having given up three oxidation levels by turning the ethyl acetate into carbon dioxide, most likely. There were a few orange-brown spots of the Cr(VI) stuff, but those were mostly confined to the front of the lab coat, in a vivid line that showed where the hood sash had gotten pulled down to.
My hood wasn't looking its best. There was smoke hanging in the air, although that was getting pulled out. There was a huge stain of the green and brown chromium mixture all over the inside, thickest in the directions of the three open necks of the flask. Which was still intact - if I'd been foolish enough to set this up in a closed system, the whole thing would have gone up as Pyrex shrapnel. Even the ceiling had a line of gunk on it, from the thin gap in the hood sash assembly.
While I was taking this in, wondering what the hell had gone wrong, and wondering what I could possibly do to TJ that was worse than what he'd just gone through, the emergency crews arrived. It was a Saturday morning, but Bob across the hall saw the explosion and immediately dialed 911. In came the fire crews, trying to talk through their breathing apparatus: "Mumph heff deff umphh cafulteff. . " "What?" "We hear there's a casualty up here"
I put my hands on my hips, and gave them the full effect of my green spots, frizzed hair, and soaking wet lab coat: "That would be me."
I have this from a lab-accidents-I-have-known discussion over on Reddit. It is, of course, unverified, but it's depressingly plausible. As a chemist, this one is guaranteed to make you bury your head in your hands - it's the second law of thermodynamics come to take vengeance, with the entropy increasing as you go along:
"A graduate student was constructing three solvent stills (dichloromethane, THF and toluene) inside a hood in Room XXXX. As a final step in this process, the student was cutting pieces of sodium metal to add to the stills. Once the sodium had been added, the student began to clean the knife used to cut the sodium. During the cleaning, a small particle of sodium was apparently brushed off the knife. The sodium landed in a drop of water/wet spot on the floor of the hood and reacted immediately making a popping sound. The graduate student was startled by this sound and moved away quickly.
In his haste to get away from the reacting sodium, he discarded the knife into a sink on the bench opposite the hood in which he was working.. Apparently, there was another piece of sodium still adhering to the knife since upon being tossed into the sink, a fire ignited in the sink, catching the attention of another student in the lab. As the flames erupted, the student noticed a wash bottle of acetone sitting on the sink ledge nearby. He immediately grabbed it to get it away from the flames, but in the process, squeezed the bottle, which squirted out some acetone which immediately ignited. The student immediately dropped the bottle and began to evacuate the lab. As he turned to leave, he knocked over a five gallon bucket containing an isopropanol/potassium hydroxide bath which also began to burn. This additional fire caused the sprinklers to activate and the fire alarm to sound which in turn resulted in the evacuation of the building.
When the sprinklers activated, water poured into the bulk sodium-under-mineral-oil storage bottle which had been left uncapped in the hood resulting in a violent reaction which shattered the bottle sending more sodium and mineral oil into the sprinkler water stream. This explosion also cracked the hood safety glass into numerous little pieces although it remained structurally intact. By the time the first-responders arrived on the scene, the fire had been extinguished by the sprinklers, but numerous violent popping sounds were still occurring. The first-responders unplugged the electrical cords feeding the heating mantles, shut off the electricity to the room at the breaker panel and waited until the Fire Department arrived. Eventually the popping noises stopped and sprinklers were turned off. The front part of the lab sustained a moderate amount of water damage The hood where the incident began also suffered moderate damage and two of the three still flasks were destroyed. The student, who was wearing shorts at the time of this accident, sustained second and third-degree burns on his leg as a result of the fire involving the isopropanol base bath.
There were some additional injuries incurred by the first-responders who unexpectedly slipped and fell due to the presence of KOH from the bath in the sprinkler water. These injuries were not serious but they do illustrate the need to communicate hazards to first-responders to protect them from unnecessary injury."
I doubt if the sodium was being added to the dichloromethane still; I've always heard that that's asking for carbene trouble. (Back in my solvent-still days, we used calcium hydride for that one). And it would take a good kick to knock over a KOH/isopropanol bath, but no doubt there was some adrenaline involved. I'm also sorry to hear about the burns sustained by the graduate student involved, but this person should really, really have not been wearing shorts, just as no one should in any sort of organic chemistry lab.
But holy cow. The mental picture I have is of Leslie Neilsen in a lab coat, rehearsing a scene for another "Naked Gun" sequel. This is what happens, though, when things go bad in the lab: we've all got enough trouble on our benches and under our fume hoods to send things down the chute very, very quickly under the wrong conditions. And were these ever the wrong conditions.
Via on Twitter (and that via C&E News), I bring you the definitive what-are-we-going-to-do-with-all-this-sodium video. The end of World War II brought all kinds of material disposal problems - you may have seen footage of virtually new airplanes being dumped into the sea and the like. Some of those disposal problems are still with us, like the unexploded ordnance that keeps turning up. But these barrels of sodium, no one ever had to worry about them again. . .
Via Sally Church on Twitter (and a post by Bethany Halford at C&E News), I bring you the definitive what-are-we-going-to-do-with-all-this-sodium video. The end of World War II brought all kinds of material disposal problems - you may have seen footage of virtually new airplanes being dumped into the sea and the like. Some of those disposal problems are still with us, like the unexploded ordnance that keeps turning up. But these barrels of sodium, no one ever had to worry about them again. . .
Well, since I was just talking about a reagent that can potentially take off without warning, I wanted to solicit vivid experiences from the crowd. What's a compound that you've made that did something violently unexpected? I can recall making some para-methoxybenzyl chloride in grad school (for a protecting group; I was running out of orthogonal protecting groups by that time). It's not hard - take the benzyl alcohol and some conc. HCl and swoosh 'em around. But the product you get by that method isn't the cleanest thing in the world, and on storage, well. . .a vial of it blew out in my hood after the acid had had a chance to work on it.
My most vivid reagent-gone-bad story is probably this one; that's a time I literally came down counting fingers. What other things have you had turn on you?
What do you have when a fire starts at a large chemical packing company, handling all sorts of oils, paints, coatings, and various industrial chemicals? Where they have hundreds of thousand-liter containers stored, surrounded by all the crates and packing material used to trans-ship them? You have this, at Chemie-Pack in the Netherlands yesterday:
And you have a black cloud that stretched across a significant part of the whole country:
Images are from Nufoto.nl, taken by people at the scene. A reader who lives 20km downwind writes me that he's been getting a pervasive smell of burnt plastic (which, he says, certainly makes a change). His main reason to be grateful is that this didn't spread to the Shell site nearby, which would have prompted an instant vacation to Germany. And then there are all the refineries 15 km to the west - if those ever go up, he tells me, "it'll look like the ending of Gulf War 1 - lowlands-style - with cows for camels".
Culturing bacteria is usually a pretty quiet affair. Bacteria aren't too noisy, and the equipment used to keep them happy isn't too dangerous. But there are exceptions. If you're going to culture anaerobes, you need somewhat more advanced technique, what with all that oxygen-is-deadly business. A professional-grade culture chamber for those beasts is usually filled with a mixture of about 80% nitrogen, 10% carbon dioxide, and 10% hydrogen. And those you'll be getting from three compressed gas cylinders, which is how they were doing it in a lab at the University of Missouri until Monday afternoon. . .
Well, regular readers will be expecting this to be a story of someone who did not remember to Treat Compressed Gases With Respect, but that's not the case. No, this is what happens when you don't Treat Hydrogen With Respect - and everyone in the audience who's had a hydrogenation reaction get frisky on them will be nodding their head in agreement at that thought. Somehow, enough hydrogen and enough oxygen got together around an anaerobic culture hood, and the mixture found an ignition source, and well. . .
Problem is, just about any hydrogen/air mixture will do. Anything from about 4% hydrogen in air to about 75% will ignite, and everything except the two ends of that range will go ahead and explode if given the chance. (Only acetylene is worse in that regard). And it doesn't take much to set it off, either, which is the other nasty thing about working with hydrogen. A static-electricity spark is plenty, as are the sparks generated by electrical switch contacts and the like.
As you can see, the lab was not improved by the resulting explosion. The latest report I have is that four people were injured, one seriously enough to still be in the hospital, although their condition has been upgraded to "good".
Initial reports were that this was due to human error, although everyone seems to be backing off that judgment until an official investigation is finished. At any rate, the local fire department stated Monday night that the problem was one or more people in the lab "not being familiar with the warning systems designed to alert them when the hydrogen level was approaching explosive limits (and) the gas was left on". If that was the case, then. . .you ignore a hydrogen level alarm at your peril. And here are seventeen blown-out windows, four people who are lucky not to have been killed, and one demolished lab as evidence.
Update: I had a link up to a commercial anaerobic culture chamber for illustration, but (as the manufacturer points out) these use cylinders of premixed gas with only 5% hydrogen which obviates this very problem. I thought it best to take down the link so that no confusion results - after all, it wasn't the model that was being used in this incident (and in fact would have avoided it completely). I should add that the email I received about this out was exactly the sort of courteous and informative request I have no problem responding to, as opposed to some others that have come in over the years.
(Photos are courtesy of the Missourian and the Columbia Fire Department).
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.
For Friday lunchtime, I have a brief but alarming video clip from a 2007 incident in Dallas, where a fire started at a company supplying industrial gases to welding shops and the like. The incident was heralded, like so many others, by the simple but meaningful phrase "I hooked up something wrong". This as smoke began to emerge from the bed of a delivery truck full of aceylene cylinders.
If there's one thing to be learned from the whole "How Not to Do It" category on this blog, it is to treat pressurized gas containers with respect. Roasting them over an open fire does not qualify.
In case anyone missed it, a commenter on this post unearthed a really extraordinary find in the chemical literature. Here's an obscure isolation paper, from an obscure Chinese journal, reporting on a profoundly boring list of marine natural products.
What's so great, you ask? Well, take a look at the list. Dum de dum. . .hold on a minute, bis(2-ethylhexyl) phthalate? From Streptomyces, you say? When it's one of the most common plasticizers in the world, a bulk industrial chemical that, well, notoriously leaches out of labware under solvent exposure? Sure thing, guys. Sure thing.
We haven't had a How Not to Do It around here in a while, so here's a companion piece to the famous Sealed-Up Liquid Nitrogen Tank. This incident happened (as far as I can tell) about ten years ago. It's been used in a number of safety presentations then, thanks to the Airgas Corp., whose safety officer assembled a number of photos (and this is the time to emphasize that they had nothing to do with the accident itself, because people who work for a pressurized-gas company actually know how to handle pressure vessels.
As opposed to the two guys who scavenged a liquid oxygen Dewar from a scrap metal yard and decided to put it back into service. According to the most detailed report, they tried to rig up a connection to refill the cylinder, but found that it vented immediately through the pressure-relief valve. So. . .well, yeah, you know what's coming next: they took the darn thing off and plugged it shut. No more pesky venting! They filled up their cylinder, which was loaded on the back of their pickup truck, and went rolling down the interstate at lunchtime. Whereupon they had a flat tire, and pulled over for a while to fix things. . .
OK, you can look out from behind your hands now. Although I can't imagine how, neither of these two cowboys managed to get themselves killed, nor did they take out anyone else, through what appears to be sheer blind luck. According to the report, one member of the Cylinder Kings ended up being blown across five lanes of traffic, while his partner was launched forty feet in another direction. You can see from the photo how the truck weathered things. I can't imagine that a pressure wave of straight oxygen hitting tank of gasoline can end well; it's a perfectly reasonable mixture to put a payload into low-earth orbit.
Which is a good note on which to take inventory here. We have the owners of the oxygen cylinder accounted for, and their truck. What about the cylinder itself? Well, similar to the nitrogen tank referenced above, it had failed at the bottom weld and thus departed the scene of the accident like an artillery shell. It re-entered the affairs of the world a quarter of a mile away, plunging through the roof of an apartment, completely trashing the place (and severing a natural gas line in the process). As I said, how a dozen people didn't end up killed by all this is a complete mystery to me. (The red circle in that photo is where the pressure-relief device used to be. )
So the moral of this story is, I suppose, that Pressure Relief Devices Are There For A Reason. Or maybe it's "don't scrounge gas cylinders from the scrap yard and try to get them to work". Or perhaps "just because you haven't seen a pressure vessel explode yet, it doesn't mean that they can't". Or "Gegen der Dummheit kämpfen Götter selbst vergebens." Or something.
Some blogs run pictures of cats to give the readers a break from the ordinary. Around here, I thought that this might be appropriate. Here are the alkali metals, from top to bottom, differentiated in the most basic way possible. No, not by tasting them, sheesh: by tossing them into a dish of water:
(Courtesy of the Open University site in the UK). One thing they don't go into is the effect of density. Up to potassium, the metals are still light enough to float. But cesium drops like the rock it is, with depth-charge results.
I will consider running a photo of a cat, as long as he's working up a reaction.
Readers may remember the incident a couple of years ago where a paper was published claiming the synthesis of some very odd-looking 12-membered ring compounds. Prof. Manfred Christl of the University of Würzburg noticed something odd about this reaction, though, namely that it had already been run over a hundred years ago and was known to give a completely different product. (As I pointed out here, though, you didn't need to unearth the ancient literature to know this; ten minutes of looking through the modern stuff would have done the trick, too).
Well, Christl's back with another takedown of some improperly assigned weirdo 12-membered rings. This time, it's Cheryl Stevenson of Illinois State that gets the treatment, with this paper from last year that claims several interesting ring structures from 1,5-hexadiyne and base. Christl had trouble believing the mechanism, and on closer inspection had trouble believing the NMR assignments. Then, on even closer inspection, he assigned the structure as a simple isomerization of one of the triple bonds, and found that this exact reaction (and product) had first been reported in 1961 (and several times afterwards). Not good.
As it turns out, I almost certainly made some of the compound myself, by mistake, back in mid-1983. That was the summer before I started my first year of grad school, and I was doing work in Ned Porter's lab at Duke. One of the starting materials I needed was. . .1,5, hexadiyne, which you couldn't buy. So I made it, in real grad-school fashion. I homocoupled allyl Grignard to get the 1,5-hexadiene (which even if you could buy back then, we didn't). Then I reacted that with bromine and made the only six-carbon molecule with four bromines on it that I ever hope to make. Reacting that with freshly prepared sodium amide in ammonia gave the smelly di-yne, in crappy yield after distillation. I can still see it: me heating up a column full of glass beads, then turning to Steve, the postdoc in the next hood, and making a bad joke about Herman Hesse while David Bowie's "Modern Love" played on the radio. . .ah, those days, they will not come again.
At any rate, I went on to react the compound with bases under different conditions, trying (in vain) to alkylate both of those terminal alkynes, and thus passed the last of my summer, in exactly the way my two previous summers of research had passed: unsuccessfully. This latest paper, though, makes me think that I was probably turning my starting material instead into exactly the diene that Christl is talking about. I should have hit the library harder myself, although (to be fair) there are references that tell you that you can do that alkylation, and digging through the literature was a good deal more time-consuming back then that it is now.
That lab accident, you say? Well, that happened when I was making a big batch of sodium amide. You start that prep off like a Birch reduction - condense a bunch of liquid ammonia into a flask, and start chucking sodium metal into it. The big difference is that you add a bit of ferric chloride to the mix, which kicks things over at the end. After you've dissolved your sodium, to give you the bronzy purple-blue of solvated electrons, you take the flask out of the cold bath to let the ammonia reflux. At that point, the whole thing suddenly clears, dramatically revealing grey lumps of sodium amide rolling around on the bottom of a pond of plain ol' ammonia, without a solvated electron in sight. (I have, in years since, seen a couple of people refer to the blue stage of the reaction as sodium amide, which it ain't, and I can get quite cranky and pedantic about it).
One afternoon I was whipping up a batch of this stuff, when something starting going on inside the flask. I don't recall what made me take a look at the ammonia solution, but since there was so much bronze gorp on the side of the one-liter three-neck, I had to lean in and look down near the central joint. Whereupon my hair wound itself immediately around the greased shaft of the overhead stirrer, pulling my head in toward the whole setup and jamming my nose into the side of the flask. I fumbled for the switch of the stir motor, feeling like George Jetson as I shouted for someone to give me a hand, and watch with interest as the dry ice bath bubbled along an inch away from my face.
Steve the postdoc came to my aid, shutting off the grinding motor which was doggedly trying to wind me headfirst around the stirrer shaft. We unreeled my hair from the whole contraption, with me cursing foully and Steve merrily making jokes of the "Hair today, gone tomorrow" kind, with side comments about me getting too wrapped up in my work. Those days, as I said, will not come again.
Patent applications are no fun to write. You have to figure out just what you're trying to cover (and how wide a space around it you want to try to clear), and the lawyers have to whip up language that casts just the right legal spell. The chemists have to write up detailed experimental procedures for all the important compounds and procedures, gather the matching analytical data, and make sure that it all fits together. Just getting the numbers assigned to the compounds right (and keeping them right through all the revisions) is a tedious job in itself. You always have to go through more drafts than you thought. No one enjoys it.
So maybe it's not surprising that things sometimes, well, slip a little. But how about when they slip a lot? Take this morning's example from Merck (a company that pitchforks out patent applications by the pile). Their WO2009091856 just published recently, directed at bicyclic beta-lactamase inhibitors. And everything looks normal for quite a while - 120 pages or so, in fact. Then the text suddenly snaps into bold face, and an authorial voice makes itself heard:
It appears that the data in Table 3 were generated in the same manner using the same enzymes as in Table 2 (Table 2 is unchanged from the provisional filing. I plan to DELETE the entries in Table 3 for Ex. 2,6,7 and 8 because this data duplicates the data in Table 2. . .Also, the entries in Table 2 for Ex.8 are both "1.6", not "16" as shown in Table 3. Please clarify these differences. . .
Is the data shown for Ex 1A data generated in a separate run, or is it supposed to be the same as for Ex. 1?. . .You don't want to include synergy data for these compouds (sic). It would be helpful to include it, at least for some of the examples (could put in a separate table). Recommend we include it for Ex. 14, since this is a likely backup candidate.
Now that's not supposed to be in there! What you're reading are the comments of someone in Merck's legal department - the sorts of comments that every patent draft collects as it's written, the sorts of comments that are supposed to be excised before you send in the application. Not this time! So if you were wondering which compound in this application represents the real candidate, and which ones are the backups to it, well, wonder no more. That query about including synergy data, for example, is an attempt to make it harder to figure out the most preferred compound itself - in vain, as it turns out. Oh, and those corrections that the comments say should be made? They weren't. So you'll want to fill in the correct numbers yourself.
That sort of thing goes on all the time in patent writing. You have to disclose your best compound - and in fact, you have to "teach toward" it in the claims. But you don't have to spray-paint it orange, and there's no sense in making things easy for your competition to figure out. A careful analysis of a patent application's claim structure will narrow down what a patent's authors are really interested in protecting, but there's often still some doubt about which exact compound is the winner. There can be other clues. Sometimes it'll be the compound with the most extensive biological characterization, or sometimes, if you look through all the experimental procedures, you'll notice that everything's being done on 50-milligram scale until one prep jumps to twenty grams. Bingo! Careful preparation of an application can scrub most of this stuff out. But all is for naught if your legal team's strategy comments are included in the Special Bonus Director's Cut version of the application.
Oh well, bonus dormitat Homerus. Anyone who's interested in beta-lactamase inhibitors (which, I should add, I'm not) now has more data to work with. The patent analysts at Thomson-Reuters are the people who caught this mistake (a colleague forwarded their writeup on to me). As they note, the wayward legal paragraphs also mention the possibility of comparing compounds to "MK-8712". That MK designation is Merck's usual method of showing a compound that's been recommended for the clinic, but this is the first that anyone seems to have heard about this one. But we can be pretty sure of something: someone in Merck's legal department has had a very bad day of it within the past couple of weeks. . .
You need access to vacuum if you’re going to work at the bench in chemistry. In fact, you need more than one kind. Reasonably hard vacuum (well, by our standards, which is laughable by the standards of the physicists) is down in the single Torr or below – that is, less than about 1% of normal air pressure. We use that for pulling out residues of water or organic solvents from our compounds. You can’t usually see it happening from the solid ones, but the syrupy liquids will foam up or blow a long series of thick bubbles when the vacuum is applied. The foam can be an irritating problem at times; some things will fill your flask with sticky bubbles and go right on up into the vacuum line if you’re not watching them.
The lesser vacuum lines are used for bulk evaporation of solvent (on your rotavap) and for filtering things off. We do an awful lot of both of those, too, and a full vacuum-pump pull is too vigorous for them in most cases. Evaporating down reactions is a constant task in an organic chemistry lab; I’d rather not think about how much of it I’ve done over the years. As for filtration, there are many cases where a solid product can be filtered out of the bulk liquid (which is good) or where some undesired solid by-product has to be filtered out before you can go on (not as good).
The low-tech way to get the sort of pull-it-though vacuum you need for these things is a water aspirator. You don’t see these as much any more, and you don’t see them at all in industry, since they necessarily pull solvent vapors into the water stream. But they work. An aspirator is basically a narrowing tube that hooks up to a hard-spraying water tap and has a sidearm fitting. The accelerating blast of water pulls the air in the tube along with it as it goes, creating a useful vacuum. If you wanted to make one rather more environmentally friendly, you’d keep a well-stocked dry ice condenser in line with it to trap out the solvent vapors before they go down the drain (which is what your rota-vap should have on it, anyway), but even with that, you’re always going to be turning the water flow into a waste stream. As I say, you don’t see them as much these days.
But we used them back when I was in grad school, that’s for sure, mostly for the rotavaps. If you wanted to keep things from splashing around back in your hood, you attached some rubber tubing to the other end of the thing and ran it further down the drain a bit.
Well, one day, one of the guys in the lab next door to me was shocked to see water blasting around in his hood. It was a real fountain, just geysering out full blast from what must have been a cracked water line or something in the back. He ran over and immediately shut off every tap – but to no avail. Roaring, showering water everywhere. Getting a look at the source, he realized, to his consternation, that the water was coming up out of the drain in the back of his hood. I remember standing there with him, staring at this in disbelief. It looked like a special effect. How on earth could you get water blasting up out of a drain pipe?
Suddenly it hit me. I ran around to the other side of the lab, where a new Japanese post-doc had taken up residence. “Masa”, I asked him, “Did you just put that rota-vap in your hood today?” “Yes, yes, just started it today”. There was a water aspirator flooshing away back in the back of his hood. “Did you put some rubber tubing on that thing?” “Tubing? Oh, yes” “How much?!” “Whoaaa. . .” He spread his arms to indicate the mighty extent of the rubber tubing he’d added.
Mighty, indeed. He’d run the stuff down his drain, through a horizontal pipe and right through a T joint, and back up out of the drain of the other guy’s hood, which backed on to his. So when he turned his water on full throttle, he immediately started irrigating his labmate’s space. We finally go thing turned off, and trimmed back the rubber tubing to a more reasonable length (like, not seven feet), and order was restored. For a while.
Note: if you want to see How Not To Do It to a really expensive vacuum rig, try here.
This post will have one of those stories that I can’t vouch for personally, and I’m very glad of that. It involves making diazomethane, which will have already gotten the attention of the chemists in the crowd.
Diazomethane’s a very useful reagent, but it has to be treated the right way. You can’t buy it – no one will ship the stuff – so you have to make it fresh. (There are several such reagents). For many years there have been chemicals in the catalogs whose only real use has been to generate diazomethane when needed. Generally this involves treating some nasty N-nitroso compound with base in ether, then distilling over the ether solution of the reagent, which is a distinctive bright yellow.
There’s where some of the trickiness comes in. That diazo group is looking for an excuse to revert back to nitrogen gas, which process comes with an inevitable no-substitutions side order of kaboom. The chemist’s job is to not give it that excuse. That means that you can’t heat the stuff up, you don’t make it very concentrated, and you don’t even expose it to sharp or rough surfaces, because that can be enough right there. They sell distillation glassware specifically for diazomethane preps, with weirdly glossy ground-glass joints.
You can keep your yellow solution stockpile in the freezer for a while, and the temptation is always to make a lot of it so you never have to do it again. That leads to phenomena like the big flask of the stuff left behind when someone leaves the grad school group. One of those surprises (“Is this yellow stuff what it looks like it is? How long has it been in here? And who the hell made it, anyway?”) was the cause of a new lab inspection requirement while I was getting my degree. You couldn’t leave until someone determined that you weren’t passing on any explosive bequests.
Of course, sometimes you honestly need a lot of these things. One of the guys in my group was in that situation early in his total synthesis. One summer afternoon, the power went out in the labs during a thunderstorm, and the head of our safety committee came rolling a big cooler of dry ice down the hall. “Anybody need to store something in the cold?” was the call. “Well,” I said, “we’ve got a couple of liters of diazomethane solution.” “That’s not very funny,” he said. “That’s because it’s not a joke”, I replied, and we moved to the front of the line.
So, what’s the stupidest way to handle the stuff? That’s the story told to me by a colleague. He attests that when he was in grad school, he looked across the hall to see someone involved in making a goodly amount of diazomethane – in a large standard ground-glass-joint apparatus. Oh, dear. How the guy was going to get his collection flask off without running the risk of grenading everything, that was the question. As my friend watched in disbelief, the guy reached up to just twist the darn thing right off. . .and it was stuck. A frozen joint – just the perfect time for it. (This is the point where the audience for this story began to bury their heads in their hands).
My colleague swears that he then watched this maniac pick up a propane torch to sweat the joint loose. I believe that someone may have stopped him in time, but I think the teller of this tale decided to adjourn for lunch at some distant location right around then, so I can’t vouch for the outcome. But if anyone has a more drooling, slack-jawed approach to an ether solution of diazomethane than running a propane torch over it, I’d like to know what it is. Short of maybe using it as an HPLC solvent, I’m out of ideas.
I was mentioning chromatography last week, and I’ve been running several columns this week myself. I’m doing them the new-fangled-with-sufficient-funds way, which has been the standard in the drug industry for many years. You buy the columns of silica gel pre-packed and plug whatever size you need into a machine. Then you load your sample on, tell it what solvents you want to use, put in a rack of test tubes, hit the button and go do something else.
You can set the machine to collect all the solvent that comes off, or (if your compound absorbs ultraviolet light) to only collect when something UV-active starts coming out the other end. Twenty minutes or so later, you come back to a rack of fractions and a printed map showing which UV-active peaks are in which tubes. All this is very nice, and would have caused me to faint with desire if I’d seen it when I was in grad school – not that I could have, since plug-and-play systems like this weren’t on the market back then.
The standard way was (and is, in less well-funded environments) to slurry up your silica gel in solvent, pour that into a glass column, push the solvent through with a stream of air or nitrogen pressure from the top (usually holding it down by hand to keep things from getting out of control), loading your mixture, and eluting fractions into test tubes or Erlenmeyer flasks until you’re sure that your stuff is all out. I wouldn’t want to guess how many columns I’ve run by hand like that over my lab career, but it’s been several years since my last one, and I don’t see another one in my future, with any luck. (For those of you who want to see how it's done, and have twenty minutes to spare, the folks at MIT will tell you all about it).
It’s hard to mess up the automated systems, although you should never underestimate the ingenuity of the user base. But the hand-run columns can easily be loused up in all sorts of ways. Perhaps the most spectacular I’ve seen was when the guy across the hall from me in grad school, who I’ll call “Bob”, since that was his name, decided to run a big column using DMSO as the solvent.
Most of my chemistry readership will have just looked at the screen and said “He did what?” DMSO is a mighty odd choice for a chromatography column. It’s a strong, strong solvent, for one thing, and would mostly just be expected to dissolve everything and sweep it right out the other end. And it’s thick and viscous, too, compared to the solvents that reasonable people use, which means that it would be no fun to get it to come out that other end at a reasonable rate.
But that was Bob’s choice, and he was working on a bunch of nasty, insoluble stuff, so DMSO seemed like a good idea to him at the time. But he didn’t run his column as I’ve described above. He was of another school of setting up columns – apostates, if you ask me – which advocated packing the silica gel into the column dry and running solvent through it before loading the sample. (That always seemed to take longer and use up more solvent, as far as I could tell).
It was a particularly ill-suited method for running a big honking DMSO column. DMSO, as you’ve probably never had the chance to notice, has rather exothermic solvation behavior with silica. In non-PhD language, it gets very hot, very quickly, when it wets the dry powder. So when Bob started, against all odds and a lot of common sense, to force a big bolus of DMSO down his dry column, things shortly got out of hand. Next door, I heard a big “POW!” and ran over to see what it was this time.
There was Bob, staring with dismay at the remains of his column, which had cracked and spewed DMSO-soaked silica all over the bench. In retrospect, he’s lucky that it didn’t shrapnel all over the place. As it was, he had an awful mess to clean up. I never got around to asking him just what he was going to do with all the DMSO fractions he would have taken off the column - evaporating that stuff off is no joke, although it's another problem that yields to sufficient funding. But, then, if sufficient funds had been available back then, Bob never would have been running a column in DMSO in the first place. . .
I’ve written before about all the fun you can have in a lab with compressed gas cylinders. We use the things all the time in chemistry, but as pieces of apparatus, they can only be pushed so far. The problem is that they demonstrate their unhappiness by venting great quantities of stuff that you’d rather not breathe (if you’re lucky), or by taking off like an unguided missile and punching holes in the walls and ceilings (if you’re not). That latter behavior is flat-out guaranteed to show up if you abuse them – for one of the more spectacular examples, see here.
I’ve never had one take off on me, fortunately, but I haven’t always stuck to the straight and narrow with the things, either. My worst behavior has usually been with lecture bottles, the dilettante-sized gas cylinders that bench chemists often use. Most chemistry departments have a few of these sitting around, generally charged with foul reagents that are needed every three or four years or so. Sulfur dioxide, boron trifluoride, phosphine – that’s the sort of thing. They’re low-use almost by definition. If there’s a regular need for a gaseous reagent, you buy larger cylinders of it, because lecture bottles are by far the most expensive way to go.
In graduate school, I was setting up some Prins reactions, which take some sort of acid component to make them run. If you use an aqueous one, you generally get an alcohol out of them (from water picking up the final cation), but if you go anhydrous you can get all sorts of other compounds. I needed a bromide, so anhydrous hydrogen bromide it was.
We had a fairly crusty lecture bottle of it around, and I eventually located some dubious-looking small regulator valves. I picked the least-corroded looking one and screwed it on. Lecture bottles have a main metal-faucet style valve up at the top, like all gas cylinders, and once you open that it’s up to the regulator valve to stop things down to manageable flows. I had my reaction set up, so I worked some tubing onto the thing and had a go at opening it up.
No dice. Boy, was that thing tight. I reached into the hood and wrestled around with it, to no avail. I took it out and got a better grip in another part of my hood, away from my reaction setup. Tighter than two quarts of fresh frogs in a half-pint pickle pot, as Walt Kelly once put it – the valve wouldn’t budge.
I’ll skip over a couple of intermediate stages and cut right to the final scene, which is me clutching the darn lecture bottle to my chest, hopping around the lab grunting and cursing as I put all my strength into trying to force the stupid valve open. You won’t see a pictograph of that method in the instruction booklet, I’m pretty sure. A recording of what happened next would have gone something like this: “Urk! Unk! Ark! WHOA!”
The valve opened, finally responding to my Conan-the-Barbarian technique, and the cheap regulator then hissed out an orange cloud of gaseous HBr right into my shirt pocket. Not a good storage compartment, actually. I whooped, shut the valve, laid the gas cylinder down fast and stripped off my shirt where I stood. You haven’t really lived in a lab until you’ve taken off your clothes in it, I always say. I staked my claim to this one by standing in it shirtless, splashing saturated sodium bicarbonate on my chest, and glaring at the remains of what used to be one of my favorite shirts.
Man, have things changed since I was in grad school. We used to pour all kinds of horrible things down the drain - mind you, this was a good twenty years ago. But you can't do that now, can you?
A respected University of Washington pharmacology professor became a felon Wednesday when he acknowledged dumping a flammable substance down a laboratory sink and then trying to conceal his actions.
Daniel Storm, 62, pleaded guilty in federal court to violating the Resource Conservation and Recovery Act by flushing about four liters of the solvent ethyl ether. He faces a maximum five years in prison and a $250,000 fine when sentenced June 18, although prosecutors have recommended probation under the terms of a plea agreement.
Well, everywhere I've worked, the safety officers have tried to put the fear of RCRA ("rick-rah") into us, and by gosh, it looks like they may have had a point. Turns out that Prof. Storm's lab had several elderly containers of ether which turned up in a lab inspection, and he decided to get out of paying the $15,000 hazardous waste disposal bill. So he decided to take matters into his own hands.
And how: he went after the metal ether cans with an ax, which means that he was lucky not to blow himself up. (A stray spark from the metal could have done the trick, and who knows how much peroxide was in the stuff, for that matter). Why the Monty Python lumberjack routine? Well, the lids were too tight, and according to Prof. Strong, the ax just happened to be handy. (How many times have the police heard that old excuse, eh?) Yep, you can't pour ether down the sink like we used to, and you can't chop open the stuff with an ax like we. . .well, actually, we never used to do that. No one ever has, most likely.
What really ripped it was when he went on to fake paperwork from a nonexistant waste disposal company to make it look as if the ether had been properly hauled away. No, if you haven't clicked on that link yet, you'll have to take my word that I'm not making this up as I go along. But you get the impression that Professor Strong sure was. Makes you wonder if he had been exposed to too many fumes. A spokeswoman for the school says that she's unware of any similar incidents there, and I'll bet she's telling the truth. No, I've seen some stupid things done with diethyl ether, but this one threatens to retire the trophy.
Over here at scenic Lowe Manor (otherwise known as the House that Pharma Has Paid And Will, With Any Luck, Continue Paying For), the dinner table conversation sometimes runs to things like the proper handling of flaming t-butyllithium. Well, OK, the conversation is a bit one-sided, since I'm the only one in the house who's used the darn stuff. My wife has done a lot of bacteriology and molecular biology, fields that don't find much use for pyrophoric organometallics, and I'll keep my kids a good distance from any bottles of butyllithium, thanks.
But I was speaking to the them the other night about the value of experience. Tertiary butyllithium catches on fire, and there's nothing you can do about it. Your best course is to be aware of that, and to expect it. That way, you won't be shocked when you put your syringe into a bottle of the stuff and withdraw it only to find a merry orange flame burning from the tip of the needle. That's a good sign - it shows that your bottle of butyllithium is still good, as opposed to the cloudy, wimpy, hydrolyzed stuff that you should carefully leave sitting in someone else's hood when they're not around.
This pilot light will do you no harm, and will extinguish itself once you put your syringe needle through the next rubber septum. But if you're not expecting it, the sight can come as a bit of a jolt. The consequences are generally not good. There's almost always a tensing of the hand and arm muscles, which tends to depress the syringe plunger a bit, and whooomph: instant flamethrower. I've heard of several completely needless fires that started this way, invariably at the hands of someone who wasn't psychologically prepared to wield some (temporarily) flaming lab equipment with the needed aplomb.
As I mentioned here before, I've had still more practice with fiery glassware. I can attest that a butyllithium flame from the pure substance has a more noble purple color to it than the common orange of the commercial hexane solutions. That magenta hue is from the emission spectrum of lithium itself, and (at least in my case) it did not have a calming effect.
There are, regrettably, even more stupid things to be done with t-BuLi. I'm not sure if I've told this one here before - it's not in the archives at right, so here goes: a friend of mine in grad school was showing a summer undergrad student (hear the chairs of the experienced chemists draw closer) how to do cannula transfers of air-sensitive materials. (This involves hooking up a hose system with needles and tubing, with dry nitrogen or argon gas bled in at the front end of the system to force the stuff over into another waiting flask). There was a double manifold set up in the hood, as usual, to allow a vacuum pump out to remove air from reaction flasks and let nitrogen in to replace it. Somehow, this summer student got the vacuum and nitrogen setting all hosed around when trying to cannulate a whacking load of t-BuLi, and reported back a few minutes later that (although everything was set up perfectly) no butyllithium was appearing.
Feeling the hair raise up on his arms, my friend came to look things over, and saw that indeed, no t-BuLi was showing up in the receiving flask. But there was nitrogen pressure going in, so surely something had to be going somewhere, right? He looked up. . .and realized with a sinking heart that the vacuum manifold at the top of the hood was inexorably filling with the stuff. Now what? I always think in that kind of situation that it's time for lunch, no matter what the clock says.
Here's a question for the readership that should generate some interesting answers: what's the most valuable item you've seen someone ruin in a lab? I'll leave it broad enough to include both equipment and materials, and I expect to cringe numerous times on reading the comments.
I can put one into the hopper to start things off. Back some years ago, the guys down the hall from me had bought one of the largest Chiracel columns that were then sold. (For the non-chemists in the audience, this is a large packed column used to separate mirror image compound isomers (enantiomers) by pumping a mixture through). This was one of the ones where the chiral packing wasn't really bonded on to anything, but just sort of layered on another powdered solid support. And as the literature included with the column made clear, this meant that you could wash the stuff right off if you weren't careful with your solvent selection.
Well, it made it clear if you, like, read the sheet and everything. Which didn't stop someone from taking up their compound in methylene chloride and pumping it right onto the barely-used $15,000 (late 1980s money) column. And in the fullness of time (say, ten or fifteen minutes), out came the solvent front from the other end: cloudy, milky, swirling with opalescent shimmers like shampoo. Which shimmery stuff was, of course, the fifteen long ones of chiral resolving agent, scoured off the packing material by the cleansing wave of chlorinated solvent.
There: clean, simple, direct, and easily avoidable by spending two minutes reading a sheet of paper. That's the kind of thing I have in mind. Some additional examples?
I had a lot of long days when I was a teaching assistant. One semester I had three lab sections of sophomore organic chemistry to TA, and sometimes after the first lab of the week I knew that I was in for a rough ride.
Such was the case when we took the students through the good ol' Grignard reaction, as I mentioned here. The problem wasn't the reaction (I have to admit that I can't remember what we had everyone quench the reagent with, actually). The problem was the ether. We didn't have nearly enough fume hood space for all these reactions, so out on the bench tops they were. Two dozen steam baths hissing and gurgling, two dozen round-bottom flasks belching out ether vapors. By the third run-through of the week, I was teaching the lab section from out in the hall, trying to avoid another whiff of the damn stuff so I wouldn't have to run back in and puke in the sink.
The isoamyl acetate synthesis was another one in that vein. That's banana ester, the source of all the cheap, penetrating, one-note artificial banana flavor in the world. It's a pretty dramatic demonstration, in its way, because isoamyl alcohol and acetic acid (the starting materials) smell nothing like bananas, believe me. The workup was extraction with saturated sodium bicarbonate, a perfectly reasonable thing to do since there's an acid catalyst in the reaction.
But you do have to remember to vent the separatory funnel, oh yes. Close the stopcock, pour everything in, put in the stopper, shake it around - but never forget that shaking acid around with bicarb makes plenty of fizzy carbon dioxide. Especially don't forget about it until you've shaken the thing for, like, a minute and a half. And then open the stopcock. While it's pointed at your lab partner. This poor girl comes up to me with this huge stain all over the front of her dress, smelling like a shipwrecked banana boat, saying "I think I need to go back to the dorm and change".
The other thing that I can't help remembering is the day we had everyone do an oxidation with nitric acid, glucose to gluconic acid, I believe. Now, no one does that any more, at least since Emil Fischer keeled over from tasting his own compounds (a fact his Wikipedia article neglects to mention; I should edit that in there). And I'm not sure why that experiment was even in the lab curriculum. We did manage to get everyone into the fume hoods for it, though, thank God, except for one lunatic. He hauled his beaker of hot nitric goodness over to me, stuck it under my nose, and asked "Is this done yet?" I was coughing for three days. No matter where I end up over the next few months, at least no one's going to do that to me.
Dylan Stiles has a post up on distilling HMPA, which will be familiar to anyone who's worked with the solvent. The problem is, HMPA doesn't come dry, and it has be be dry to be any good. You can take the wimpy way out and dump a load of dessicant into a fresh bottle, but the only way to be sure is to distill the stuff.
Well, that's one of the problems with it. The other one, as Stiles mentions, is that it's carcinogenic. His advice - not to soak your genitalia in it - is sound. And that prompts me to update an old post from a few years ago here, to get it into the "How Not to Do It" archives. Longtime readers may recall it, but it's worth bringing back:
One fine afternoon in graduate school, I was peacefully advancing the cause of science when one of the guys from down the hall came into my lab. "What's HMPA smell like?" he asked. "Holy (excreta)!" I answered, "You think I know? Probably like it tastes, I guess." He told me that one of our recent postdocs was distilling it, and he was afraid that the smell in the lab was, well, HMPA, odd as that might sound.
I went down the hall to investigate, and came upon the single stupidest distillation rig I've ever encountered. There it was, a two-liter round-bottom flask with a heating mantle on it, boiling and bumping away on the high vacuum line. (OK, fair enough, if you're going to distill the stuff, you might as well get it over with.) On top of this lunking load of toxic solvent was the smallest still head in the group, a tiny little 14/20 short-path job that looked, in that context, to be about the size of the cherry on top of a triple banana split. This thing wasn't even slowing the hot HMPA vapors down much. My friend had a right to be suspicious, because yes, that smell probably was the springtime-fresh aroma of HMPA itself.
Unfortunately, I can't say much about the bouquet of this caricinogenic substance. You'd have to track down our Spanish post-doc and ask him; he was basically showering in the stuff. I stayed out in the hall while I ranted at him, and as he informed me that this was how they did it in Barcelona. "Well, go to Barcelona and do it!" I shouted, looking around to see if there were drops of solvent starting to run down the walls yet.
I left him peacefully distilling away, confident in his technique. Sometimes I wonder what's become of the guy. . .
(Note: no slur is intended on Spanish post-docs - I've worked with lunatics from all over the world, and as far as I can see, none of us are safe. Besides, if there were any country where people didn't make idiotic mistakes, they'd have taken over the world long ago, you'd think. . .)
Update: I should add that I haven't used HMPA myself in at least fifteen years. It's a no-win solvent in drug research. For one thing, it can be quite difficult to remove from your samples, and you have to make sure that it's all gone before you assay anything. If your reaction will only work with HMPA in it, it's as good as dead, because no scale-up group will use it. And if it'll work with something else besides HMPA, well, you should have just used that to start with and saved yourself the trouble.
A colleague of mine forwarded a copy of an accident report from Texas A&M. It seems that in mid-January they had a bit of a blowout there, thanks to a big liquid nitrogen tank. Now, liquid nitrogen cylinders are normally fairly benign, as long as you don't freeze your external organs off with the stuff or leave the liquid sitting around where it can condense oxygen out of the air. But idiocy will find a way - note the regular cylinder on the right and the new, improved model next to it.
These guys are usually equipped with pressure relief fittings, since nitrogen does tend to want to be a gas, and gases do tend to want to expand quite a bit. This tank, though, which seems to have been kicking around since 1980, had been retrofitted by a real buckaroo. Both the pressure relief and rupture disks had failed for some reason in the past, so they'd been removed and sealed off with metal plugs. You may commence shivering now.
Why it didn't blow long ago is a real stumper, but presumably people were taking nitrogen out of it quickly enough to keep things together. Not this time, though: at around 3 AM, things came to a head as the internal tank (these things are double-walled) expanded until it pressed against the outer one. That kept it from expanding anywhere else except on the ends, and as fate would have it, the bottom blew out first. The engineer's best guess is that this took place at around a 1200 psi load. It must have been quite a sight, although it's a damn good thing that no one was around to see it. I'll let the engineer's report take it from here:
The cylinder had been standing at one end of a ~20' x 40' laboratory on the second floor of the chemistry building. It was on a tile covered 4-6" thick concrete floor, directly over a reinforced concrete beam. The explosion blew all of the tile off of the floor for a 5' radius around the tank turning the tile into quarter sized pieces of shrapnel that embedded themselves in the walls and doors of the lab. The blast cracked the floor but due to the presence of the supporting beam, which shattered, the floor held. Since the floor held the force of the explosion was directed upward and propelled the cylinder, sans bottom, through the concrete ceiling of the lab into the mechanical room above. It struck two 3 inch water mains and drove them and the electrical wiring above them into the concrete roof of the building, cracking it. The cylinder came to rest on the third floor leaving a neat 20" diameter hole in its wake. The entrance door and wall of the lab were blown out into the hallway, all of the remaining walls of the lab were blown 4-8" off of their foundations. All of the windows, save one that was open, were blown out into the courtyard.
No one seems to have heard the celebrations, but someone noticed that the building's water pressure had gone a little wimpy and went to investigate, which I'll bet was a real eye-opener. I get the impression that they're still trying to track down the Mr. Fix-It who inadvertently rigged the tank for takeoff. The company engineer who came in to investigate noted that he's seen these kinds of "repair" jobs before, generally after they've powered through something.
We've had a few incidents recently at the Wonder Drug Factory where people have been using some common solvents like ether or tetrahydrofuran (THF) and ended up with odd results. When they check their reactions, there's something else in there that hasn't turned up before. The same impurity turns up in completely different reactions, too, which narrows the possibilities down a lot. In a couple of these cases, the chemists involved went to the trouble of isolating this pesky impurity and getting NMR spectra of it.
The experienced chemists in my audience are already cringing; I can feel it. No, we didn't blow anything up. But the people involved are now the proud owners of clean NMRs of ether peroxides. These explosive little beasts are an unavoidable byproduct of storing ethereal solvents where ambient oxygen can get to them. Oxygen is just too reactive - which is fine for us, biochemically, since that keeps us alive, but it can be a real nuisance in other situations.
These solvents are usually sold with some inhibitor added, a free-radical sponge like BHT, for example. But over time - or if someone in the supply chain stored things improperly - this will get used up, and then peroxidation moves right along. In extreme cases, such as with the unstoppable di-isopropyl ether, you can even get crystals of the peroxide coming out of solution. I have never seen this in person, and I will be very glad if I never do
Biologists and physicians have, among chemists anyway, a reputation for treating bottles of ether much more cavalierly than we do. A colleague of mine witnessed this at first hand at a former company of hers. A note had gone out to all the departments to check for old ether bottles, went into the possibility of crystal formation, and told everyone to notify the haz-mat team if any bottles were uncovered. In the molecular biology department where my source was working, one of the lab heads promptly marched out and rooted through the cabinets, emerging a few minutes later with a can of ether of uncertain age. This person then held the can up to his ear while shaking it, listening to see if any solid material was sloshing around in there. Which is one way to find out.
I wasn't an eyewitness to this one, although I wish I had been. I pass it on secondhard from a former colleague of mine, on which it made an understandable impression.
Bromine's an odd element. The two lighter halogens leading up to it are nasty gases. Fluorine, the first one, is actually beyond merely nasty, being a hazard to life from several different directions. Chlorine is something you can handle,although it was still nice enough to be used on the battlefield in World War I. But bromine is the first one of the series that makes the grade as a liquid at room temperature and pressure.
All the halogens have colors - for example, I'm told that liquid fluorine is green, not that I hope to see any of the damned stuff, and liquid chlorine is supposed to be yellow. Iodine is notoriously purple, and For its part bromine is a deep, almost opaque red-orange. It's one of those liquids that hasn't forgotten its gaseous heritage, and you always see it with a red haze of vapor above it. It's unmistakeable.
You'd think. Our story begins, as do so many fine lab disaster tales, with the phrase "There was this summer student. . ." In this case, there was this summer student whose grad-student supervisor thought he was ready for a spot of bromine work. They'd ordered a fresh bottle, which had come in from Aldrich the day before, and everything was ready for a good old-fashioned bromination reaction. As the chemists here know well, if you add bromine to a compound with an exposed carbon-carbon double bond, it'll react with the alkene, breaking it down to a single bond with a bromine on each carbon. Sometimes it's fast enough that you can see the red color disappear as the stuff drops into the reaction, and you can just go until the color persists, but sometimes it hangs around as an orange solution for a while.
So, our grad student leaves for lunch, entrusting this small-scale bromination to his ready-to-solo summer undergrad. And he wanders back presently to check out the reaction, but there's something wrong. The flask has no color to it, for one thing. For another, there are chunks of floating crud whirling around in the clear solution, and that's not right, either. He turns to the summer student and asks him if that's the bromination.
"Sure is!" Hmmm. Did he, in fact, use. . .bromine in this reaction? Oh yes, indeed. The bromine that just came in? Absolutely! What's that stuff floating around in the flask? Well. . .the bromine, right? Show me this bromine, then, by all means. And the summer student goes over to the opened shipment and lifts out the vermiculite packing material that's in there to keep the bottle from breaking during shipment. Behold the bromine!
No, you cannot make anything foolproof. Fools are too tricky. Needless to say, anyone who can't tell the difference between bromine and box-filler is someone that you don't want within a hundred yards of a working research lab. My colleague had not recorded the reaction of the grad student to this revealing answer, but I'm sure it involved raised voices and plenty of adjectives.
I wandered into the lab one Saturday morning while I was in graduate school - (OK, scratch that, I wandered into the lab most Saturday mornings while I was in graduate school, which was one of the things I hated about it.) And as I walked past the vacuum pump, I noticed something a little odd.
For those who don't work in a synthetic chemistry lab, the vacuum pump is where you put flasks of stuff after you've evaporated most of the solvent off of them. The pump pulls the last volatiles residues out of your syrup, crystals, or powder, leaving you with a dry weight that you can use to check your reaction yield, get pure spectra of the material, and so on.
The pump was making a different sound than usual. There was more of a rattle in it this morning, and less of a hum, if that makes any sense. I looked the thing over, trying to see what was going on, and finally I checked the row of stopcocks. Pay dirt! One of them was wide open, and the reason the pump was making that unusual sound was that it was trying to pump the air out of the entire chemistry building.
That isn't good for them. And in an academic lab, it's not like you could just reach into a cabinet and pull out another vacuum pump when you burned one out. They aren't cheap, and we spent time fixing the ones we had rather than attempt to ever buy new ones. So I twisted the glass stopcock closed, muttering foul gerunds, and left an unpleasant note taped to it. Something about how if you were the last user of this pump, you left the procreating stopcock open, and you shouldn't reproducing do that, etc. And I went about the rest of my merry morning's work.
The next morning, I wandered into the lab yet again. (I was there most Sunday mornings, too, damn it all.) And as I walked past the pump, I could swear that something was odd yet again. Surely not. I went back, looked it over, but couldn't see anything out of whack. Then, hardly believing it, I moved my taped note back to find the same stopcock, left wide open again. The slackjaw that did it it had to hold the note up to get to the stopcock, which really defied belief. There was the poor vacuum pump, trying to evacuate the air out of the state of North Carolina again.
I went stomping through our labs, looking for the culprit. But I was the only person there. No one in my group had done anything like that, so I wondered if someone else had been in there. . .then I remembered a Moroccan biochemist from another group who came over sometimes to hang out with the guy around the corner from me. Maybe. . .I went over to the next hall, and there was the pride of Marrakesh himself, humming tunelessly as he wandered about his lab.
So in I went, demanding to know if he'd been in our lab that morning, used our vacuum pump, and so on. He gave me a big grin: "No, I have not used this pump. But I have gone past it this morning, and I have thought, Hmmm, she is sucking ze air, no?" My testy reply was that if I found out that he'd been leaving the pump open then he would be sucking ze air, yes. For whatever reason, our apparatus went unmolested after that.
Summer students are showing up at academic and industrial labs around the country right about now. A certain percent of them will blow something up within the next three months, and that percent will be several standard deviations above the ka-boom rate of the other lab members. I'm not trying to say mean things about summer students. I merely speak the truth.
I had a summer undergraduate working with me for a while in grad school, and he taught me several lessons, of varying utility. One day he needed some dry benzene for a reaction, so I helped him set up a still in my hood. One-liter round-bottom flask, some benzene, a little sodium. My intern, who I'll refer to as Toxic John, put a heating mantle on the thing and turned it on.
A little while later, I walked past my hood and noticed that the stuff was boiling merrily. A bit too merrily, actually - it was really hopping around in there. I turned down the Variac (basically a big dimmer-switch type AC transformer that's used to step down the voltage to equipment like heaters) and went on my way. But I came back a little while later, and it was still rolling away in there.
If anything, it was worse. I turned down the Variac again, wondering just what was going on, and why my guess about the inital setting had been so wrong. A few minutes later, things hadn't improved much. The benzene was really leaping around, splattering and erupting. I looked a little more closely at the Variac this time, and noticed something that had escaped me: the heating mantle wasn't plugged into the damn thing at all.
Nope, it was plugged right into the wall socket, as some of my experienced readers will have guessed. As soon as I noticed that, I dropped the lab jack that was holding the heating mantle, which gave me a good look at the glowing red coils showing through the woven glass lining. I could feel it on my face like a sun lamp. Cursing, I pulled the thing out of my hood and heaved it into the hallway, right into a shopping cart that we kept out there for visits to the stockroom.
I went looking for Toxic John as the mantle popped and clicked. It was cooling down, but I wasn't. "John!", I shouted (I was pretty crabby back in grad school), "you plugged the mantle into the wall! No wonder it looked like a volcano in there!"
"What's the matter," he asked me. "Benzene doesn't burn, does it?" "Doesn't. . .burn. . ." I said slowly, as a nearby post-doc put a warning hand on my shoulder. "Well," said John, making his case, "it's inert to bromination!" That line of reasoning didn't impress me much, and as I recall, I told him that if he had any more insights like that we were going to find out if he was inert to bromination himself. Then I went off looking for the professor who'd just taught him sophomore organic chemistry, to let him know that his work, once again, had been in vain.
I thought about writing a whole post welcoming back my fellow Arkansan, the Ivory-Billed Woodpecker, but decided that not many folks would sit still for that. It does strike me, though, that the bird was rediscovered not that far from where I went to college.
I didn't have much time to hang out in the Cache River Wildlife Refuge, anyway, even if I'd been so inclined. I was living the life of the virtuous chemistry major, which for one semester led me to have all-afternoon lab sections every single day of the week. There's no doubt that I learned something in them, but not all of it was in the official curriculum.
Take one time in a physical chemistry lab, for example. We were doing something that no one in their right mind does in the real world, a determination of molecular weight by boiling point elevation. That's a relic of the sealing-wax era of chemistry, but it's not quite as much of a waste of time as those qualitative organic tests I was railing about the other week. I still wouldn't keep it in a lab course these days, but the way this one was set up, it did have some value.
We were figuring out the molecular weight of benzoic acid by adding increasing amounts of it to a solution (toluene, I seem to recall) and seeing how much the boiling point increased. We then plotted this out, running it through Raoult's Law to get the answer.
Well, of course, we already knew the molecular weight of benzoic acid. But as we took boiling point after boiling point, in a finicky apparatus that splashed boiling solvent over the lowered bulb of our thermometer, we could tell that something was going wrong. We were getting something way over 200, and benzoic acid weighs 122. Hrm. We checked everything again, but the data all looked pretty good. Way over 200. . .about 244, actually.
Then it hit one of us. "Dimer!" he said suddenly. "Benzoic acid forms a dimer, and that weighs twice as much!" We all slapped our foreheads and grinned. But the guys next door to us had put their minds at rest even before we did.
Not that they had the right answer. In fact, their plot showed the molecular weight of the dissolved benzoic acid at 122, right on the nose. Pretty good curve, too, pointed right at the seemingly-correct-but-impossible answer. How did they get there? By taking so many data points that the freaks and throwaways could be assembled, Frankenstein-style, into a zombie plot that gave the answer that they just knew it had to give. Never mind that the other 90% of the data pointed to twice that number - that can't be right. What do the data points know about right answers, anyway?
I was using hydrogen chloride gas straight out of the cylinder today, first time I've done that in many years. That's a very different substance from regular hydrochloric acid, which is technically a solution of HCl gas in water. The straight stuff will really snap your head back if you get a whiff of it, which you'd better not since it does Bad Stuff to your lungs, as you'd imagine.
You need to rig up a trap for the vented gas, since it's rather bad form just to send it up the fume hood. The standard way is to bubble the excess through aqueous base to neutralize it, preferably rigged up so that the water doesn't have a clear path to go siphoning back into your reaction if the pressure goes haywire. Bubbling the HCl into a solvent like methanol always looks a little odd. You can send a pretty vigorous stream of the gas in one side and have very little coming out through the trap at all, since the methanol is soaking up so much of it.
These gas cylinders are under pressure, and the large ones look just like the helium tanks that non-scientists are familiar with from balloon vendors. The regulator valves on top of them need to be made of rather more robust material for an HCl tank than for helium - which is, after all, totally inert under all conditions short of the interior of the sun. Back in grad school, a corroded regulator (on a whopping big HCl cylinder) gave me a real scare as it threatened to give way and vent all the gas at the full tank-neck pressure of about 1500 psi. I had to go sit down for a while after that one.
But today's work was with a lecture bottle, a much smaller cylinder that holds only a couple of hundred grams of the gas. That's enough to do some damage, true, but not on the scale of the free-standing ones. I saw that happen back in graduate school as well. One day I was sitting in the library, looking up some references, when I noticed the occupants of the third floor research labs pouring out onto the lawn from the rarely-used side stairwells. They were hustling right along, too, which suggested some sort of liveliness upstairs.
As it turned out, it was in the lab next door to mine. One of the guys had another HCl tank, a medium-sized one, which was also corroded and jammed. He went for the cylinder wrench, which he then used for the non-standard purpose of vigorously whanging the valve with strong overhand strokes. One of the other guys in the lab summed up the sound of this process as "PING. . .PING. . .PING. . .hisssssSSSSSSS oh @!?#!"
The hood wasn't enhanced by having a kilo or two of hydrogen chloride vented all over it, that's for sure. It looked as if it had been subjected to some sort of accelerated aging process - if there were a market for antiquing lab equipment, this would be a good way to do it. All the exposed metal was pitted and flecked with green. The stainless steel was hazed with rust, having reached its carrying capacity for corrosion. And everything still had a fine mist of concentrated hydrochloric acid all over it where the gas had sucked the water out of the air and condensed on the nearest surface. Cleaning it up is not the way you want to spend your Friday afternoon.
None of that for me today, though. I ran the stuff in uneventfully, with the reaction turning to a clear yellow, which is nothing compared to the colors I'd turn if you sprayed that much on me. I'll find out tomorrow if things have worked according to plan. One thing's certain: something will have happened. Nothing escapes from HCl gas unchanged.
Someone came to my lab today to borrow some thiophenol, a request that made me think of something that happened in my first summer of undergraduate research - twenty-two, gulp, years ago. Now, thiophenol is not known as a great inducer of nostalgia. Like the other small-molecule sulfur compounds, it reeks without letup. It's a major part of the smell of burning rubber, so if you can imagine that concentrated and put into a bottle, you've got a pretty good idea. It's distinctive.
I was using this cologne as a starting material, reacting it with sulfuryl chloride, which is another reagent that no one is going to dab behind their ears. It's a reactive chlorinating agent and a fairly strong oxidizer, and it'll make you shake your head and snort if you come across its fumes, for sure. The two of them together make an eyebrow-raising mixture - I was exiled to a lab at the other end of the building while I ran this one, just because of the potential smells.
Heating this brew gives you phenylsulfenyl chloride, a red oil which combines the foul properties of its parent compounds. You distill the stuff out of the reaction, cap it up, and store it in the cold. I think it's too reactive to be an article of commerce; you have to make it fresh. And make it fresh I did, even though everything around me smelled as it had been dead for weeks. In the freezer with the stuff for the weekend (we didn't work grad-school hours at my all-undergraduate school, not even in these summer research programs.)
Monday morning I went down and picked up the flask. Hmmm. . .no longer a red oil. Odd. The stuff had changed to a pale yellow solid, which didn't seem right. I wasn't sure of the compound's freezing point, but the color change alone made me wonder. I stood there puzzled for a minute or so with the tightly stoppered flask in my hand, holding it up to see what I could make of the stuff. And then, with a loud gunshot bang, the top of the flask exploded in my hand.
I jumped straight up in the air, flinging away the lower part of the flask that I was still holding. I came down on the balls of my feet, in a fight-or-flight stance, looking around wildly. I didn't feel as if I'd been injured, but I'd never had anything blow up while I was holding it, either, so who kenw? After a second or so I looked down to see if I was OK. And weirdly enough, I was. I can't imagine how I managed not to pick up some glass shrapnel, at least - perhaps even at that early point in my career I had enough sense not to point the neck of a round-bottom flask toward me. My hand was fine - I kept flexing my fingers in wonderment. After a minute of two of stalking around the room, shaking and gibbering, I started looking around to see what had become of the chemical.
I found it about ten feet away, a lens-shaped piece of light yellow stuff, molded smooth by the inside of the flask. Whatever it was, it wasn't melting again, and it sure wasn't phenylsulfenyl chloride. We figured out what it was pretty quickly, but I think I'll leave its identity as an exercise for the technical portion of my readership - guesses to go in the comments below. If you get it right, you'll know why it blew up, too!
We have a lot of pyrophoric substances in an organic chemistry lab - things that burst into flame when they encounter normal air. Liquids are handled by syringe and needle, using bottles fitted with gas-tight septae. That works pretty well, once you get the hang of it. The main things you have to learn are to provide some inert gas as a replacement if you're removing a large volume, and to not twitch your arm and hand muscles while you're holding a syringe full of stuff. (That can provide a spectacular flamethrower effect, which is fine if that's what you're after, but we rarely are.)
Pyrophoric solids are a bit trickier. Some of them (like sodium hydride) are often sold as a fine powder mixed with mineral oil, which coats everything and keeps it from igniting. Of course, you have to get rid of it at some point, because it's sure not going to go anywhere. You can either wash off the mineral oil before you begin, while the solid is inside your reaction flask, or clean it away from your compound at the end of the reaction. I usually opt for the latter, on practical grounds: by the time, I'll know if the reaction worked, and I won't have wasted the initial effort on a loser. Besides, most of those reactions need purification anyway.
You can buy dry sodium hydride, but I'm not a fan of the stuff. It goes bad too quickly, and an NaH fire is a beast to put out if it really gets going. A carbon dioxide extinguisher usually isn't up to the job, and it'll blow the powder all over your lab, which isn't recommended. And you most surely don't want to throw water on the stuff, although you'd have to be the village idiot to try that one. You'd certainly get a village idiot's reward for your efforts. The only reliable way to put one of these things out is to bury it in sand or some other inert powder. Then you have to let it cool down for a while - if you don't, it'll just whoomph up on you again when you try to clean the place up. If that doesn't make you feel like you're wasting your day, I don't know what will.
The related potassium hydride is invariably sold in a hard-to-handle suspension that's mostly oil. I've never seen it packaged dry, and I don't want to. It ignites much more easily than the sodium compound, to the point that people even manage to start fires with the oil-soaked item. At least the flames are prettier.
And the metals themselves are usually sold and stored under oil. Sodium metal, as my fellow chemists know, is interesting stuff. It's soft enough to cut with a metal spatula (the texture is rather like cold butter), and is very shiny indeed until the air hits it. You can work with it out like that, if you move with reasonable speed and get it under some inert solvent. Potassium metal is much less forgiving, and I have no desire to work with the heavier metals in the series (rubidium, cesium) as their elemental metals, because they just get worse as they go up.
You can mess up the area with just plain sodium, though, oh yeah. Some collagues of mine had a summer undergraduate (here's where experienced chemists start to grin and pull their chairs closer), and they were teaching him how to handle sodium metal: take it out of the oil, have your beaker of hexane ready, cut it like so, pick it up (the point of the spatula works well), drop it in the solvent, and so on. Everything went fine. So the next day, one of the guys tells him to go down and weigh out, say, five grams of sodium for a reaction. The summer student scampers off, and a few minutes later, the grad student wanders down to have a look, just to make sure things are going OK.
And there the guy is, with his beaker of hexane, sawing away with a spatula at a big cylinder of sodium metal which he is gripping in his bare left hand. Well, he hadn't been told not to do that, true, but neither had he been warned not to fetch it in his teeth. You just sort of take these things for granted. The summer student had no doubt taken the skin of his left hand for granted, too. (It took a few weeks, but he recovered.)
I'm not sure if I've told this story before - perhaps on my previous blog. The memory is with me so vividly that the retellings run together.
I've never liked hydrogenation rooms. For my non-lab audience, that's where we keep the equipment for running reactions in pressure vessels under hydrogen gas, always with some sort of metal catalyst to make the hydrogen come in and reduce things. It's about as close to witchcraft as modern organic chemistry gets.
And it's just those ingredients that make me nervous. Big metal cylinders of hydrogen gas can't help but bring to mind visions of the Hindenberg, for one thing. If something fails on the apparatus, it generally fails with spraying, fizzing, and/or flames. And any hydrog room that's been around a few years invariably picks up black residues of spilled catalyst everywhere. It's in the cracks of the lab bench and in the fittings of the equipment.
You want to be careful with that stuff. Most of the time, we use powdered charcoal impregnated with palladium or platinum, which looks like, well, charcoal. But under the right (um, wrong) conditions, it can come to life like you wouldn't believe. In the presence of hydrogen gas and some air, such as when you mess up and open the flask, the powder gets so hot it glows bright orange. It looks like it's just come out of a furnace, and that's about when it ignites your reaction solvent. Then you might as well get out the hot dogs and suntan lotion, because the fireworks are going off already.
Once about thirteen years ago, I was in my company's old hydrogenation room getting a balloon full of the gas to take back to my lab. You can run less vigorous reactions under just that much gas, so there's generally some fitting for people to fill things up to go. In this case, you opened up the valve near the gas cylinder just a crack, and then opened up the needle valve near the balloon a lot.
Or, anyway, that's how we'd been doing it. When I got there that afternoon, someone had just decided that they liked it better the other way around, where you just barely dinked the balloon takeoff valve. News to me. I stuck my balloon assembly on, opened the valves up the way I was used to, and KA-BLAM!
Off goes my whole balloon and tubing rig, flying off the fitting from the blast of gas. And up I went, straight into the air, with my hair, I swear, standing straight out from my head in some sort of mad-scientist perm. My adrenal glands hit the Emergency Squirt button and dumped their entire load of adrenaline into my bloodstream, convinced that at long last a sabretooth had shown up at the mouth of the cave.
I landed on my toes and bounced back up like a rubber ball. A videotape of my actions would be worth watching; I've often been sorry that I don't have one. I was pinging around the room kicking at the cabinets, waving my arms and gibbering obscenities. It took a couple of circuits of the place before I could slow down enough to be sure that I had all my extremities attached. I hate hydrogenation rooms. . .
"Twenty-five years of being a laboratory chemist, says Gregory Hlatky today, and this is the first time I've had an incident." Hey, maybe he's been doing the wrong kind of chemistry. Some kinds can almost guarantee you an incident every month or two!
Actually, the kind of chemistry he does (organometallics) is already pretty lively, and I have to say that I'm impressed by his 25-year safety record in a field like that. But I'm not surprised that it was an organoaluminum compound that took off on him, because I've had several of them do the same thing to me (without injury, fortunately.)
And the most nerve-wracking part of them was the time delay. Most reactive compounds are very forthright about their reactivity. They burst into flame on exposure to air (like tertiary-butyllithium, or for the hard-core pyromaniacs, the dialkyl zincs.) Or they give off great clouds of choking fumes (I can recommend neat titanium tetrachloride for those who want to experience this special effect - the one molar solution in dichloromethane is for dilettantes,) or hiss and splatter violently if they encounter water (chlorosulfonic acid is a winner in that category.) At any rate, you know very quickly, if you didn't already, what kind of substance you're dealing with.
But the alkylaluminum compounds have their coy ways. I recall a large aluminum alkyne reaction that I set up in graduate school, one of the once-and-for-all reactions which get scaled up a little more than is prudent. I don't believe that I've told this story on the blog, so this one can go into the file with my other lab stories.
With this reaction, the fun started early. I first had to add a large amount of n-butyllithium, which is less reliably pyrotechnic than the tertiary kind, and I'd done that by running the solution into a Pyrex dropping funnel. That, for the non-chemists in the readership, is one of these. As the stuff dripped slowly into the reaction, the BuLi had dried into a crust on the glass tip of the funnel. At the end of the addition, I had to pour in a larger quantity of toluene and switch the funnel to another piece of glassware, so I just grabbed the thing and swapped it out. Whereupon it burst into lovely orange/purple flames.
Well, they went out shortly. But I was standing there, pouring my (flammable) toluene with one hand, and holding this flaming funnel with the other, thinking that this would be a good time for a member of the department safety committee to show up. After that, the next steps of the reaction went along relatively quietly, and eventually it was time to quench the reaction. I did that by very carefully adding a few drops of methanol to the liter or so of solution. Nothing. So I added a few more.
Nothing. I waited to see if anything would happen. Nothing did. A brief squirt of methanol this time. Zilch. I was starting to wonder if there was going to be any reaction at all - surely there was some leftover organoaluminum stuff that needed to be quenched. A longer, more vigorous squirt of methanol. No sign of life.
Then, at the very bottom of the flask, down near the magnetic stir bar, a bubble formed and rose to the surface. And another. Several more. A stream, several streams, a vigorous fizzing mass that came roaring and foaming up the sides of the flask - Well, I managed to catch most of it in a bucket. When the bubbling started to really roll along, I had bolted for something to catch things in, because it was clear that the reaction could be on its way to a spectacular conclusion. It didn't disappoint me.
Like most once-and-for-all reactions, I had to do that one again, eventually. But the next time, at least, I was ready for it.
Several of us started in on a "stupid lab tricks" conversation at lunch the other day. When chemists get together, we always know that that topic's available if we run out of things to talk about. Anyone with reasonable organic lab experience has a stock of favorites.
One of mine is a fellow grad student who was trying to save money by recycling acetone, the wash acetone that he used to clean out his flasks. He had a four-liter round bottom with a side-stopper on it, rigged up to a big distillation head, and the thing was always cooking away.
Of course, the acetone in the pot got nastier and nastier as time went on, as he kept refilling it with whatever gorp he washed out of his dirty glassware. During the time I knew it, it was a deep, opaque chocolate brown with sort of purple overtones - not the usual color you look for in your acetone supply. The stuff he distilled off was pretty decent, but even so. . .
Well, eventually this guy took a vacation (for the first time since I'd joined the group.) He took off for a few days, and turned off the still before he left. There it sat, just as ugly in repose, until he came back and flipped on the power to the heating mantle. All was quiet, for a while.
I was right around the corner when I heard it: a loud "PING-whUUUrrrsh - splattt!" This alarming noise was followed by a really fearsome stench, a knock-you-back brew of who knew how many stinky carbonyl byproducts (no doubt including, as fellow organic chemists reading this have figured, a generous helping of mesityl oxide.) I charged through this to find that the still had blown the side-arm stopper clear across the room - the pinging sound was its ricochet off the wall - and the foul brown concoction had come geysering out after it. The far wall was a Rorschach blot of dripping slime, the source of the eye-crossing aroma.
The problem - as those who've made similar mistakes well know - was when the heating was turned off. All the grunge in the pot had, for the first time in months, finally had a chance to settle down to the bottom of the flask. Where it coated the pile of boiling chips down there with resinous mung, rendering them useless. Which allowed the whole shebang to superheat once the power was turned back on, until something finally oozed aside long enough to allow the first bubble to form. And then, as Louis said, aprez-moi, le deluge.
A solid majority of lab-accident stories start out "We had this solvent still. . ." That's why you won't find any of them in any large industrial environment. It's just not worth the opportunity to add to the story file!
After swapping stories, which some of did again at work recently, you wonder how anyone physically survives their academic chemistry training. Chemists usually come out of their degree programs with a stockpile of good yarns, filed under headings like "Idiotic Lab Explosions" and "Maniacs I Have Worked Next To, And Their Life-Threatening Ideas." Your own explosion stories usually start "One time when I was up in the lab at three AM. . ."
I'll pass some of these along every so often, to give folks outside the field an idea. Before doing that, though, I should mention that the litany of explosions drops off dramatically when you get into industry - and no, I don't miss them. The responsible factors are experience, better facilities, not working all hours of the night, and a certain weeding-out of the real hard-core crazies.
I recall one party I went to back in grad school. Several of us from Chemistry were standing around telling ball-of-flame stories, to the great interest of some law students. One of the guys down the hall from me, though, piped up and said "I don't know what y'all are going on about - I've never had an explosion in my life."
Well, the Chemistry Gods listen to you when you say things like that, and they reach for their bottles of laxative. The next morning, my friend was cleaning out a solvent distillation pot. . .and here's where my organic chemistry readership all start to grin. Cleaning out solvent stills is the all-time leading method of starting lab fires in synthetic chemistry, because you tend to distill many solvents from mixtures involving metals like sodium or (God help you) potassium. Bits and chunks of these lively substances tend to hide under layers of sludge as you try to inactivate them, only to jump out and do their thing long after you're sure everything's been quenched.
Which is just what happened that morning. My friend was sure everything was fine, and rinsed the (theoretically) tamed mixture down the sink (which they won't let you do any more, for the most part, but this was back in the mid-80s.) A couple of seconds went by quietly, then there was a muffled "thoongh!" from deep in the pipes - followed at speed by a three-foot geyser of flaming gunk straight up from the sink drain. I heard the shouting, and came down to find him standing wide-eyed in a thin haze of smoke, still holding the flask. "I never should've said that, should I?" were his words. .