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DBL%20Hendrix%20small.png College chemistry, 1983

Derek Lowe The 2002 Model

Dbl%20new%20portrait%20B%26W.png After 10 years of blogging. . .

Derek Lowe, an Arkansan by birth, got his BA from Hendrix College and his PhD in organic chemistry from Duke before spending time in Germany on a Humboldt Fellowship on his post-doc. He's worked for several major pharmaceutical companies since 1989 on drug discovery projects against schizophrenia, Alzheimer's, diabetes, osteoporosis and other diseases. To contact Derek email him directly: derekb.lowe@gmail.com Twitter: Dereklowe

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February 1, 2012

Potassium Hydride Is Not Your Friend

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Posted by Derek

Noted chem-blogger Milkshake seems to have had a close call with a fire started by a tiny potassium hydride residue. It looks like he made it through without serious injury, but that sort of thing will definitely shake a person up.

I hate potassium hydride. Its relative sodium hydride is a common reagent, but it's much tamer (and even so, can cause interesting fires - I knew someone who ignited a heap of it on the pan of a balance while he was weighing it out, which slowed things down a bit). Sodium hydride is usually sold as a 60% dispersion, a dark grey powder soaked with mineral oil to keep it from deteriorating too quickly (and to keep it from setting everything on fire). You can buy 95% sodium hydride, the dry stuff, and there are people who swear by it, but I tend to sweat at it. You never know if it's been stored properly; you may be adding a slug of sodium hydroxide to your reaction without knowing it. And there's the fire part. You'll want to move briskly if you're using the 95%, and I'd pick a day when the humidity is low.

But potassium hydride, that's another beast entirely. It makes the sodium compound look like corn meal, in terms of how forgiving it is. You can't get away with the clumpy oily powder form at all - traditionally, KH is sold as a gooey dispersion of grey powder sitting under a few inches of mineral oil. If it's well dispersed, it's supposed to be 35%. You shake the stuff up until you think it's even mixed, then pipet out the amount of gunk that corresponded to the KH contained therein. Sure you do. What actually happens is that you pipet out the stuff, noticing while you do that it's already settling out inside the pipet, thereby to clog it up when you try to transfer it. No fun.

It's becoming available now dispersed in a block of wax, which is not such a bad idea at all. Wax isn't any harder to get out of your reaction than oil is, and you can carve off chunks and weigh them without so many what-am-I-doing moments. But Milkshake worries that this ease of use will lead to more fires during workups (which is where his reaction ran into trouble), and he may well be right. If you're going to use KH, don't let your guard down.

Comments (47) + TrackBacks (0) | Category: Chemical News | Life in the Drug Labs | Safety Warnings


COMMENTS

1. Ben Zene on February 1, 2012 9:06 AM writes...

Someone has to say it. I hope Milk wasn't too shaken up by the experience!

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2. exGlaxoid on February 1, 2012 9:13 AM writes...

This sounds familiar. I had one series of reactions years ago that required KH in a moderately large scale. (We tried NaH to no avail.) Of course, this was summer in the South, so the humidity was ~100%.

Fortunately, I was warned, and the reactions were run in an argon filled glove bag, and everything was carefully done, then quenched carefully. But we also learned that the reactions were best done in an empty hood, far from anything flammable, as it was not uncommon for a pipette or spatula than we thought had been quenched to catch fire within seconds of being removed from the bag.

Also, if you wash the KH (or even NaH in the humid south) with hexanes, be aware that the washing are EXTREMELY easy to have spontaneously ignite due to humidity. They should be quenched very carefully, AWAY FROM OTHER CHEMICALS, and not handled open, keep a watchglass or cap on the container when handling.

I had several times when objects contaminated with KH would spark or ignite, but since there was nothing else near them, and I had an empty glass baking pan to sit them in, no harm was done, and they would burn out in seconds. And it is always a good idea to have a bucket of sand nearby when working with any flammable organometallic reagent.

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3. opsomath on February 1, 2012 9:22 AM writes...

Worse still, when we bought NaH from Alfa Aesar, it came in a plastic bag with a Twistie, inside a tin can. What in the world. I store my coffee in a better sealed container than that, and I have definitely seen this method from other manufacturers as well.

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4. ddd on February 1, 2012 9:25 AM writes...

Glovebox/Schlenk line?

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5. cirby on February 1, 2012 9:38 AM writes...

Every time I read this blog, I think up CSI plots.

"Well, we either have three very clever serial arsonists working together, or..."

"Or what?"

"One clumsy industrial chemist."

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6. IchDich on February 1, 2012 9:53 AM writes...

@3 Opsomath: The polymer bags with NaH are truly terrible. I had to quench one of them last week. Bag torn open, all globules that are white on the outside but nicely grey and reactive on the inside. A big problem ther is that once the container is opened and more than two weeks old, people start questioning the quality.
Great way to spend your afternoon, pouring ~50g of old NaH dispersion in isopropanol...

Often use commercial NaH 95% though, and ALWAYS make sure to transfer in a closed vessel. I do find it more convenient and effective than the 60% dispersion.

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7. Just Sayin' on February 1, 2012 10:13 AM writes...

KH + 18-C-6 + ethereal solvent = PARTYYYY!!!!

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8. RB Woodweird on February 1, 2012 10:19 AM writes...

KH in wax is genius. I only have two questions: What is grade 4.8 nitrogen? and What happens to the wax?

Now if they can get tBuLi in wax, I'll be impressed.

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9. annon on February 1, 2012 10:23 AM writes...

Big fire in Steve Burke's lab at South Carolina during the 80's (Bill Murtishaw, a terrific chemist) involving KH and quenching of a small amount resulting in a major fire with extensive damage. Even with caution difficult to deal with.

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10. cynical1 on February 1, 2012 10:26 AM writes...

Years ago, they used to sell sodium dispersion in mineral oil too. It was lab clean up day at our small company which meant we disposed of our nasty chemicals ourselves. One of my co-workers blew himself up that day. He was pretty lucky that Johns Hopkins has a great 'Burn Unit' but he looked like Freddy Kruger afterwards for awhile. Nothing more fun than hearing the explosion and then looking up and seeing a fireball traveling across the ceiling tiles. Funny thing was that the sprinkler system didn't go off. That was a very good thing. It was a big bottle. I got my on-the-job fire extinguisher training that day. Fun, fun.

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11. Industrialist on February 1, 2012 10:29 AM writes...

Grade 4.8 nitrogen is 99.998% pure - 4 nines and an eight.

Matt

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12. Andrew on February 1, 2012 11:01 AM writes...

@6 IchDich

Was bagging and tagging the unwanted NaH for collection an option?

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13. startup on February 1, 2012 11:05 AM writes...

Don't you have gloveboxes? Or people who know how to use a Schlenk line?

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14. Anonymous on February 1, 2012 11:10 AM writes...

@12 Andrew

Our safety officers can be really unrealistic in the procedures for the removal of dangerous chemicals. It is the theoretical option, but quenching it yourself is less painful. It's a job for the experienced post-docs, and done in a safe manner.

This does lead me to another point. Let's say you bag it and ship it. What happens next? Its related to something I saw over at Chemjobber. They just burn it?

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15. BV on February 1, 2012 11:12 AM writes...

@9: Bill Murtiashaw was my first supervisor in pharma; he delighted in telling stories of chucking huge chunks of Na into lakes.

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16. Vader on February 1, 2012 11:17 AM writes...

My high school chemistry teacher told us about a student who swiped a big chunk of sodium, wrapped in in a hankie, stuffed it in his pocket, and made a quick getaway out of the lab.

Tripped into a puddle of water on the way to wherever he was headed with the sodium.

Didn't kill him, but I believe he was still a candidate for a Darwin Award.

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17. CG on February 1, 2012 11:23 AM writes...

@9 and 15: Bill was a real gentleman. Was an honor to have known him. RIP Bill.

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18. anonymous on February 1, 2012 11:26 AM writes...

Reminded me of that video. http://video.google.com/videoplay?docid=-2134266654801392897#

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19. molecular architect on February 1, 2012 11:31 AM writes...

@6: Many years ago in grad school, we developed a safer method of quenching pyrophoric chemicals when faced with the task of cleaning out a storeroom with many old bottles of organo-lithium reagents, NaH, KH, Na/K alloy, etc.

In a large deep bucket or pan in an otherwise empty hood, place a large beaker containing the isopropanol for quenching. Surround the beaker with additional iPrOH and add dry ice to the surrounding alcohol. Slowly pour, shake or spoon the pyrophoric reagent into the beaker. The cloud of CO2 will instantly smother flames. If done properly, you can actually pour n-BuLi rather quickly without exceeding the ability of the CO2 to quench the flames.

Have used this method many times in my career without problems.

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20. Andrew on February 1, 2012 11:32 AM writes...

@14 A few years ago the lab I was in at the time had an NaH fire. There was about 100 g left, I arranged with the waste company to pick it up. I reminded them thrice of the hazardous nature, and they told me not to worry: the Schlenk flask of NaH, the silicone oil we drowned it in (don't ask!), and the 3 L beaker we put it in went straight from the truck into the furnace. I called them to check after, and they said there were no problems.

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21. InfMP on February 1, 2012 11:45 AM writes...

I am a big fan of NaH 95% because you don't have the grease at the end that reduces chromatographic resolution. Also, if it's a clean reaction you can easily carry on to the next step without having grease in your flask.
When I used to live next to the atlantic, the summers were so humid. I ended up knowing I had exactly 2 min to get it all into the flask before it would ignite and i had a metal pan i would throw it into if it started to ignite. At the same time, adding too fast can lead to a bubbling overflow. But isn't that what makes chemistry fun?

Speaking of waxed reagents, J. J. La Clair used this sort of technique in his famous Hexacyclinol synthesis.

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22. Chemjobber on February 1, 2012 1:15 PM writes...

@14, @20: Thanks for the link!

I'm under the impression that the better places have some sort of separated drum-crushing apparatus that chews up a drum and its contents before it goes into the furnace, but I really don't know.

Nevertheless, from that one example you linked to, there are places where a person actually opens the drum, removes the materials, etc.

Permalink to Comment

23. newnickname on February 1, 2012 1:55 PM writes...

First: I would often (a) tare flask + septum (b) add NaH or KH oil dispersion (c) wash / rinse by syringe with pentane [CAUTION: small amounts of MH get sucked into the syringe, so squirt into iPrOH, not directly to waste] until the pentane washings leave no oily residue on a piece of ground glass (e.g., a stopper) (d) Argon stream or quick vacuum to remove pentane (e) reweigh OIL-FREE MH (f) Use oil-free, O2-free, under-Ar MH for reaction.

Second: I try to avoid alcohol quenches until I'm down to a safer stage. I quench BuLi and LAH with EtOAC diluted in hexane, NOT alcohols. You avoid generating butane and hydrogen gases. You can also
use it to consume Na metal in your Na stills by an acyloin but that takes a while. EtOAc should also work to tone down NaH and KH quenches via the ester enolate, but I don't recall ever having done that so I don't want to assume the kinetics will always be in your favor. (I would usually just use alcohols or aqueous workups with my NaH / KH.)

Another milder STOICHIOMETRIC quenching agent is a solution of a cheap, dry phenol such as BHT in an inert solvent. I also use BHT to titrate my metallic bases: BuLi, LDA, etc.. A LOT easier to get accurate amounts that using tBuOH or some of the other old recommendations.

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24. Retron on February 1, 2012 2:11 PM writes...

i've personally lit potassium hydride on fire after ~50% of the quenches, used it about 8 times. i used a similar approach to @19 but worse. a big deep recrystallizing dish with a shallow layer of i-PrOH. clear the hood out of flammables. have a bigger recrystallizing dish that will fit over the top of the quenching one so the purple flame can be extinguished. add KH REALLY slow. enjoy.

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25. Andrew on February 1, 2012 2:42 PM writes...

I like to use dry ice for Grignards or BuLi.

To destroy sodium metal, I know Prudent Practices recommends adding a large volume of high boiling non-reactive solvent like toluene, and then adding ethanol. I've done it before, but it works very slowly. I also don't like the idea of adding more combustibles to a fire hazard.

The last time I quenched sodium I transferred the sodium balls into a Schlenk, and stored them temporarily under vaccum. Working one ball at a time, I cut them into tiny bits, and dropped them into water. Just like you might a sodium-water demonstration in a classroom. Water won't burn, and if your pieces are small enough, the hydrogen won't catch fire either. But it's really slow.

Is there a better way?

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26. bbooooooya on February 1, 2012 3:17 PM writes...

"The polymer bags with NaH are truly terrible. "

Good memories though. I recall a British PDF explaining to me that when the NaH was white it was no good, and proceeded to pour a bunch of it in the sink to dispose of it. Turns out, it still had some activity. Thankfully, no one was hurt.

Permalink to Comment

27. Dave on February 1, 2012 3:52 PM writes...

There are times, when dealing with pyrophoric materials, where the appropriate response is to run screaming from the lab. After all, labs can be rebuilt. But, trained chemists are rare.

Actually, some fire departments use that philosophy, too. Sometimes, it's better to just pull back and let the problem resolve itself, rather than rushing men and equipment into a disastrous situation.

Dave

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28. CMCguy on February 1, 2012 4:00 PM writes...

#14/#19 I served as Chemical Hygiene Officer for number of years and found out often many things that would make perfect sense to Chemists in regards to pre-quenching or deactivation of hazardous waste are actually illegal unless have proper permits, facility, safety equipment, PPE, Training and other bureaucratic compliance. Although very costly, nor logical and appeared even more potentially dangerous it was better to pay for contracted experts to handle in the end.

Of course this was many years after the days when would enjoying take the shaving of Na or K (indeed must nastier as indicted here) down to a local fountain or creek for "fireworks" displays.

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29. PMP on February 1, 2012 4:28 PM writes...

The good ol' Vogel says that potassium hydride residues _will_ catch fire when they contact water. I, too, found out that this prophecy is 100% accurate - at the sink, just like Milkshake, but fortunately the fire was much smaller.

Some chemicals simply leave no margin for error.

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30. CR on February 1, 2012 4:34 PM writes...

@15, BV:
"Bill Murtiashaw was my first supervisor in pharma; he delighted in telling stories of chucking huge chunks of Na into lakes."

Is there anyone out there that hasn't heard some form of this story? It's either throwing chunks of Na into lakes, or those that worked near the ocean doing it there.

@16, Vader:

"My high school chemistry teacher told us about a student who swiped a big chunk of sodium, wrapped in in a hankie, stuffed it in his pocket, and made a quick getaway out of the lab.

Tripped into a puddle of water on the way to wherever he was headed with the sodium.

Didn't kill him, but I believe he was still a candidate for a Darwin Award."

Sorry, this never happened. Seriously. Think about it. Seems a story your HS teacher made up for gullible kids.

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31. metaphysician on February 1, 2012 5:28 PM writes...

#30-

If I might ask, further context? My own thought is that it sounds unlikely. A decent sized chunk of sodium isn't going to react *that* fast to water. Even if it did happen, it would be more painful and embarassing than dangerous. That said, I might be misremembering the reactivity.

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32. Greg Hlatky on February 1, 2012 6:45 PM writes...

"After all, labs can be rebuilt. But, trained chemists are rare."

You obviously have no future in management.

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33. anonymous on February 1, 2012 7:03 PM writes...

@31 - Hmmm, I've done it (Na into a pond) and it is impressive!!
Signed, A Stupid (Ex)Chemist

Permalink to Comment

34. milkshake on February 1, 2012 7:27 PM writes...

Derek, thank you for the reference. (The Wordpress servers are deluged by the incoming traffic.)

@31,33: A big chunk of sodium catching on fire in a pond: One reasons why K metal is a much worse fire hazard than Na in the lab settings is that even a tiny shaving of K metal can self-ignite on contact with moisture or oxygen whereas a tiny Na chunk will just as likely fizzle out or crust up. My understanding is that very low melting point of potassium metal is the main reason for this difference: once the alkali metal chunk had heated itself enough the reaction of the molten metal suddenly takes off at runaway rate and the metal surface promptly becomes hot enough to ignite the H2/air mixture. For this reason the K+Na alloys that are liquid at room temperature are super-pyrophoric, far more unforgiving than pure K metal.

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35. chirality on February 1, 2012 7:28 PM writes...

#31 Sodium will dance on the surface but the generated hydrogen will eventually blow up.
I once had to clean a lab space which had previously been occupied by an idiot who, unknown to me, used to dump leftover sodium into a large plastic container originally used to store potassium hydroxide. Of course, he did not bother to remove the original commercial KOH label or write any kind of warning on the container. It was a perfect setup as I then proceeded to dispose the gunk 'of the KOH that obviously absorbed a lot of water' by first diluting it with more water. To my horror, large chunks of sodium metal appeared on the surface. I did not feel like being a hero that day, so I just left the lab to watch the inevitable fireball from a safe distance.

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36. Secondaire on February 1, 2012 8:19 PM writes...

Oh, so many memories.

Used to use NaH for some Dieckmann condensations in grad school. We used to use the 95%, and yes, it does go bad in the polymer baggie (as does LiOMe - who ships alkoxides this way?!!). Opened once to find hole burned right through the bag, with hydride spilling out, and we quenched it in a merry brazier of flames in a big crystallizing dish.

Also, when I was an undergrad, my labmate and I got a 4-L beaker and took turns throwing in chunks of sodium. The best part was when we got a compression wave which blasted the water out and propelled the sodium into a 2-inch pancake in the ceiling light fixture. I think it's probably still there, only now oxide/hydroxide...

These days, I handle sodium under hexane, and quench in alcohol at 0 ˚C - then leaving the residue somewhere to hydrolyze the (comparatively) innocuous alkoxide

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37. Curious Wavefunction on February 2, 2012 10:43 AM writes...

Reminds me of Oliver Sacks dropping every alkali metal from sodium through cesium into a river from a bridge, as recounted in his book "Uncle Tungsten". Once you see that visual display you won't ever have to memorize alkali metal reactivity trends from a textbook. The good old days.

Permalink to Comment

38. Hasufin on February 2, 2012 10:59 AM writes...

To this day, I regret not walking off with the kilogram of undocumented potassium I found in my high school chem lab storeroom. I cannot readily imagine the kind of explosion that would make, though I'm sure it would have been quite impressive.

On the other hand... well, you know, hands. I have both of them. So maybe it's for the best that I put it in the hazardous waste bin like I was supposed to.

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39. Tt on February 2, 2012 11:09 AM writes...

Once worked with a Russian who swore by his native country's sodium quench protocol... Quarter fill a large beaker with some water and top it off with a large top layer of ether (or MTBE), add sodium metal ( smallish chunks) and watch the bouncing ball. Not nearly as dangerous as it sounds.

Permalink to Comment

40. okemist on February 2, 2012 12:19 PM writes...

In the pysics dept at nyu in 1969, my step dad was doing PhD, one of the students left 500g of Na on sink shelf, it fell into the sink and the explosion killed the father of someone I later went to high school with. IDK what the pysicists were doing with sodium, especially in that quantity, but it wound up being a horrible tragity.

Permalink to Comment

41. newnickname on February 2, 2012 3:07 PM writes...

Sodium in the pond (or Charles River) tricks: MIT undergrads used to (still do?) toss lump Na or K into the Charles River basin. There were (still are?) several youtube videos. A few years ago, a lump of encrusted sodium was picked up by a river cleanup crew and it caused a fire and seriously injured two people on the boat. There was an out of court legal settlement.

I'm attaching URLs / links, if they get through:

http://www.boston.com/news/globe/city_region/breaking_news/2007/09/mit_prank_may_h.html

http://www.upi.com/Top_News/US/2010/08/11/MIT-frat-settles-suit-over-explosion/UPI-50561281503987/

Permalink to Comment

42. molecular architect on February 3, 2012 12:12 PM writes...

@41: Just read the two links you posted. Contradictory information as to whether the Na was in a container or not. It's hard to believe that, if not in a container, that the metal would not have all reacted within a short period of time. The metal is usually stored in a light mineral oil.

Whatever the specifics, it's unfortunate that two innocent people were seriously injured by this prank; one which I'm sure many readers of this blog have done sometime in their chemistry career.

What made me shake my head in disbelief though is the statement that the clean-up organization incurred thousands of dollars in decontamination costs for their boat and equipment. Damn, sodium hydroxide must be difficult to wash away!!

Permalink to Comment

43. non-pharma chemist on February 4, 2012 10:18 PM writes...

@ #10 You mean they don't sell sodium dispersion in mineral oil anymore? I used it once a couple of years ago and it was truly terrifying. Nothing like flames shooting out from underneath the cap as you screw it back onto the bottle.

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44. JH on February 5, 2012 8:39 AM writes...

@#19 molecular architect

This sounds like a good idea for many cases. However, I understand that you cannot put out a LAH fire with CO2, because it is such a strong reductant. I guess it depends on details, whether it is H2 that is burning or LAH.

Permalink to Comment

45. haxwithaxe on February 13, 2012 1:39 AM writes...

@42
to wash away just add water :P
you'll know it's gone when your equipment stops fizzing XD

Permalink to Comment

46. Richard Ash on May 22, 2013 3:19 PM writes...

This reminds me of a tale told of industrial research being done on sodium-sulphur rechargeable batteries. These have to be kept hot (I think around 90 C) to keep the chemicals molten and so working. The problem is if they work too well they get hotter, and so thermal run-away occurs.

The reaction to a run-away seemed to be to get the offending battery out of the lab, and throw it as far as possible across the sports pitch so it could catch fire on it's own.

This went wrong for one researcher who aimed a bit far to the right - into the fire-fighting water reservoir (a large, deep pond). The resulting bang and fireball was apparently spectacular!

Permalink to Comment

47. Anonymous on May 23, 2013 7:02 PM writes...

#25 - unfortunately, of course, the faster the reaction, the more energetic it is.

There's several YouTubes of the "we do it so you don't have to" variety, that react sodium metal with a variety of solvents, like sodium (violent, natch), glacial acetic acid (even more violent, especially once the sodium melts) and ethanol (pretty tame, but slow as you are aware).

Acetone was the most violent that I saw with any reasonably-sized chunk of sodium (which is odd because it's less acidic than acetic acid), and I wouldn't try anything more violent than that (though of course others have, like reacting it with hydrochloric, sulfuric and even nitric acids; at least nobody's been foolhardy enough to try perchloric).

Maybe a flask or bath of one-third water, two-thirds hexane or heptane, which separates the sodium's heat from the ambient oxygen, so the hydrogen doesn't have an ignition source? Kind of like #39's bouncing ball.

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