About this Author
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: Twitter: Dereklowe

Chemistry and Drug Data: Drugbank
Chempedia Lab
Synthetic Pages
Organic Chemistry Portal
Not Voodoo

Chemistry and Pharma Blogs:
Org Prep Daily
The Haystack
A New Merck, Reviewed
Liberal Arts Chemistry
Electron Pusher
All Things Metathesis
C&E News Blogs
Chemiotics II
Chemical Space
Noel O'Blog
In Vivo Blog
Terra Sigilatta
BBSRC/Douglas Kell
Realizations in Biostatistics
ChemSpider Blog
Organic Chem - Education & Industry
Pharma Strategy Blog
No Name No Slogan
Practical Fragments
The Curious Wavefunction
Natural Product Man
Fragment Literature
Chemistry World Blog
Synthetic Nature
Chemistry Blog
Synthesizing Ideas
Eye on FDA
Chemical Forums
Symyx Blog
Sceptical Chymist
Lamentations on Chemistry
Computational Organic Chemistry
Mining Drugs
Henry Rzepa

Science Blogs and News:
Bad Science
The Loom
Uncertain Principles
Fierce Biotech
Blogs for Industry
Omics! Omics!
Young Female Scientist
Notional Slurry
Nobel Intent
SciTech Daily
Science Blog
Gene Expression (I)
Gene Expression (II)
Adventures in Ethics and Science
Transterrestrial Musings
Slashdot Science
Cosmic Variance
Biology News Net

Medical Blogs
DB's Medical Rants
Science-Based Medicine
Respectful Insolence
Diabetes Mine

Economics and Business
Marginal Revolution
The Volokh Conspiracy
Knowledge Problem

Politics / Current Events
Virginia Postrel
Belmont Club
Mickey Kaus

Belles Lettres
Uncouth Reflections
Arts and Letters Daily
In the Pipeline: Don't miss Derek Lowe's excellent commentary on drug discovery and the pharma industry in general at In the Pipeline

In the Pipeline

« Vial Thirty-Three Rides Again | Main | Peter Kim, So Far »

June 7, 2006

Best When Used By. . .

Email This Entry

Posted by Derek

Most of the things we use in an organic chemistry lab can sit around for reasonable periods of time. I've used reagents from bottles that are older than I am (OK, this was twenty years ago, so it's getting a bit harder to do). As long as the stuff isn't air- or moisture-sensitive, it can hang around a long time. That lets out the violently reactive things - don't expect to find the same piece of potassium metal you left if you're silly enough to leave it out while you answer the phone, for example. (In fact, you'd better pick up a fire extinguisher on the way back, just on general principles).

But there are some reagents that don't react with air, but rather react with themselves, which can make them particularly hard to handle. If the reaction is exothermic, things can get dangerous. The heat given off by the first bit that reacts tends to set off some others, which really gets a good amount going, and ba-doom. Even if you're not in the ba-doom category, there are some things that you need to look out for. Styrene, for example, is always sold with some free-radical inhibitors in it, because if it gets a chance for a radical chain reaction to start, the whole bottle will seize up into a warm gunky block of polystyrene. (That's for small bottles - larger ones won't be able to transfer their heat so well to the surface of the container and can make a much bigger mess).

Benzaldehyde isn't so violent, but it slowly forms a six-membered-ring trimer on standing. Update: I've been carrying this idea around for twenty-five years now, but it's wrong. The nonaromatic aldehydes love to trimerize, but benzaldehyde and the other aromatic ones don't. The solid gunk is benzoic acid, from air oxidation, which is a separate category of How Reagents Go Bad. Old bottles can have some crusty crystals of the stuff around the neck of the bottle. The reagent is one of those things that you really have to distill before using it if you want to trust your results. Update: This point is definitely still true!

The extreme case in the self-condensation category is probably cyclopentadiene. It does a Diels-Alder reaction with itself first chance it gets, so it's always sold as the dimer. If you want the pure monomer, you have to distill for it. The dimer cracks thermally, so the vapor condensing at the top of the still is a different substance than the stuff down in the pot. Collect it, keep it on ice, and use it - the diene's a ware that will not keep.

Comments (18) + TrackBacks (0) | Category: Life in the Drug Labs


1. eugene on June 7, 2006 10:19 PM writes...

I did the cyclopentadiene cracking in an undergraduate lab where we then used it to make Ferrocene. Afterwards, we used the ferrocene in a Friedel-Crafts acylation. That was one of my favorite labs in undergrad and I still remember it today. And it is the only time I've ever used Ferrocene or cyclopentadiene. The sublimation with dry ice to purify it was also something I did for the first time -- but not the last -- in that lab.

What made it even better was that we did it all from instructions in a huge binder; the TA was not allowed to help us or talk to us too much and was only there to make sure we didn't hurt ourselves too badly.

That lab played a major role in me deciding to go into Organometallics.

I think.

Maybe I just like to say that though, because I'm not certain of the exact factors which led to the decision. But regardless, that was one neat reaction.

Cyclopentadiene is now getting less use as a ligand since the new polymerization catalysts are found to have better activity when they are relatively electron poor. So six electron donor aromatic ligands are sort of seen as old timey these days. And a twelve electron dicyclopentadiene ligand motif makes for a very classic Ziegler-Natta catalyst system. This is of course, my general feeling, and will probably make some people upset.

Permalink to Comment

2. secret milkshake on June 7, 2006 11:42 PM writes...

I am pretty certain that benzaldehyde actualy does not form cyclic trimer. Benzaldehyde forms simply benzoic acid because for some reasons benzaldehyde is much more sensitive to radical-chain air oxidation than many other aromatic aldehydes.

Problem with (cyclic) oligomers is typical for very reactive aldehydes like ethyl glyoxylate. (Always crack-distill fresh to depolymerise the Fluka toluene solution)

Permalink to Comment

3. SP on June 7, 2006 11:48 PM writes...

1- Did you go to MIT, or is that experimantal sequence just a common undergrad experiment everywhere? I'd guess the latter, because it's the easiest column in the world to run- separating ferrocene from acetylferrocene lets you actually watch the bands separate as you run the thing.
For the biologists out there, acrylamide is fun- I don't think it will ever begin polymerization without the appropriate catalyst, but once it does it's quite exothermic. Don't try to dispose of a bottle by just tossing in some APS to make a bottle of gel because the bottle won't remain intact.

Permalink to Comment

4. AJ on June 8, 2006 3:38 AM writes...

It's always good to know some principles an in my chemistry lab i also work with some very old bottles. They always end up in the warehouse for some time.
I just regret in reading this post that you didn't mention diethyl ether, that is something most people have in the lab, sometimes it stays open a bit to the air, and most people don't know is that it forms peroxides very easily, and that can make a nasty fire.

Permalink to Comment

5. Michael G on June 8, 2006 4:47 AM writes...

SP - I was thinking the same when I read the first comment and I went to university in Scotland (Strathclyde Uni, home of Peter Pauson who discovered Ferrocene but got the wrong structure - d'oh! Nature (Lonson) 1951, 168, 1039). They love it there and we were constantly reminded of how he'd missed out on the Nobel Prize. It is a nice easy column, the first one I ever did (and never used alumina since...)

Permalink to Comment

6. Justin B on June 8, 2006 5:27 AM writes...

Formaldehyde trimerizes into 1,3,5-trioxane, and acetaldehyde trimerizes into 2,4,6-trimethyl-1,3,5-trioxane (paracetaldehyde), as far as I know those are the two most common aldehyde trimers.

It's a little off topic, but sometimes the trimers (or hexamers) can be much more useful and easier to use. An example is hexamethylenetetramine (the hexamer of ammonia and formaldehyde). In acid, like TFA, it hydrolyzes into formaldehyde, so it is much easier to use (in my opinion) than an aqueous formaldehyde solution. HMTA is also sold as "esbit" in backpacking and outdoorsy-type stores as a solid fuel tablet.

Permalink to Comment

7. Lou on June 8, 2006 5:39 AM writes...

SP - I know a 1st year Ph.D. student in Biochemistry who did indeed try to dispose of an old bottle of acrylamide-bisacrylamide by chucking in some APS. At least she transferred the liquid into a beaker before doing it, and she tells me that there was a big blot of gel bouncing around after the addition.

Recently, a postdoc in our lab tried to dispose of an old plastic bottle of acrylamide-bisacrylamide solution by adding APS - she added it directly to the bottle. The plastic bottle softened considerably due to the exothermic reaction (but didn't melt totally, thank god).

Funny how biologists can't keep their hands off that stuff.

Permalink to Comment

8. NJBiologist on June 8, 2006 6:27 AM writes...

#7--I keep my hands off the stuff by using precast gels.... They're more consistent than anything I can pour, and the lab has fewer bottles of stuff that makes me a little nervous (acrylamide???).

Permalink to Comment

9. Derek Lowe on June 8, 2006 7:59 AM writes...

I left acrylamide off the list because not so many chemistry labs have it around (gels are alien to us, fortunately).

Secret M, I just sat down with SciFinder (which I neglected to do when I wrote this post last night), and you're right: the aromatic aldehydes don't like to form the trimers like the aliphatic ones do. I've been carrying a benzaldehyde myth around in my head for 25 years, now expunged. I'm changing the post to reflect this.

Permalink to Comment

10. Philip on June 8, 2006 8:04 AM writes...

To answer SP, the ferrocene experiment goes on everywhere. I learned it at UNC-CH as an undergrad (in 74)and taught it to undergrads at Cornell a few years later.

I agree with milkshake, the crusty stuff is benzoic acid. The trimer is more of an alkyl aldehyde problem. We had problems with hexadecinal forming the trimer.

Interestingly, the formaldehyde trimer is a controlled substance.

Permalink to Comment

11. Milo on June 8, 2006 10:22 AM writes...

I recently found a bottle of benzaldehyde that was solid (!). An NMR showed almost > 95% acid.

I have found that piperidines often come as a liquid, but have the nasty habit of forming a nice white solid that tends to gum up the cap of the bottle (N-oxide?).

Permalink to Comment

12. eugene on June 8, 2006 10:42 AM writes...


For that reaction, I went to undergrad at McGill (Montreal), so it's a general reaction that is used everywhere. I forgot that we had to do the column as well which we had to prepare by our favorite method (dry loading is still the best method for me). Overall, that reaction teaches you a lot of useful techniques.

Permalink to Comment

13. Thomas McEntee on June 8, 2006 11:07 AM writes...

Re Philip's comment, paraldehyde (the cyclic trimer of acetaldehyde) is a DEA controlled substance (see 21 USC Section 812) but 1,3,5-trioxane, the cyclic trimer of formaldehyde (metaldehyde) is not. However, simple mixing of formaldehyde or its trimer with ammonium hydroxide followed by evaporation yields hexamethylenetetramine (hexamine, urotropin) which can be nitrated under very specific conditions (the details of which, dear reader, I shall not divulge) to produce the energetic materials known as RDX and HMX.

Permalink to Comment

14. SP on June 8, 2006 12:23 PM writes...

That reminds me of a chemistry prof's explanation of why he got into chemistry- he's sort of a hippie type, and he said, "In the 60s there were two reasons to get into chemistry- drugs and explosives. I don't like blowing things up."

Permalink to Comment

15. secret milkshake on June 8, 2006 12:26 PM writes...

Old leaky amine bottles could develop solids from CO2 absorbtion.
It is only strongly-basic primary and secondary amines that form solid CO2 adducts. I think the product is a carbamate salt R2NCO2- (R2NH2+)

Permalink to Comment

16. Derek Lowe on June 8, 2006 2:39 PM writes...

That's my understanding, too. And that would explain why you often see that stuff on bottles of piperazines and pipidines, but I don't ever recall seeing it around the cap of a bottle of morpholine.

There will be a future post in a few weeks on watching things slowly oxidize, provisionally titled "Quinoline Isn't Really Black".

Permalink to Comment

17. Anonymous on June 8, 2006 7:46 PM writes...

There is a legend about 2 Czech chemists transporting vinyl chloride. (They could not order it from Aldrich, under communism). Apparently they went into a PVC-producing plant and had the monomer condensed for them into a steel pressure vessel. They took the pressure tank and were driving back to university a in a small car. After some time the guy who was holding the pressure vessel said: "listen, it is getting awfully hot, I wonder if they added any stabilisers. Oh shit, it is polymerizing!" They managed to threw the tank out of the moving car and it exploded as soon as it hit pavement. There was a policeman nearby and the guys spent a night in jail as terror suspects. (It did not help it happened just before May 1st, a big commie holiday)

Permalink to Comment

18. Rigel on June 11, 2006 4:34 PM writes...

Runaway polymerization of Cp species is a problem. The thought of a cylinder of some neat Cp monomer in the cargo hold of the plane I'm riding to Basel brings little comfort. Despite the regs there is still lots of room for error in judgement on the part of suppliers.

Being in the Cp ligand business, I can say that we encourage people to consider buying the lithium salts of specialty Cp ligands. Cracking dimers or oligomers of expensive and scarce Cp species can lead to even greater scarcity...

Permalink to Comment


Remember Me?


Email this entry to:

Your email address:

Message (optional):

The Last Post
The GSK Layoffs Continue, By Proxy
The Move is Nigh
Another Alzheimer's IPO
Cutbacks at C&E News
Sanofi Pays to Get Back Into Oncology
An Irresponsible Statement About Curing Cancer
Oliver Sacks on Turning Back to Chemistry