August 8, 2010
I have been working in the same chemistry laboratory for more than three years and by now I feel like quite the seasoned veteran. However, what happened yesterday served as a good reminder to me to always treat chemistry with respect. Perhaps I have become a bit too cocky. Here is the story.
About a week ago, I decided that I needed to synthesize a sulfoxide from the corresponding thioether. The thioether starting material I was using is a commercially available compound, but the sulfoxide I needed to make has, as far as my literature searches could discern, never been made before. Now although I had not performed a novel organic synthetic reaction in quite some time (I have been focusing much of my time studying electrochemistry as of late), this one-step synthesis did not seem particularly challenging or interesting to me. A quick look in an organic textbook taught me that sulfoxides can be easily made in good yield through the use of a variety of different oxidizing agents. The main problem of many of the older reaction conditions utilized is that they produce significant amounts of the over oxidized sulfone byproduct.
Although there were many different reaction conditions and reagents form which to choose, I elected to use peracetic acid as the oxidizing agent and acetone as a solvent. I chose this route because this is what was described in a paper outlining the synthesis of sulfoxides from thioethers that have very similar moieties on them as my compound does. In fact, let me quote the experimental procedure from the paper published in 1997 in Tetrahedron that I followed:
The thioether (2.0 mmol) “was dissolved in acetone (10mL), cooled to -15C, and 34% peracetic acid (10mL) was added. The reaction mixture was kept for 24h at this temperature, 15mL of chloroform were added, and the mixture was stirred for 0.5h at 18 C, the chloroform layer was separated, washed with 10% NaOH (3 X 15mL) and water (2 X 20mL), after which it was dried over anhydrous MgSO4, filtered, and the solvent was removed under reduced pressure.”
This procedure definitely sounded easy enough to me. From the passage above, it seems that the yields were so good and the product was so clean, that even purification by column chromatography was deemed unnecessary. So I went ahead and performed this procedure to the letter. I had a bit of trouble devising a way to maintain a constant temperature of -15C overnight, but I managed to come pretty close by immersing my reaction vessel in a large salt-ice bath. I charged the round bottom flask with ~200 mg of the starting material and 10 mL of acetone before cooling it to the requisite temperature. I then added 10 mL of 34% peracetic acid dropwise to the reaction mixture while stirring. Nothing spectacular seemed to occur.
The next morning I came to the lab and proceeded to take a TLC (thin layer chromatography) of the reaction mixture to monitor its progress. I performed a microscale workup of the solution in a test tube in order to basify and dry it before subjecting it to TLC. I determined that a good solvent system for eluting the starting material was 95% dichloromethane-5% isopropyl alcohol. The starting material was, as expected, very visible under UV light due to a benzene ring in its structure. Once I tried to spot my reaction mixture on my TLC plate, I immediately ran into problems. No matter how many drops of solution I spotted on the TLC plate, the compound was not UV-active. I tried staining it with iodine, but this also was to no avail.
So I let the reaction mixture stir for another 6 hours and for another hour at room temperature after adding chloroform (giving it a total of more than 24 hours) and after failing to get a UV-active spot again, I decided to work up the entire reaction mixture. As per the instructions, I added chloroform, washed with base and water, dried with anhydrous magnesium sulfate, and rotavapped the solution down. After doing so, I got about 150 mg of a white solid. With this solid in hand, I was able to make a very concentrated solution of the product which I was able to make a nice UV-active TLC spot with. The TLC showed a faint spot corresponding to the starting material and another spot that was much more nonpolar. This was surprising since sulfoxides are generally more polar than their parent thioethers.
At this point, I was a bit confused by the difficulties I had in analyzing the product via TLC and the unexpected polarity of my product. However, it looked to me that by TLC, my starting material had been consumed and that I had generated a new product. In order to see if I had isolated the sulfoxide, I took an IR of my product, but the characteristic sulfoxide S=O stretching frequency (1030-1060 cm-1) was not present. So it looked like I didn’t have what I wanted, but I had no idea what I did create since I obviously had a pure white solid. I thought maybe I had just gotten the starting material. Then I took an NMR of my compound to see what it was and that was when I was really confused. I used deuteurated acetone as my solvent for NMR since the starting material was not very soluble in chloroform. Much to my surprise, I only saw three peaks- three singlets. Two of the singlets could be easily ascribed to water and acetone. This meant that my compound only had one NMR peak!
Instead of interpreting the data I had obtained as real, I dismissed it, thinking that I had performed recorded the NMR spectrum incorrectly, or that the compound was so wet with residual solvents and water that the three peaks I saw outweighed the actual signal. After talking it over with a friend, we both agreed that I should try the reaction over again at room temperature. Perhaps that reaction simply was not pushed hard enough.
So I went ahead and performed the same reaction this time on a 100 mg scale at room temperature, but still with 10 mL of acetone and 10 mL of 34% peracetic acid. I left the reaction stirring overnight and when I came back the next morning, a good deal of white precipitate had formed. Aha! This was a sign that something different was happening! I went ahead and monitored the reaction using TLC and had the same frustrating experience. I worked up the whole reaction, rotvapped the organic layer down, and got a white solid with the same appearance as before. It looked like I had a lot of it. When I weighed my product, I was surprised to find that I had greater than 400 mg, a greater than 400% yield!
This was troubling to me as I knew the solid I had was relatively pure by its appearance and by NMR. I sat down at my desk to think and for the first time, tried to think about how all of the chemicals I had added to the reaction mixture might interact. What could be causing the 400% yield? Since the compound was not very UV-active, perhaps it was masked by some ionic compound. The only thing I could think of was that somehow a whole lot of magnesium sulfate (the drying agent) had gotten into the organic layer. This did not seem at all possible given that magnesium sulfate is not at all soluble in acetone or chloroform.
I thought about it harder. Then suddenly, I remembered a story written by a fellow amateur chemistry in which he unknowingly made large quantities of acetone peroxide. I immediately recognized that this was what I had done. I took some about 5 mg of the white solid and put it under a Busnsen burner and it immediately ignited into an orange fireball. I took a slightly larger amount, lit it, and an even larger fireball resulted. For comparison, I took some of the starting material and tried to ignite it. The compound just melted, smoked, and formed a waxy residue.
All of the experimental puzzle pieces now fit together in my head. Looking at the structure of acetone peroxide, the dimer would indeed only have one chemically distinct type of hydrogen as the NMR spectrum I took revealed. The compound would not be very UV-active since it is aliphatic, which explains my troubles visualizing it during TLC analysis. And I would also expect it to be less polar than my starting material which is what I did see once I finally spotted enough on the TLC plate. Acetone peroxide is synthesized from the acid-catalyzed reaction of hydrogen peroxide and acetone. Peracetic acid probably decomposed into hydrogen peroxide and acetic acid. The acetic acid formed would be more than enough to catalyze the reaction.
Although I was initially relieved that I figured out the mystery, the more I thought about it, the more lucky I realized I was. I had formed the more unstable dimer which forms preferably at elevated temperatures. When pyromaniacs synthesize acetone peroxide, they always make sure they make the trimer by keeping the temperature low through the course of the reaction. It is said that acetone peroxide is more unstable when it is dry and when it is impure. I had made extremely dry acetone peroxide that was laced with the starting material. I was lucky that I did not blow off my fingers while taking the product off of the rotavap. I thought I was being safe by performing the reaction on a small 100 mg scale. However, the two things that reacted, peracetic acid and acetone were present in large quantities. I could have easily made grams of the stuff even though in my head I was performing the reaction on a small scale.
One thing I realized that it is very important as a chemist to learn how to synthesize explosive materials. Although, unlike many amateur chemists, I completely shied away from pyrotechnics, I still read a lot about other chemists’ exploits and now realize that this provided me with invaluable knowledge that can help keep me safe. The most frustrating thing about the whole experience was that the procedure I used to attempt the synthesis of the sulfoxide must have been inherently flawed. There was no mention of a warning that it would be possible to make acetone peroxide as a byproduct. I learned that you have to always question the validity of literature and that you should never blindly follow an experimental procedure as I did.