Hot Buttered

Orville Redenbacher is on the TV, telling us again how great his microwave popcorn is, and by the way, it doesn't contain diacetyl. Any more.

Diacetyl (emphasis on the first syllable) is also called biacetyl (emphasis on the last syllable) and the latter is what we called in when I was working on the photooxidation of aromatic hydrocarbons a couple or three decades ago. Biactetyl, in fact, occupies an important place in the history of smog chemistry, though I have to admit the notion of "important" is open to interpretation.

There are basically four kinds of "reactive organics" that are important in smog photochemistry: paraffins, olefins, aromatics, and carbonyl compounds (aldehydes and ketones), the latter being more commonly formed in the smog process than emitted outright. I'm taking a bit of a liberty here by omitting alcohols, ethers, and other oxygenated compounds, partly because, ethanol and MTBE notwithstanding, they still don't amount to a large fraction of the mix, and partly because their photochemistry is pretty close to that of paraffins, or ketones that don't photolyze, i.e. break up by the direct action of sunlight.

The early days of smog chemistry were dominated by research into the chemistry of paraffins and olefins, so much so, in fact, that it wasn't until the mid-1970s that researchers realized that the photolysis of aldehydes and ketones was the primary source of catalytic radicals in the smog formation process. In fact, that was the biggest single difference between the first photochemical kinetic mechanism that I worked with, the Hecht-Seinfeld mechanism, and the later, Hecht, Seinfeld, Dodge mechanism. The former used oxygen atoms (from the photolysis of NO2) as its primary radical source, whereas the latter used formaldehyde and higher aldehydes to that purpose.

Both of these mechanisms were based on smog chamber experiments involving butane and propylene (or propene, if you're a nomenclature purist). Aromatics chemistry was tacked on as an afterthought, not because it was believed to be unimportant, but more because nobody had any idea what to do with it.

Aromatic hydrocarbons, as they are called, all have a "benzene ring" somewhere in them, and that makes everything very complex. Perhaps you remember the story about Friedrich Kekule literally dreaming up benzene's structure. It's formula is C6H6, and its structure "bites its own tail," so each carbon atom, with four chemical bonds, has, after accounting for the hydrogen, three bonds to share with its two neighboring carbon atoms. That could work out to two and one or one and two, i.e. a paraffinic bond with one neighbor and an olefinic bond with the other, but the wonders of quantum mechanisms allows it to actually be one and a half bonds with each neighbor. Such are the wonders of quantum electrons being able to be in several places at the same time.

Benzene itself is almost dead, photochemically speaking; put it into a smog chamber and it mostly just sits there, making a little tang of phenol after a while, but phenol is deader still, so…boring.

But if you replace one or more of benzene's hydrogens with a methyl group (-CH3), now you're talking. One added methyl group gives you toluene. Two, and you get xylene, which comes in three isomers, meta, para, and ortho, depending upon whether the methyl groups sit right next to each other (ortho), on opposite sides of the ring (para) or one over (meta). There are also, of course, trimethylated benzenes, and compounds where the substituted groups are more complex than methyl groups. But actually, toluene and the xylenes make up the bulk of aromatic compounds in air pollution. There is even a refinery stream referred to as "TBX" which stands for toluene, benzene, and xylene.

Okay, so I'm going to tell you how the photochemistry works, then how it got figured out. The tricky part had to do with how the aromatic rings would open up. Everyone knew it had to happen sometime, but how, and what the products were was a mystery for years.

What happens to something like toluene in smog is that, when it encounters an hydroxyl radical (-OH), the hydroxyl adds itself onto the ring somewhere, usually at the carbon that sits next to a methyl group, because of the way that methyl groups mess with the electron distribution of the aromatic ring. This is what hydroxyls do with olefins, incidentally, so you can look on it as the hydroxyl briefly looking at the ring and seeing, not that "one and a half bonds" thing I mentioned above, but a double carbon-carbon bond, which hydroxyls just love to glom onto.

This breaks one of the carbon-carbon bonds, and one end of it now has a romantic relationship with the hydroxyl radical. But the other end, like a jilted lover, is on the rebound, ready to pick up with just about any pretty face that comes by. That face, almost always, belongs to oxygen, a really promiscuous molecule. It's diatomic (i.e. O2), but not so committed to the relationship that it passes up some good carbon bond action.

So an O2 gloms onto the other, lonely, carbon and you now have a peroxy radical, an aromatic ring with an oxygen tail. The radical characteristic of the thing tends to be concentrated at the free swinging tip of the tail, and in most peroxy radicals, that tip winds up reacting with some other molecule.

Not so with the aromatic peroxy radicals, however, because it so happens that the radical tip is just right for swinging around and hooking up with another carbon, somewhere else on the aromatic ring. You may now consider all of the other sexual double entendres that I could use for this situation.

Anyway, another oxygen now gloms onto the group, but now the situation is stable enough (maybe) so that it waits around for some outside compound (usually a molecule of nitric oxide—NO) to take the last lonely oxygen atom away from the daisy chain.

All the oxygens then decide to settle down with their new carbon best buddies. The oxygen-oxygen bonds call it quits, and that leaves another oxygen bond for each oxygen connected carbon. If you're counting, and remember that carbon only has four bonds to its name, this means that it has a double bond with an oxygen, one for either a hydrogen or a methyl group, and, whoops, only one left for another carbon in the aromatic ring. In short, the ring opens, in multiple places, once for each oxygen. At some point, the poor hydroxyl group, which is now the radical of the bunch, meets yet another oxygen molecule and the hydrogen leaves the party to for hydroperyoxyl (HO2).

The aromatic ring is pretty much finished at this point, and it cleaves into at least two pieces, one with two ring carbons, the other with four. The one with four has, in addition to two oxygen atoms, a olefinic bond (there was some belief for a while that the fragments might all have two ring carbons, each, meaning that there would have been another oxygen molecule bridge on the ring, but later product yield measurements indicate otherwise).

Both ring fragments are called "dicarbonyls" because they each have two carbonyl (C=O) bonds. In one of the fragments, the two carbonyl bonds are right next to each other.

The simplest dicarbonyl is called "glyoxal." It's just H(C=O)(C=O)H. The next one is methyl glyoxal, with a single added methyl group: H(C=O)(C=O)CH3. Both of these are very hard to measure; they tend to stick to gas chromatographic columns nigh onto forever.

Ah, but the next in line is a dicarbonyl with two methyl substituants: CH3(C=O)(C=O)CH3. This is called biacetyl, or diacetyl. And it comes through a chromatographic column.

If you photooxidize orthoxylene, with it's two adjacent methyl groups, when the ring opens, a certain percentage of time you get biacetyl. A group at the University of California at Riverside, (Darnall, Atkinson, and Pitts, 1979) saw the biacetyl coming off of their chromatograph and realized that they had seen the first evidence of ring opening products.

It so happens that both biacetyl and methylglyoxal photolyze like crazy, so much so that they last only a few minutes in sunlight before splitting into radical fragments. I had been looking for something exactly like these dicarbonyls in my own studies of aromatics photochemistry, because I'd found good evidence of very powerful radical sources in toluene experiments. My calculations indicated that the radical formation rate from toluene was twice what it would be if toluene were going to pure formaldehyde, which of course it does not. It forms a significant amount of methyl glyoxal, and that was what I was looking for.

Later, I heard that biacetyl/diacetyl was used to flavor margarine; I also heard that microwave food products use excess flavoring agents because the microwave heating process drives the volatiles away faster than regular cooking.

I had some vague suspicions that it might not be a good idea to use a compound as photochemically unstable as biacetyl in food. Light causes biacetyl to break into two pieces, both acetyl radicals, and when there is any oxygen around, you get peroxyacetyl radicals. Add some nitrogen dioxide and you get peroxyacetyl nitrate (PAN), which is biologically active. Actually, it's a good bet that any give peroxy compound is biologically active. These are some pretty potent radicals.

So then we see a story about the guy who loved the buttery smell of microwaved popcorn and got a rare lung disease, bronchiolitis obliterans. More to the point, "popcorn lung" has been added to the list of industrial diseases affecting production workers.

All I had were a few suspicions, of course. Nothing to go on, really. But I can't say that I'm surprised in the slightest.