Surfaces

Something in the smog biz that used to drive me nuts was when someone would look at some smog chamber experiment that had some unusual feature to it and remark, “Well, that’s just a chamber effect.” The subtext was “We’re studying gas phase kinetics, and that’s something having to do with a surface phenomenon, so we shouldn’t pay any attention to it.”

I didn’t think that should let us off the hook. What kind of surface effect was it? How did it behave? And were we absolutely sure that such effects didn’t occur elsewhere?

Eventually I wrote a paper, “Background Reactivity in Smog Chambers.” Google scholar tells me that it’s been cited at least 17 times, as recently as last year, so it did okay for a paper published 20 years ago.

In the 60s and 70s, there were a lot of smog chamber experiments done on all sorts of individual compounds; there was a belief that one could produce a “reactivity scale” that would let you reduce those things that had the most smog forming potential. As the complex nature of smog chemistry began to dawn on people, such experiments became less common, because “reactivity” has multiple components, sometimes 2 + 2 = 6 in smog chemistry, making the development of a single scale problematic. There’s a fellow at SAPRC in Riverside, Bill Carter, who has developed a much more complicated way of estimating “incremental reactivity,” which has its own problems, but it’s better than “one size fits all.”

Anyway, one of the “pure compound” experiments involved methyl chloroform, and I found it fascinating.

Methyl chloroform is also called 1,1,1 tri-chloroethane. If you start with ethane (CH3CH3) and replace all the hydrogens on one methyl group with chlorine, you get methyl chloroform. It’s pretty unreactive stuff; the only reaction sites for hydroxyl radicals are the ones on the methyl group and methyl hydrogens are bound pretty tightly. So for the first part of the chamber experiment, using very high concentrations of MCF with some added NOx, the thing just sat there.

Then, after a couple of hours of induction, something began to happen. The NO began to convert to NO2, some of the MCF began to decay, then suddenly, wham! The whole system kicked into high gear, NO went down like a shot, the MCF began to oxidize like crazy, and ozone began to shoot up. Then, just as suddenly, the ozone just disappeared, all of it, in just a couple of measurement cycles.

Everyone who looked at it said, “Ah, chlorine chemistry,” which was a sure guess. Chlorine will pull hydrogen off of even methyl groups with almost collisional efficiency (if a chlorine atom hits the molecule, it pulls off the hydrogen almost every time). Moreover, chlorine atoms destroy ozone; that’s the “stratospheric ozone depletion” thing.

But I was puzzled. Where did the chlorine atoms come from? Yes, there was plenty of chlorine in the MC, but that was bound. To get one off, you need to create a free radical and those ain’t cheap. If you create an HO radical, that can pull off one of the hydrogens, and that, after the usual reactions, gives you chloral, a tri-chlorinated version of acetaldehyde. Put in a high enough rate of photolysis for chloral in your simulation and you can get the whole system to react.

The problem was, it didn’t look right. With a high rate of photolysis for chloral, the simulation kicked off too quickly. Lower the rate and you never got the sudden takeoff. I’m pretty good at fitting the curves, and I could never get it to work.

So I started looking at the other actors in the system. The end result of chloral oxidation is phosgene (see why I was looking up all those post-WWI gas papers?), but phosgene itself didn’t fill the bill. So maybe the phosgene was converting to CO and Cl2 on the chamber surfaces like it does in someone’s lungs. No, that didn’t work either.

I kept returning to the problem over the years, trying yet another idea, each time getting no further.

In 1985, the “ozone hole” over the Antarctic was reported, and everyone in the stratospheric ozone community, including Gary Whitten, my boss at SAI, immediately suspected that it had something to do with the ice clouds that only form in the stratosphere over the Antarctic. In 1987, Mario Molina published a series of papers describing the surface reactions of stratospheric chemical species on ice crystal surfaces. The really critical reaction was the reaction of chlorine nitrate with hydrochloric acid to form nitric acid an molecular chlorine (Cl2). Cl2 photolyzes so rapidly that it might as well be two chlorine atoms.

I’m not sure when I first tried the Molina reaction on the methyl chloroform system, but it worked much better than anything else I’d tried. It makes the whole thing a very strong positive feedback system. It worked well enough to convince me that it was probably the missing factor; if I wanted to get a better simulation, I’d have to get very specific about some details of the original chamber experiment, and that one’s 35 years old. It’s pretty well moot at this point anyway.

Molina won the Nobel Prize for his work on stratospheric ozone depletion, and it was well-deserved. I was just looking at a single smog chamber experiment, one with a surface reaction that no one was interested in. The chance that I would have figured out the right answer to the peculiarities of that experiment is pretty small. The chance that I would have made the leap from the chamber walls to the stratospheric ice clouds is smaller still; I’d never heard of them before Whitten told me about them, and I certainly didn’t make the connection between them and the chamber experiment until Molina worked out the correct surface chemistry. So I’m certainly not trying to say that I coulda been a contenda.

But I will say that we all should have been paying more attention to the chamber wall effects. You don’t get to say beforehand what will turn out to be important.

Swimming Pool Air

In “The [Widget], the [Wadget], and Boff, Sturgeon mentions in passing a plethora of alien civilizations and lifeforms, including “fluorine fellowships.” Alliteration aside, the likelihood of a lifeform being based on fluorine in any significant way is very slim; the likelihood of an environment having free fluorine in it is so close to zero that I’m saying it’s impossible. Fluorine will burn water, so the most you could get would be hydrogen fluoride, and that will etch glass, so your putative fluorine world would need to be free of silica. Good luck on that one.
Chlorine, on the other hand, seems more likely. It’s actually a bit less electronegative than oxygen, so it isn’t obviously less likely to exist in a free state. But an atmosphere of chlorine has a serious problem: sunlight.

Chlorine photolyzes all the way down to yellow light, the chlorine molecule splitting into two chlorine atoms. And chlorine atoms are really promiscuous, especially to hydrogen. While they won’t strip hydrogen from water, they will do it for any hydrocarbon, including methane, as well as molecular hydrogen. So you’re really talking about an atmosphere of hydrogen chloride, i.e. intensely acidic.

We use chlorination as a disinfecting process for drinking water and swimming pools, with the latter getting a lot more chlorine than the former. Anyone who has ever had to maintain an outdoor pool can tell you that a sunny day can chew up your chlorine mighty fast. There’s a trick you can use to get rid of the chlorine taste in over-chlorinated water, which consists of just putting the water in a glass jar and leaving it in the sun for an hour. Then you can make your coffee or tea from it without the chlorine taste.

Part of the chlorine in swimming pools goes to hydrochloric acid, which will acidify your pool. We used to use sodium hypochlorite at the old YMCA pool where I was a lifeguard, and the sodium took up some of the acidity. Nevertheless, we sometimes had to add some sodium carbonate to the mix.

Chlorine doesn’t just kill bacteria though; it chews them up and grinds them down into small molecules, including ammonia and organic amines. Then you get chloramines which are even more toxic than chlorine itself. A substantial part of the “swimming pool smell” that most people think of as chlorine is actually from chloramines. For us former life guards, a little whiff of chloramines (or even Clorox) triggers memories.

A “chlorine producing photosynthetic algae” like the one mentioned in “Tranquility” and in Steve Gillette’s World Building articles (and book), probably wouldn’t affect a planetary atmosphere very much, at least at first. But it would begin killing sea life pretty quickly, in a kind of “green tide.” Depending on how it sequestered the alkali elements (I had in mind carbonate in a silica diatom shell), it would also raise the pH if the upper water layers. That would disrupt the carbonate-bicarbonate balance of the sea surface, releasing CO2 into the atmosphere while also depleting surface CO2 as the biologically sequestered carbonate precipitated down the water column. Oceanic primary productivity would plummet, while atmospheric CO2 would skyrocket; a combination of global climate change and a die-off of oceanic life that I figured as apocalyptic.

But of course no one would ever do so foolish a thing as to tamper with the global environment just out of curiosity, or while trying to make a buck, so I'm sure we're safe from this particular danger.

Paying Attention

Chlorine gas is quite the little charmer, from a photochemist’s point of view.

Most gases need ultraviolet light to get frisky, but not chlorine. No, Cl2 absorbs into the mid-range of the visible (somewhere around yellow, as I recall), and does so pretty emphatically, photolyzing at several percent per minute in full sunlight. When it absorbs light, it decomposes into two chlorine atoms, radical atomic chlorine, which is mega-reactive, yanking hydrogen atoms off of even methane at a pretty fast rate. For larger hydrocarbons, the rate is “collisional” which means that when the Cl atom hits, the probability of reaction is 1.

So Whitten and I got interested in chlorine because it can goose the smog process, and it was probably doing so in Henderson, Nevada, and we had a small contract to assess the situation. Henderson had been having smog episodes on cold December mornings, when the thermal inversion was very low (trapping the pollution very near the ground). Usually smog does not go so well in the cold, but in Henderson, apparently, the chlorine trumped the cold. The chlorine came from a company called Timex, which made titanium, not watches, with a process plant that was new in WWII, and leaked like a sieve.

As part of the work, I did a literature study on chlorine photochemistry, which was interesting. The “smog accelerator” effect from chlorine had been used since the mid-sixties as a way to enhance radical initiation in smog chamber experiments, and after that, the phenomenon of ozone destruction in the stratosphere had led to a fair number of papers.

But before the 1960s, there wasn’t very much about chlorine photochemistry in the literature—until I got back to the 1920s and 1930s. Beginning around 1920, there was a sudden boom in the study of chlorine in the gas phase. That boom lasted until the mid- to late 1930s and then trailed off. By the time of WWII, it was a moribund field.

It takes no great insight to realize what caused the sudden interest. The Great War had left a wake of disabled veterans who had been victims of gas attacks, chlorine, phosgene, and bis-(2-chloroethyl) sulfide (“mustard gas”). I also got interested in phosgene chemistry a while after our Henderson study (phosgene also contains chlorine, and it’s a decomposition product of chlorinated hydrocarbons like methyl chloroform), and all the information I could find on it also came from this time period.

This is, of course, a very ordinary description of the ordinary behavior of science, or, indeed the ordinary behavior of human beings. We study things that interest us; we pay attention to our interests.

Some of the more extreme advocates of relativism and deconstruction claim that science is an enterprise no different from any other, and that the results of science are conditioned/determined/biased/required by the social, cultural, or gender outlook of the scientific community. One can, of course, deconstruct such claims by noting that they are made by one class of university academics who are locked in the usual struggle with a different class of university academics, but that is a petty, or at least trivial insight.

Scientists generally hold themselves to be “reality based,” to use another culturally biased phrase and I strongly agree that no amount of cultural relativism will make N-rays exist and X-rays imaginary. But that is not the be-all and end-all of bias and cultural determinism. No, the phrase “where one’s interests lie” says a lot about science. We study chlorine in the 1920s and 1930s, not because of some rational, objective criterion of how chemistry should progress, but because cultural factors made it important to do so. I can point to a number of objective truths that my colleagues and I uncovered over the course of my career in science, but it would be folly to think that we were not looking “where the light’s better,” under the streetlights of funding, social interests, and group status within our little part of the scientific community.

If I had to make my own philosophical stand, holding up my own interpretation of both free will and the forces that work against it, I’d give that as a capsule summary: it has to do with paying attention, to what, and for what reasons.