Unbending the bends

Sometime before dawn this morning, we took our oldest son to the airport. He's bound for the Caribbean for a pre-orientation trip for college (learning to sail with a team of other freshmen). They will get the chance to do a little snorkeling, but when his dad asked him about whether or not they'd be doing any scuba diving, he replied enigmatically,"There is no hyperbaric chamber in the Virgin Islands. They'd have to fly you to Puerto Rico, I guess."

My first response was to wonder how they would do that, given that most aircraft are pressurized to something around 10,000 to 15,000 feet, which would certainly exacerbate the bends - the outgassing of nitrogen from the blood, which can cause embolisms (blockages) in your blood vessels and painful swelling in your joints.

Henry's law governs the amount of gas dissolved in a liquid: the amount of dissolved gas depends on the external pressure of the gas. For example as the pressure of carbon dixoide increases, so does the amount of dissolved carbon dioxide. Some portion of that dissolved CO₂ turns into carbonic acid (H₂CO₃), and lowers the pH, which gives soda water it's characteristic bite. It also means that acidification of the ocean is a risk of fossil fuel burning, and the resultant carbon dioxide in the atmosphere. Climate deniers will say that there is no data linking CO₂ levels with changes in the ocean pH, suggesting it's because the oceans aren't plain water, and that this will complicate the chemistry. True. But your blood is pretty chemically complicated, and this is essentially the system that is used to control your blood's pH.

So why would flying make the bends worse? As the external pressure of nitrogen falls with altitude, more nitrogren comes out of solution in your blood stream and joints. Neither are places where you want more bubbles. If possible, victims of the bends are evacuated on planes that can be pressurized to lower altitudes (an expensive proposition, and one often not covered by travel insurance).

Bariatric chambers allow the external pressure to be increased, and then slowly decreased to prevent the formation of large bubbles. It can take several "dives" to assuage the symptoms. I sat with my mother while she underwent treatment in a hyperbaric chamber, it's not for those with claustrophobia is all I will say.

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Photo is from Wikimedia.

Red Dwarfs

A version of this was written as a guest post for an artist friend's blog.

If you see a colored compound in chemistry, you can almost bet that it will contain a transition metal. Though we think of metals as being a shiny grey hue (with a few exceptions, gold being one), metals are key elements in producing colors for artist. The visible frequencies of light are relatively low in energy, and conveniently correspond to the small gaps in energy that electrons can leap in metals (what chemists call d to d transitions). Cobalt blue, one of my favorite hues, is (as its name suggests) a cobalt salt: CoAl₂O₄. To get different colors, you have to use different metal salts. You can get a brilliant, though not long-lasting, yellow pigment using lead chromate, the same chrome yellow that Vincent Van Gogh made famous. Tweaking colors to get slightly different hues requires either mixing materials or finding a different salt altogether, the gaps that the electrons leap over when they absorb light aren't adjustable.

But there are other ways to capitalize on the properties of metals to create color. Red stained glass has been made for centuries by adding gold to molten glass and carefully controlling the temperature. The gold clusters together in small particles which then become evenly distributed and suspended in the glass.

These tiny clusters are called nanoparticles, because they are 100 nanometers or less in size. One nanometer is 1 billionth of a meter, the period in this sentence is about a million nanometers across, the little gold balls in red glass are about 25 nanometers in diameter. (The prefix nano, comes from the Greek word for "dwarf," hence the title of this post.)

The gold nanoparticles are not dissolved in the glass, but form a colloid. And one property of colloids is that they scatter light. Different frequencies of light scatter differently, which is why the sky is blue, though the scattering of light by a colloid is a slightly different process. (Scattering isn't the only process involved in the color, but unless you really want to fly off the math cliff with me, let's leave talk of quantum dots and wavefunctions to another day.)

The color of light that a colloid scatters depends on the size and shapes of the particles dispersed. It turns out just by varying the size and shape of the particles involved you can tune your gold nanoparticles to be red, red-violet or even green and many colors in between!

If you are interested in knowing more about the history and chemistry of color, Bright Earth: Art and the Invention of Color by Philip Ball is a terrific introduction. He has a recent blog post about color here. For a readable introduction to nanoparticles, quantum dots and color, try this article in the NY Times.

Handheld chemistry

There was a time when chemists regularly reported the taste of newly synthesized compounds as well as other physical data (density, color, etc.). There was also a time when chemistry kits suggested doing chemistry in your hand, for fun. For a piece I wrote for Nature Chemistry (Homemade chemists) I found these instructions in a 1937 manual for a Chemcraft chemistry kit:

I'm a little cautious about using calcium oxide (CaO) as the reaction when it comes in contact with water is famously exothermic (you can cook an egg with it, see the video, and back in the day transporting CaO, or quicklime, by wooden ship, was hazardous duty). I wondered how exothermic was this reaction, and how much ammonia did it make relative to what you might encounter in a barn (the breakdown of urine yield ammonia) or your cat's litter box.

I'll admit to using Hess' law for fun. For those who have not enjoyed (endured?) an introductory chemistry class, Hess' law makes use of the fact that the energy content (heat of formation) of a molecule is a state function. Like altitude, it doesn't matter how you get to the top of the mountain from the valley, climbing straight up the side or meandering up a series of switchbacks, the change in altitude remains the same. So if I know where I am starting (the reactants, in this case calcium oxide (CaO) and ammonium chloride (NH₄Cl)) and where I end (the products, calcium chloride (CaCl₂, ammonia and water), I can figure out how much energy is used up (endothermic) or given off (exothermic).

The handheld reaction is 2 NH₄Cl(s) + CaO(s) → 2 NH₃(g) + H₂O + CaCl₂(s). I looked up the heats of formation in a handy table. To get a sense of magnitude, for 60 grams of CaO, which is about a tablespoon of material, the heat of formation is -635 kJ...or about the same amount of energy you can get from eating 3 Oreos. Overall, this reaction needs about 100 kJ to use up those 60 grams of CaO, in this case the energy comes from your warm hand. [Ed. note: While handheld chemical synthesis is an interesting way to "burn" calories, this is not a recommended weight loss technique!]

So your hand won't melt. Good to know. But if it were me, I'd do this in a test tube and warm it with my hand!

What the reaction does produce a surprising amount of ammonia. If you let the reaction go to completion (and since I don't know how fast the reaction proceeds, I can't tell you how long that will take), using about a 1.5 grams of ammonium chloride, and all the ammonia stays in a 1 cubic meter area around your hand, the concentration would be about 450 ppm. Since the CDC considers the IDLH (immediate danger to life and help) for ammonia to be 300 ppm, this would not be a great experiment to try in the tiny basement bathroom I used as a lab when I was a kid. Still, if you did this just until you could smell the ammonia, for most people that is about 50 ppm, a level considered reasonable for a brief (less than 5 minute) exposure. Levels inside a barn might be around 120 ppm.

Wash those hands.

Chemistry by accident

I just finished another Thesis column for Nature Chemistry, this one on the notion that chemistry sets are an essential part of turning kids into chemists — more particularly, what I called the Uncle Tungsten trope: risky chemistry is more fun and makes better chemists. As part of the article, I wondered how many accidents there are in home labs (not counting home meth labs). It turns out that in the US, the Agency for Toxic Substances and Disease Registry (ATSDR) keeps track of hazardous substance events. The data suggests there are around 1000 chemical incidents in private homes each year, and the vast majority involve carbon monoxide (nearly all the fatalities are caused by CO) or inappropriate mixing of common household chemicals (usually of bleach and something else: ammonia, pool acid, pesticides). As far as I can tell, none of the accidents were part of amateur chemistry gone awry.

There are no narratives linked to the data, but a chemist can read between the lines. When the primary chemical listed in a chemical accident is sucrose — table sugar — (a) what is the secondary chemical likely to be? (b) What was the intended goal of the experiment?

Answers: (a) potassium nitrate (or potassium chlorate) and (b) solid rocket fuel (or sparklers or smoke bombs or...). Sucrose oxidizes readily (toasted marshmallows, anyone?), and potassium salts (KNO₃, KClO₃) are good oxidizing agents.

It should go without saying, but do not try this at home. Especially do not try mixing bleach with anything. It will not make a stronger cleaner, bug killer, or weed killer. But it might kill you.

Cordial Chemistry: Syrup of Violets

Today's talk at the Chemical Heritage Foundation was by one of my fellow Fellows, Rebecca Laroche, on syrup of violets and Robert Boyle. It had long been known that adding an acidic material, such as lemon juice, to syrup of violets turned it a rose color. (More creepily, kids apparently used to hold pansies, also a member of the viola family, over ant hills to watch them change color, presumably from the formic acid produced by the ants.) Boyle is credited with the discovery that this botanical extract also changed color when exposed to alkalis, turning green (see his report here). This led to the development of a panel of pH sensitive indicators, helpful in chemical analysis in Boyle's time and now.

The color changes are due to the anthocyanins in the violets (the same thing that makes red cabbage change color with pH). Syrup of violets is not hard to make, you can find a modern recipe here, not much changed from the older recipes (see an assortment here), and you can buy it.

After Rebecca's talk a group of us went to lunch and, quite serendipitously, on the menu were drinks made with syrup of violets. Since some of us had writing to do this afternoon, we eschewed the vodka versions, but gave the club soda tonics a whirl. I wanted to see what happened when you added acid, would I get a pale rose drink? Alas, it seems not.

Turns out that commercial syrup of violets has citric acid added to it, which turns the pure syrup red, or it would if artificial colors were not added to make it violet again. Since it's already in the red form, adding more acid doesn't change the color.

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Twisted Threads: The history of the LCD

Learning to tell time when I grew up was a challenge. Clocks were analog - not digital. Everywhere. I can still see the little stiff pink paper clocks we were issued in first grade, with a brass brad fastening the hands to the face. We practiced setting the hands and reading off the time. Thanks to LCDs (liquid crystal displays) my sons had a quantiative sense of time much earlier, digital clocks blinked at them in every corner of their lives.

It took quite a bit to turn the initial discovery into a technology so smoothly integrated into modern life that we rarely notice it's there (how many LCD screens are in the room where you are now? Don't forget the ones in your pockets...). You can read more about the history of the LCD in this blog post by Ben Gross, a fellow at the CHF where I'm currently a short term fellow.

One of the pivotal developments was the leveraging of the twisted nematic effect...which made me wonder what worms (nematodes) and my iPad might have in common. The Greek root of threads....

The most Zen of molecules

Chemists are the Zen masters of science. Chemistry is a minimalist art. Its structures and mechanisms resemble the spare ink characters which trickle down scrolls. We seek elegant syntheses in which a few, carefully chosen pieces collapse into a whole. There is particular pleasure chemists take in crafting a molecule that strains the bounds of possibility — such as cubane — which evokes the aesthetic of Noh, where nearly impossible movements are made to look effortless. And despite our abilities to peer into the depths of a molecule with lasers or beams of neutrons, we haven't lost our connection our history. We are still distilling and crucibles are not merely historical artifacts. Zen sees a beauty in the old and well-used, a touch of wabi.

I've a piece in this month's Nature Chemistry on what makes a molecule beautiful (here, $), through the lens of the ten molecules that I consider to be most beautiful. I've already had a couple of emails suggesting gorgeous molecules that didn't make my list. What's on your list of elegant molecules?

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My list of the ten most beautful molecules

azulene
carvone
ferrocene
ethanol
vanillin
penicillin
insulin
snoutane
cubane

An element by any other name would smell as sweet

Elemental naming was as fraught in the 19th century as it can be today (though now the IUPAC has rules and committees). Alternate names and symbols for elements persisted not merely for decades, but in some cases more than a century.

I've recently skimmed a number of articles about glucinium (Gl). Not familiar? It has 4 protons and these days is known as beryllium for the gemstone beryl, in which it can be found. Beryllium salts can taste sweet, hence glucinium. Beryllium was suggested early on an option, since the sweet taste of its salts was not a unique characteristic. Other metals, including lead and yttrium, form sweet tasting salts. Still, in 1890 many authors were insisting that glucinium was the preferred name, suggesting that the arguments were continuing nearly a century after the initial discovery. It took more than 150 years for the chemistry community to settle on beryllium.

Other elements have endured dueling names, including colombium (now niobium) and the sounds-too-awkward-to-be-real jargonium (hafnium!).

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In searching for an appropriate image, Google turns up lots of bathtubs, including this one. Not only does an antiquated elemental name appear in the description of this wild tub, but the term angstrom as well. Translation software, I'm sure, but what is being (mis)translated?

And I couldn't resist the post title, as one of my fellow Fellows at the Chemical Heritage Foundation is a Shakespeare scholar.

Elemental tales: Strong waters

I ran across a reference to aqua fortis in one of the commentaries in Chemical News (1891). The conversation is about a suit in court where a chemist was injured when an inappropriately packaged bottle of aqua fortis spilled. (It had a cork, and according to the rather snarky commentator, the judge — and the chemist in question — should have known that aqua fortis should not be capped with a cork.)

Aqua fortis, literally strong water, was once the common name for nitric acid. Concentrated nitric acid is a strong oxidizing agent (I can still see the small scar on my mother's hand from a spill in her undergraduate days), and I imagine would rather quickly eat away any organic matter, such as a cork. Glass would obviously be the preferred medium for storage. The suit is a frivolous one!

The term aqua fortis has fallen out of fashion, but its companion term has not: aqua regia, the royal water that would dissolve even gold. Aqua regia, as any general chemistry text will tell you, is a mix of concentrated nitric acid and concentrated hydrochloric acid (a 1:3 ratio by volume). Neither acid alone with dissolve gold (or a variety of other hard to oxidize metals), but the trick lies in the shifting equilibria.

Nitric acid is able to oxidize small amounts of gold, turning elemental gold into ions, Au³⁺. These ions then react with the chloride ions from the hydrochloric acid to form the complex ion AuCl₄^—. As the gold ions are pulled into the chloroaurate complex, the nitric acid oxidizes a bit more elemental gold. This goes on until all the solid elemental gold has been turned in chloroaurate ions floating around in solution. Imagine putting out a bowl of pretzels, as the pretzels get eaten, you try to keep it full by adding more pretzels. Eventually you run out of pretzels. The trick of using complex ion formation to drive something that isn't very soluble into solution is a common one.

Arguably the most famous example of this happened when the Nazis invaded Copenhagen. Franck and von Laue had given their 23 karat Nobel prize medals to Bohr to prevent the Nazis from confiscating them. Bohr was reluctant to bury them, sure that wherever they were hidden, a search would eventually turn them up. A chemist on staff, de Hevesy, thought to use aqua regia to dissolve the medals. After the war the gold was precipitated out and recast into medals; Franck received his recast medal in the early 1950s. Those were strong waters indeed that Bohr and de Hevesy waded into.

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You can read a bit more about the saving of Franck and von Laue's medals at the Nobel site and see a video of aqua regia in action here.

Recording science

Bruce Gibb mused in a Thesis column in Nature Chemistry a few months back about taking small chunks of time to tune up your research apparatus. I'm on sabbatical leave this semester, and in addition to the research projects I've got going, I'm trying to devote some time on a regular basis to just this. I'm playing with an simple animation app, that would let me quickly put together animations for research talks or classes — and test driving apps for electronic research notebooks.

As a computational chemist, I've been balanced on the knife edge of digital record keeping my whole career. What goes into paper archives (hand kept, or printed), what stays electronic? Who backs stuff up, how often? Long term storage? I've encourage my students to think about how they want to track their data and, at least as importantly, their thinking about their data. Through it all (from punch cards to mag tape to memory sticks) I've always kept at least some of my work on real paper, in a traditional hardbound notebook. In ink. Dated. You know the drill.

I've been reluctant to let go of pen and paper. Just as I still outline just about any piece of writing, including this one, on real paper, I find I think differently off the keyboard. Keyboards tend to enforce a certain linearity of thinking, while a sheet of paper (or several and lots of stickies) lets me move into multiple dimensions, with fewer restrictions on insertions and more flexibility in formatting.

The work I'm doing now in the archives is facilitated by having photos of what I'm reading, many of the bound copies are too fragile to routinely scan or photocopy. Ironically, reading 19th century journals has catapulted me into the 21st century as far as my own record keeping is concerned. I'm using an integrated notebook app on my iPad which allows me to scribble and sketch by hand, take and incorporate photos (and mark them up if I wish), and input text from the keyboard. Finally, I can tag pages, and filter the notebook by tags (more consistent than my own hand written indexing procedures). The only thing I don't care for is that I can't write as small as I wish, making it harder to get an overall view of where I'm going. It's an experiment still,

Today's Nature [Nature 481, 410(2012)] has an editorial and an analysis piece on digital record keeping in science. One scientist notes that paper has nothing to offer her - she's gone entirely to her iPad. I may be right behind.