The Who, What, When, Where and Why of Chemistry
Chemistry is not a world unto itself. It is woven firmly into the fabric of the rest of the world, and various fields, from literature to archeology, thread their way through the chemist's text.
Before chemists became adept at synthesizing and purifying single molecules, materia medica relied heavily on plant based materials. The chemicals in plants are not uniformly innocuous, or safe at any dose, a point I tried to make in this article at Slate a couple of weeks ago. A case in point: St. Ignatius' beans.
Last fall, I was digging through a 1903 organic chemistry text (looking for examples of eponyms for this article), when a familiar name caught my eye. What was St. Ignatius doing in a chemistry textbook, an organic one at that? Jesuits, I could understand (quinine is extracted from cinchona, also called Jesuits' bark), but Ignatius (the founder of the Jesuits) himself?
"Strychnine, C21H22O2N2, is found in St. Ignatius' bean..." What is a violent poison doing in a bean named for Ignatius? Despite the fact that I was up against an impending writing deadline and had a couple of dozen exams to grade, I had to know. Faba Sancti Ignatii were first described by an Austrian Jesuit living in the Philippines in the 17th century, George Kamel, S.J. (his description was published in the Philosophical Transactions in 1699 - and yes, I looked up the Latin version). Later authors speculated the plant was named for Ignatius because of its many medicinal virtues (which they do not list). At the turn of the last century strychnine was part of the US Pharmacopoeia, prescribed as a stimulant — it was implicated in a early Olympic doping scandal — and for gastric upset; in the Phillipines it was often (more sensibly) the bean was worn on a string around the neck for protection against various diseases. These days it forms the basis for a homeopathic nostrum prescribed for grief and melancholia, particularly when associated with an abundance of tears.
I'm at ScienceOnline2013 where Carmen Drahl and Dr. Rubidium just finished running a terrific session on chemophobia: how can we bridge the gap between "better living through chemistry" and ads for "chemical-free sleep aids." The thrust of the session was not how to convince people chemistry and chemicals are good, but more about how to inject nuance into the public conversation. Chemicals have risks and benefits — and of course, are unavoidable. But we current view chemical as synonymous with toxic, hazardous, unnatural or just plain bad.
What are the roots of this cultural shift? Can understanding these help scientists and writers communicate more clearly and in the end help people not only understand what is in their "stuff" — chemicals, it's all chemicals — but give them tools to work with and make decisions about the materials that make up the world — chemicals. As @docfreeride (ethicist Janet Stemmwedel) noted at another session yesterday, we can agree on facts, and still make different decisions based on them.
Today's New York Times has a perfect example of the various ways chemophobia presents in the Magazine: The Boy with a Thorn in His Joints. The piece chronicles Susannah Meadow's search for an effective treatment for her son's rheumatoid arthritis. She agonizes about the decision to give him methotrexate (which in high doses is used in anticancer treatment) and turns to alternative treatments, in particular four-marvels powder. There are intense arguments with the pediatricians and with her husband over the issue. I was struck by two things in this piece. First, the language Meadows uses to limn the controversy, and second her ignorance, not so much of the chemistry that is in your face (methotrexate), but of the ways in which chemistry is couched in alternative cultural schemes(four-marvels powder).
It makes me wonder how chemophobia is linked to the language we use to talk about it. It can be nearly impossible for an non-chemist to figure out what methotrexate is (beyond "a chemical"). The very name sounds harsh. Four-marvels powder is easy to parse: a powder with four effects. Its name rings with hope.
I also wonder if we worry more about stuff we are familiar with, we've heard more talk on the street about their risks. So we obsess about vaccines, because we hear and read about the side-effects of vaccines, but how many people know anyone who has died of measles? (One of my sister's friends died of measles when I was a child, before there was a vaccine.) So we get in the Times' piece "I was desperate to find a way...without the drugs." pushed up against "[My husband] has always been more comfortable with pharmaceuticals, more trusting in general."
Of course, four-marvel powder is a pharmaceutical, it's just from a different pharmacopoeia — the traditional Chinese — than the one Meadows or her husband is familiar with. Meadows can read the package insert with information on the side-effects of methotrexate, she may be unaware of the routine advice given in Chinese medicine programs (and yes, there are formal academic programs in Chinese medicine, e.g. at Nanyang Technical University) about four-marvels powder (it should never be given to pregnant women, for example, which might make you hesitate before giving it long term to infants or young children).
The session at SciOnline2013 brainstormed about effective ways to help people develop a better sense of nuance around what is a chemical and what are the risks of this particular chemical? What strategies do you think would be most effective?
Most winters Punxatawney Phil is the furry face of Pennsylvania, but last year, he had competition: meet the purple squirrel of Jersey Shore (which should not be confused with either a television show or a town in New Jersey).
The news report offers a number of theories about the squirrel's unique coloration. A dye job seems the likely culprit, whether from the squirrel's nesting material or an inadvertent bath in a violet solution. Computer scientist Krish Pillai had a novel suggestion: "This is not good at all. That color looks very much like Tyrian purple. It is a natural organobromide compound seen in molluscs and rarely found in land animals. The squirrel (possibly) has too much bromide in its system."
Leaving aside that Tyrian purple (produced by a particular class of marine snail and to the best of my knowledge and research abilities by no mammal) is a much redder color, this assertion is roughly equivalent to saying that if I eat too much chloride, say from table salt, my body could start synthesizing Splenda, an organochloride. No, just, no.
Pillai is apparently extrapolating from reports that bromide (bromine anion - Br-) has been found contaminating wells near fracking sites. Calcium bromide is used in drilling fluids to increase density, by some estimates 20% of the bromine used in the US ends up in "clear brine fluids" — mixtures of various bromides. But it is a long way from bromine ions to 6,6′-dibromoindigo along very specific biochemical pathways. Which squirrels don't have. Or humans. (What can and does happen is that the bromide reacts with various chlorine compounds used in water purification to form organohalides, which aren't healthy to ingest....)
It's worth noting that direct ingestion of dyes can have interesting effects on pigmentation. Flamingos get their characteristic color from ingesting shrimp pigment, and you can change the color of a canary's feathers by feeding it paprika. Humans who eat too many carrots can develop carotenemia — they turn orange. These processes are reversible, stop eating the shrimp or carrots and feather or skin return to their normal coloration. Unfortunately consuming silver or gold can produce a permanent change in skin coloration, as in argyria.
An alternate definition of a purple squirrel via Urban Dictionary.
Grasping the abstract models chemists use to understand what holds a molecule together — its bonding — is one of the major goals of the general chemistry course I am teaching this semester. Understanding the bonding in a molecule is the key to predicting and understanding its structure and reactivity. The models chemists use to describe bonding in molecules range from what can be done on the back on an envelope, such as Lewis dot structures or VSEPR structures, and those that require hefty amounts of computer time to set up and solve (ab initio MO theory). The text we are using includes many full color diagrams of bonds, but student still struggle with how these two dimensional representations "work" in three dimensions.
Ad hoc models — mock ups of molecules built by hand from mundane materials such as cardboard and wire — have a venerable history in chemistry. Watson and Crick used cutouts of the bases to figure out how they paired along the helix; Smalley built a paper model of pentagons and hexagons to see how C60 could be constructed.
So last week, I brought paper, tape and some simple molecular models (tubes and small metal centers which I buy by the bag to hand out to students) to class and asked students to build a valence bond model for acetaldehyde (implicated in hangovers - acetaldehyde, not valence bond theory, though the latter certainly can make students queasy). Students cut, paste, built and discussed, producing what you see in the slideshow.
With thanks to Danqui Luo, Tess McCabe, Kai Wang, Ben Kaufmann and all the students of Chem 103 who built and photographed the models.
Hurricane Sandy left us without power for several days and while a basement chest freezer remained solidly frozen, thermal equilbrium was unfortunately reached by our refrigerator and kitchen, at roughly 55oF. Saturday morning found us rooting through the refrigerator, deciding what had to be chucked (milk) and what didn't (ketchup). Butter? In this cool weather, it could stay, it would be unlikely to have turned rancid.
But coincidently, while breezing through a depression era Chemcraft chemistry set instruction book, I encountered directions for "renovating" rancid butter. Around the same time that margarine made its debut, so did process butter, butter that had been treated to remove the objectionable materials. As near as I can tell, it's an extraction process, presumably the rancid materials (such as butyric acid) dissolve in the cream and the remaining materials can be reworked into a solid mass.
Laws remain on the books in many places forbidding the sale of process butter without making clear to the consumer what is being purchased. In the early part of the 20th century this was widespread enough for the US Department of Agriculture to print a booklet which "enable[s] any housekeeper, with only the usual facilities of the kitchen, to distinguish in the great majority of cases between genuine butter, renovated butter, and oleomargarine."
Next time the power goes out, I'll know how to "renovate" my butter, as long as I don't intend to sell it!
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.
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 Colorby 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.
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.
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.
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.
