The Culture of Chemistry

Exploding expertise

2013-05-06T14:01:00.002-04:00

How do we decide who to listen to about something chemical? I have a piece in Slate this week (on the brouhaha around the teenager in Florida and the exploding water bottle), and someone in the comments feed there thought to comment on my expertise:

Jack Stephens

The author is apparently ignorant of chemistry. The active ingredient in toilet bowl cleaner is not hydrochloric acid, it is sodium hydroxide. Aluminum and sodium hydroxide react to form hydrogen gas.

16 Hours Ago

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Franz Liebkind

Some toilet bowl cleaners (e.g. Lysol's) contain hydrochloric acid in 10-ish percent concentration. It is there to dissolve calcium carbonate deposits (i.e. scale) found in hard water areas. Both HCl and NaOH react with nonbulk Al to produce, among other things, H2 gas. Iirc in the HCl case the reaction should go faster.

Sodium and HCl (or just water!) is much neater and more spectacular. Adolescent pyromaniac curiosity inspired many of us to major in chemistry.

9 Hours Ago

from slate.com

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Franz Liebkind

P.S. USGS says that the water from the Upper Floridian Aquifer, which supplies most of Polk County and Bartow, is moderately to very hard.

I just finished a new Thesis column for Nature Chemistry about the ways in which chemists can (or cannot) communicate with general audiences about chemistry. Is it possible to have nuanced conversations using the word "chemical" and chemistry, or has the word itself chemical accreted so many toxic associations that it can't be rehabilitated? Can chemists have a role in these conversations by virtue of their expertise? (Short answers: Problably not, probably yes, probably not.)

A session at ScienceOnline2013 earlier this year still has me thinking about the disconnect between how chemists want to talk about their field ("did you know that everything is chemicals? just look at how this works! isn't it cool?") and how people process the information we are so enthusiastically providing ("she is a working mother who probably feeds her kids fast food five nights a week and can't possibly care about her family's nutrition so why should I listen to what she has to say about the molecular structure of NutraSweet™?") Addressing the deficit in science knowledge may not in fact help people assimilate what they need to know make informed decisions about things chemical.

For a society who in many ways is so keen on credentials (or how else do those online diploma mills spammers make money), social science research suggests we don't necessarily consider purely those credentials into our decision when we decide who is an expert in a given field. Dan Kahan and colleagues at the Culture Cognition Project suggest that we assess expertise through the lens of our cultural and social affinities as much (or more) as we do through objective credentials.

So when it comes to deciding who you should believe about aspartame, you believe Dr. X who is an "nutritionist, aspartame victim and single mother of three boys" (her doctorate comes from an unaccredited online school) not Dr. Y who does research on molecular structure and is the mother of two boys and is not an aspartame victim (her doctorate comes from a top accredited school).

As for the Slate commenter, I'm rather fascinated that someone could generalize from you don't know what is in toilet bowl cleaner (is the subtext here that I would be above cleaning my own bathrooms?) to you apparently know no chemistry. Could you imagine that a well-trained scientist would not think to look this up even if she doesn't do bathrooms? (The police report gives the brand of cleaner, which the manufacturer says contains 20% HCl. Of course, practically it doesn't matter -- both the acid and the base oxidations of aluminum produce three equivalents of hydrogen gas.)

All natural, locally sourced liquid nitrogen?

2013-04-25T23:08:00.000-04:00

Robyn Sue Fisher wants you to know that she would never cook with chemicals not found in nature. Smitten, her ice cream shop in San Francisco’s Hayes Valley, may at moments resemble a high school chemistry lab, but that’s because Fisher uses liquid nitrogen to freeze her product.

Nitrogen is “a natural element,” she notes. “It’s all around us.” [The original lead to this NPR blog post.]

I imagine not a few chemists reading this want Robyn Sue Fisher to know that liquid nitrogen is not found in nature on this planet. I suspect if she posted this photo of a cryogenic nitrogen plant in the ice cream shop, she'd have a hard time convincing her customers that liquid nitrogen was natural.

Her comments beg the question of what constitutes a chemical in the mind of a non-chemist. If we take IUPAC's Gold Book as the arbiter of the technical definition, a chemical is a material of "constant composition best characterized by the entities (molecules, formula units, atoms) it is composed of." Everyday language has drifted from the technical. The Oxford English Dictionary offers this definition for the non-technical speaker: "a distinct compound or substance, especially one which has been artificially prepared or purified."

Most people would agree that in common usage chemical carries the connotation of both artificial and noxious, while chemists attach no such presumptions as to source or toxicity to the term. Much as chemists wish it were not so, there is a growing language gap, and I think it unlikely we are going to regain the ground lost. Molecule still comes across as more neutral in tone to a non-chemist. So we are in a moment where we have people who are aghast at chemicals in their food, and others who are fascinated by molecular gastronomy (and likely some overlap in that population).

In principle, I do like the idea of locally sourced ingredients, maybe I should start my own shop and advertise that I use only locally sourced, artisanally produced liquid nitrogen?

I have to say I was also fascinated with how the NPR post morphed throughout the day in response to the comments on the blog. By the end of the day the introduction read:

Robyn Sue Fisher's ice cream shop, Smitten, in San Francisco's Hayes Valley, may at moments resemble a high school chemistry lab, but that's because Fisher uses liquid nitrogen to freeze her product.

Nitrogen is "a natural element," she notes. "It's all around us."

_____
Andrew Bissette has a good piece about chemophobia on Carmen Drahl's blog Grand CENtral today (In Defense of Chemophobia) which, along with this post from 2011 by Sciencegeist touch on the language issue.

(H/T to Fran who sent me the link to the original NPR post)

Doing the math around artificial sweeteners

2013-04-11T16:18:00.002-04:00

"The controversy surrounding these products natural, natural-like, or artificially made sugars will likely continue for years to come. We do know that artificial sweeteners increase our threshold for sweet taste, and yes, cause us to crave more sweets. If one “diet” soda leads to another “diet” soda, the “diet” effect is soon lost. " [source, emphasis is mine]

The 12 oz diet soda on my desk has 0 calories (meaning less than 5 calories under the FDA rounding rules). If I drank ten of them, I might consume 50 calories, at most; in all likelihood, far less. The diet effect is safe, I think.

I'm thinking about how we evaluate information: what does the BS detector look like for scientists versus non-scientists?

(As an aside, there is no evidence that artificial sweeteners increase the consumption of sweets.)

St. Ignatius' Beans: Strychnine and herbal remedies

2013-02-17T21:53:00.001-05:00

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.

A version of this post appeared at Quantum Theology.

Chemophobia: The Boy with a Thorn in His Joints

2013-02-02T14:56:00.002-05:00

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?

Will bromine turn squirrels purple?

2013-01-09T22:28:00.001-05:00

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.

Building Skills

2012-11-30T18:38:00.000-05:00

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 Chemistry: Renovating Butter

2012-11-07T21:37:00.000-05:00

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 55^oF. 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!

Unbending the bends

2012-07-25T14:38:00.004-04:00

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.

Photo is from Wikimedia.

Red Dwarfs

2012-07-20T11:00:00.006-04:00

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

2012-07-18T15:59:00.006-04:00

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, due out in September, $) 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

2012-07-10T14:26:00.003-04:00

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

2012-03-13T21:59:00.005-04:00

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

2012-02-23T14:29:00.005-05:00

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

2012-02-21T16:16:00.003-05:00

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?

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

2012-02-13T11:40:00.003-05:00

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!).

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

2012-02-07T16:10:00.004-05:00

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.

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

2012-02-06T16:49:00.003-05:00

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.

Ephemeral Elements

2012-02-02T17:49:00.005-05:00

The late 19th and early 20th centuries were hotbeds of elemental discoveries (literally and figuratively). New elements came — and on occasion went. Some were known elements in unknown guises, such as previously unrecognized allomorphs. Others, like didymium, weren't elements at all, but mixtures of as yet to be identified elements (in this case neodymium and praseodymium). Some were more ephermeral than others.

Yesterday I ran across a description of the discovery of a new element in an 1890 issue of Chemical News: damarium, oddly enough reported in the Notes & Queries section and not among the research papers. The report of the gaseous element, collected in Damara Land (present day Namibia) was a bit over the top, even for a time when flowery prose was in style in scientific papers: "One of the party had in his hat a branch of a shrub, which in a very short time lost its green colour and assumed a violet blue..."

One contemporary report assumes it is a hoax, but several sources were not so quick to dismiss the claim, particularly in a period when elemental identity was in flux. At least one commenter wondered if it might be "helium" — an element as yet undiscovered on earth.

I wonder if it's worth tracking down the original cite if I can (the Chemiker Zeitung is available on microfilm at the Othmer). Ah...Google books has it here.

German chemical humor or not? What do you think?

A curious invention: drawing chemical structures

2012-02-01T21:11:00.006-05:00

I am currently wending my way through fragile but fascinating volumes of Chemical News - a journal published by Sir William Crookes in the late 19th and early 20th century. It was a major journal at the time, looking rather like the current Nature in it's breadth of coverage. The society journals of the time typically reserved their pages for papers read by members and abstracts of papers thought to be of interest to them, while Chemical News and it's ilk included book reviews, reports of papers from a wide swath of journals in several languages and two robust arenas for conversation between scientists, readers and editors: Correspondence and Notes & Queries. They were a bit more open, too, to offer space to offbeat bits of science.

The volume I just finished (1890) has a rather contentious conversational thread winding through the Correspondence on what it means to be a FCS (Fellow of the Chemical Society) and should membership be more tightly policed vis a vis their chemical credentials. (At one point the secretaries of the Chemical Society accuse a former board member of having used fake letterhead to secure support for his position!) Many participants in the conversation resort to pseudonyms, some of which carry a bit of snark with them, and it's interesting that this controversy is playing out primarily in a commercial journal and not in organs internal to the Society.

My project involves tracking the correspondence around primary reports of research findings, so these raucous conversations, while fun reads, are of peripheral interest. I'll admit to finding other interesting tidbits to tag in my electronic notebook. It doesn't pay to be overly focussed when doing archive work, as long as I can avoid being completely dragged down the rabbit hole.

The Notes & Queries section appears just above the one page of adverts included in each issue, and yesterday this ad caught my eye: "The Benzene Nucleus. — An India-rubber Stamp in nickel-plated locket with ink-pad enclosed" 3s. At the top of the page, the last bit of editorial content appears — a report of a curious invention: a stamp for making benzene rings. The first benzene ring in a journal appeared in Chemical News (in 1879, eight years after the first graphical structure was used), so perhaps it's apt that it report this "little contriviance" in its pages. (And the inventor is a Fellow of the Chemical Society!)

Nowadays chemical structure drawing programs are commonplace, but when I was a graduate student chemical structures had to be hand drawn, using India ink (permanent, not water soluble!) on vellum. The Rapidograph pens used were expensive and notorious for getting clogged (irreversibly so). Rings were made using stencils, text added using mechanical lettering guides. Jiggle your hand and you had trouble that white-out might not be able to rescue you from. Blots? Argh.

I don't miss the days of chancy ink drawings for slide and papers, though I do miss the delight of pulling out pens and ink and paper. I do wonder, though, if note taking organic students would appreciate a little ink stamp of a benzene ring on the end of their pencil or pen?

Read about K&E lettering sets here.

Weird Words of Chemistry: Frigorific

2012-01-31T11:15:00.006-05:00

I ran across this word when my youngest, who I'm coaching for the thermodynamics event for Science Olympiad,asked me why the freezing point of water was 32^o on the Fahrenheit scale. The Celsius/centigrade scale was originally pinned to the freezing point and boiling point of pure water at 1 atmosphere of pressure. (Now it's pinned to absolute zero and the triple point of water.) What physical property was 0^o linked to? The freezing point of something other than water? I had to admit I didn't know and now that my curiosity was piqued, went off to hunt it down.

The zero of Fahrenheit's temperature scale was essentially pinned to the temperature of a "frigorific" mixture of ice, water and solid ammonium chloride in a 1:1:1 ratio, along with the freezing point of water and the temperature of the human body. Frigorific seems to have been coined by Robert Boyle to describe particles of cold that were transferred from body to body, and ultimately got attached to mixtures that achieved a particular temperature regardless of the starting temperatures of the materials. Wandering through the old chemistry literature, I found this table of frigorific mixtures "sufficient for all practical and philosophical purposes, in any part of the world in any season," useful in the days before refrigerators, still useful for those who need a constant temperature bath at low temperatures.

The size of a degree was set by bisecting the difference between the point at which ice and water were in equilibrium and body temperature six times, or 64 degrees (2⁶). Binary was easier to use when you had to make your own instrument than decimal.

Frigorific has essentially vanished from the chemist's vocabulary, though it's still apparently alive and well in the engineering literature. As words of science go, it sounds awkward to my ears — as roughly sharp as heaved Arctic ice.

Nova has an excellent piece on the hunt for absolute zero. Thanks, Kathryn J for the reference!

For more on what I think about well-formed science words, you can read "Neolexia" at Nature Chemistry.

What makes a molecule beautiful?

2012-01-25T14:06:00.006-05:00

I just finished a piece for the March issue of Nature Chemistry on what (in my mind) make a molecule beautiful. I will admit a preference for sparer, less baroque structures. (If you want to know more about my molecular aesthetic, you'll have to wait for the piece to appear!). In the meantime there is an article in this month's Nature Chemistry with the intriguing title "Quantifying the Chemical Beauty of Drugs" [Bickerton et al. Nature Chem. 4, 93-97 (2012), full text is free]. It's not so much beauty in the abstract these chemists are trying to quantitatively capture, but desirability. How attractive is this molecule as a target for drug development? Would a chemist be willing to surrender time and bench space to the synthesis of this molecule?

The model takes as its inspiration Lipinski's rule of 5. If most or all of Lipinski's five characteristics are present, a molecule has a good chance of being a viable candidate for an oral drug [Lipinkski et al. Adv. Drug Dev. Rev. 23 3-25 (1997)]. The goal is to develop an expert model system, one that mimics (or improves on) a chemist's intuition about what makes for a good drug.

Earlier work had suggested that chemical fashion sense is drifting toward more baroque structures for their drugs, despite various rule sets that suggest that bloated molecules are less likely to survive to clinical trials. Chemists apparently like their molecules "tractable" (which would seem to mitigate against molecular overelaboration?), synthetically and otherwise! Molecular docility is desirable.

For a somewhat darker take on chemical intuition and seat of the pants drug design read "Chemists in the Shadows" by Adam Piore in March's Discover Magazine. The article focuses on underground chemists who are developing new recreational pharmaceuticals that skirt current drug laws (steroids for athletes, and rave drugs). The conceptual framework used by some of these chemists would be familiar to any medicinal chemist (particularly in the early days, before QSAR).

Being a philosophess

2012-01-13T14:14:00.005-05:00

"...as a Philosophess she will not be discouraged by one or two Failures" Benjamin Franklin, in a letter to William Brownrigg dated 7 November 1773, where he wonders if Mrs. Brownrigg has succeeded in making Parmesan cheese (which I have to admit, I did not think was a cheese that colonial Americans knew of).

I appreciate Franklin's confidence that a woman could conduct rational experiments, particularly as at this moment I am virtually sitting on top of the site of Franklin's house in Philadelphia — I can see it from my window — working at being a Philosophess myself. I began a two month stay at the Chemical Heritage Foundation in Philadelphia today, as the Herdegen Fellow in the History of Scientific Information. My project is looking at how chemists, now and in the 19th century, deal with critical commentaries on the primary literature. Where are the commentaries located and does location change their tenor and/or content. I'm off to learn a bit about ways to computationally evalauate emotional tone, and to find some compelling narratives of critique in the 19th century and the 21st century.

I briefly wondered in my most recent Nature Chemistry Thesis column about what it meant for me to be working as both a historian of chemistry and a chemist, and how much of one field should we be exposing students of the other field to. Just how much history of chemistry does a chemist need to know to function well as a chemist? And if you do need to know something, what sorts of things? Dates? People? Materials? Methods? You can read my musings at Nature Chem, and those of Qian Wang and Chris Toumey on the same topic here. (Sorry...you or your institution need a subscription to see these, or if you would like a reprint of mine, drop me an email.)

Don't drink the water

2011-12-29T21:08:00.003-05:00

"Don't drink the water from the sink!” read a sign taped to the mirror. As I was in rather desperate need of a glass of water before rehearsing the piece I would sing solo at Christmas, I was glad to find someone had left a gallon jug of distilled water and a stack of paper cups. Rehearsing the next day, as I went to grab a cup of water, a colleague pointed out that yesterday someone had mistakenly put out distilled water, which he had swapped for spring water. “Hopefully no one drank it!” he said.

“Why not?” I inquired.

“You’re not supposed to drink distilled water.”

Ah. Yes and no.

Distilled water is water that has been boiled, trapped as steam and condensed, leaving behind the non-volatile impurities (the stuff that doesn’t easily turn into a gas, like metal salts). Other components, like alcohols can still be carried along into the distillate.

Distilled water lacks most of the ions that tap water has, and thus, much of its flavor. Some of the ions (such as fluoride) in regular tap water may have health benefits, so a steady diet of distilled water may deprive you of certain useful trace elements. Conversely, drink water that is too hard (has a lot of ions in it) is correlated with kidney stones. It’s unlikely that the ionic content of your drinking water has a huge impact on your health (despite claims found here and there). All of the trace elements (including fluoride) can be found in other food sources. And distilled water’s osmotic pressure isn’t so different from that of plain water, therefore drinking it will not cause the cells in your body to suck up water until they burst and you begin to bleed internally (yes, this theory is out there, for both distilled water and deionized water). Bottom line, yes, you can drink distilled water.

That said, you probably shouldn’t drink the distilled water in most labs, as it is not tested to be free of bacterial contamination (which it can pick up in storage tanks) or volatile organic compounds. The same goes for bottled distilled water that hasn’t been tested to be certain it’s potable.

And while we're on urban myths about water, it's impossible to completely remove all the ions from water. Water is always in equilibrium with hydronium ions (H₃O⁺) and hydroxide (OH^-).

Changing exams

2011-09-02T14:00:00.005-04:00

I just handed out a math assessment in my physical chemistry class, the same one I’ve used for the last several years. I generally don’t re-use exams (though I know colleagues who do), though I do re-use questions. By now I’ve been creating exams for more than a quarter of a century, and I wonder what the drift has been like over that time. How are the questions I ask now different (or not!) from what I asked 25 years ago? Or have the questions remained the same, and just the answers changed?

Fueling my introspection are the selections from the University of London’s 19th century bachelor’s degree exams. (H/T to a tweet from Nature Chemistry and the RSC). The chemistry question is one I could envision asking my students on an exam: “Explain the nature, from a chemical point of view, of the chief operations involved in the production of a photograph.”

The only catch, of course, is that the answer I’m expecting could be quite different than what the examiners in 1892 expected. In 1892, production of a photographic print necessarily involved silver, developers and fixing agents — and a darkroom. In 2011, production of a print could involve silicon and germanium, and a clean room. The theoretical underpinnings are less about pH and solution chemistry and more about semi-conductors and quantum mechanics.

What other reasonable exam questions might I ask, where the answers have changed so dramatically?

(And you have to love the example English question - just how important were werewolves in the 19th century?)

Photo of 39/365 Kodak Vigilant Six-20 Antique Camera, by M.Christian on Flickr.