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.
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.
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....
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?
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?
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.
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 [Nature481, 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.
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.
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?
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 32o 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 0o 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.
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).
"...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.)
