Field of Science

Five Books: A short reading list for chemistry

Only five books? And the five best books? Last month I did an interview via email with Caspar Henderson (who wrote a marvelous bestiary for the new century: The Book of Barely Imagined Beings) on the best five books I would put on a reading list titled "Chemistry."  It's now up on the site — Five Books.  But the hardest part was not answering the great questions Caspar posed, but figuring out what five books to list. What did I want this list to do? Teach you chemistry? Maybe. Or give you a sense of what I find fascinating and beautiful and compelling about chemistry? Definitely!

I thought about various friends, curious and readers, but who don't have much background in the sciences and math.  What would I pull from my shelves for them to read?  Something that teaches you to decode a bit of the chemistry, a biography - what is the life of a scientist really like.  Something that is compelling, that drags you into a story you can't put down. Something that shows off the beauty of the world at the atomic and molecular level.

Something that teaches you to decode a bit of the chemistry:
Why does asparagus make my wee smell? And 57 other curious food and drink questions by Andy Brunning of Compound Interest. A bold graphical look at the chemistry of what we eat, with lots of quick explanations of weird (but useful) words of science like chromatography. 
What is the life of a scientist really like:
Obsessive Genius: The Inner World of Marie Curie by Barbara Goldsmith.  Of course there had to be Marie Curie. And this unsparing biography of her pulls the curtain away on what it can mean to plunge into research with all your being.
Compelling stories with chemistry at their heart:
The Poisoner's Handbook: Murder and the Birth of Forensic Medicine in Jazz Age New York by Deborah Blum. Some molecules are thugs, some turn witness for the prosecution. Real crimes, real molecules.  (And her new book on the rise of food safety, The Poison Squad, which is in the stack on my desk, is just as good.) 
The beauty of the atomic and molecular world:
H2O: A biography of water by Phillip Ball Chemistry laid out for the layperson with care and delight. Clouds are not what you think!
The Disappearing Spoon: And Other True Tales of Madness, Love, and the History of the World from the Periodic Table of the Elements by Sam Kean. There's a dark side to the periodic table.
Read the whole essay to find out more about what is fascinating about chemistry (at least to me), what I do as a chemist, and of course, about these five books. Want more book recommendations about chemistry? Want to know what the runners up were? Leave me a note in the comments!

Chemistry not your thing? Go read Caspar's bestiary about the wildly improbable creatures that inhabit the very real world, from sea butterflies to yetis (or at least yeti crabs), it's a wide ranging exploration of the corners of the biological world. To quote a reviewer: "There is something lovely about a book that takes on so many disciplines and tackles them with confidence." There is indeed.

[Cross posted from Quantum Theology]

Trying to explain earthing with atoms

My ungrounded feet in rubber boots.
This week the Washington Post has an article headlined "Could walking barefoot on grass improve your health? Some science suggests it can."  The link itself is subtitled: The science behind grounding.

The article gets a lot of things right about atoms (they make up everthing!), but it confuses "free-radicals" with positive ions. (Free radicals don't have to be charged.) Then it tries to explain why negative ions can help. And while it is true that a positive ion and a negative ion can react in some circumstance to produce a neutral compound (think of hydroxide and hydrogen ions reacting to make water in an acid base reaction), random negative ions won't necessarily disarm a free radical.  You need an antioxidant for that, a molecule that can participate in a reaction that can soak up extra electrons.  You still need to eat your vegetable and wear sunscreen.

Negative ions and positive ions co-exist quite nicely in your body. You need those positively charged potassium ions, in fact, to keep your heart beating rhythmically. So on its face, the "science behind grounding" given in the article is bunk. If all those negative ions in the ground started neutralizing all the positive ions in our bodies, we'd be dead.

While I get this is a not a science news piece, but a perspective piece (a "[d]iscussion of news topics with a point of view, including narratives by individuals regarding their own experiences"), I wish someone at the Post had fact-checked the science.  Yes, it feels nice to walk barefoot on the grass, or to be outside.  I'm pretty certain the negative ions aren't the reason why.

Hunting up the ghosts of elements

This post originally appeared at the UNESCO International Year of Light's blog, in October 2015.  The site is no longer available.

Interior of an antique spectroscope.
If you’ve seen the flash of yellow-orange flames when a pot boils over on a gas stove, you’ve gotten a glimpse of the ghost of an atom.  The color is part of the atom’s spectrum.

In the late 17th century, Isaac Newton used the Latin word for ghost, spectrum, to describe the bands of colors he saw when light shone through a prism. One hundred and fifty years later, Joseph von Fraunhofer noticed he could see bright lines instead of the bands of colors when looking at certain flames through a prism.  He went on to develop an instrument to measure these spectral lines, called a spectroscope, and used it to catalog the lines seen in the sun’s light and in the light from other stars.

It would take almost another fifty years to figure out that Fraunhofer’s lines were the ghosts of chemical elements, when Gustav Kirchhoff and Robert Bunsen (the inventor of the ubiquitous Bunsen burner) teamed up to create a spectroscope that used Bunsen’s new hotter, gas burner to ignite the samples.  They noted that that each element produced a characteristic set of lines when burned, a spectral fingerprint, that could be used to identify it.

In October of 1860, Kirchhoff and Bunsen announced they had used their spectroscope to discover a new chemical element, which they named cesium, for the blue color of its principal line. Chemists quickly began to use Bunsen’s spectroscope to find new elements.  A few months later Kirchhoff and Bunsen found two bright ruby red lines in an extract of a silicate mineral lepidolite, the spectral traces of another new element, rubidium.

Thallium’s ghostly green emanations were first observed by William Crookes, indium, ironically named for its violet lines by its color blind discoverer Ferdinand Reich.  Paul-Émile Lecoq de Boisbaudran spectroscopically identified element 66 in a sample painstakingly extracted from his marble hearth, and instead of naming it for the colors of the lines, called it dysprosium, from the Greek for “hard to get” — because it was.

Hunting for new elements spectroscopically meant you didn’t actually need to have any of it in your lab or even on your planet, as long as you could observe the light from a burning sample.  In 1868 several chemists and astronomers independently observed a faint line in the spectrum of the sun, and assigned it to a new element, helium, which as far as they knew did not exist on earth.  It would take nearly 30 years for two Swedish chemists to confirm that it was present on earth — by matching the spectrum with that of a gas found in a uranium ore.  (The helium to be found on earth comes from radioactive decay.)

Spectroscopy certainly helped chemists fill out the periodic table, adding more than a dozen new elements to the collection.  But it also played a significant role in confirming predictive power of periodicity. When Dmitri Mendeleev proposed his version of the periodic table, he left blanks for yet-to-be-discovered elements, underneath elements which should have similar properties.  In 1875, Lecoq, the same man who had so patiently extracted dysprosium from his fireplace, sifted through 4 metric tons of  zincblende to show that it contained traces of a new element which fit neatly into the space Mendeleev had reserved for it underneath aluminum.  Lecoq named the element gallium, in honor of his home, France, and perhaps playing off his own name, as the Latin for le coq, the rooster, is gallus.  It was a powerful demonstration of Mendeleev’s theory.

These ghostly lines produced by elements helped fuel yet another critical discovery that would have far reaching consequences for chemists’ understanding of the periodic table:  quantum mechanics.  Niels Bohr’s quantum mechanical model of the atom opened the door to explaining chemical elements line spectra. Though more accurate and sophisticated quantum mechanical models of the atom now exist, Bohr’s model showed the relationship between the lines and an atom’s electron by insisting that the electrons’ energies were quantized, that is, they could only have certain energies.

So why do atoms have ghosts?  When an atom is heated to high temperatures, as in a flame, the energy it absorbs excites its electrons.  You can think of the electrons in an atom as being on an energy ladder.  They can only have energies that match the rungs of the ladder, and each type of atom has a unique arrangement of the rungs.  When the atom absorbs energy, its electrons move to higher rungs.  Excited electrons are unstable. They quickly return to their original arrangement, giving off some their excess energy in the form of light as they do.  The color, the wavelength) of the light emitted depends on the difference in energy between the rungs.  The colors of light emitted are the ghosts of the energy rungs.  Since each element has a unique pattern of rungs, it will have a unique spectrum of emitted light and so revealing their presence to the sharp eyes of spectroscopists.

Chemists still use the light emitted and absorbed by atoms and molecules to identify their presence.  We hunt for the structure of the universe in its ghosts.

More Information

If you want a way to see the ghosts of atoms, try this DIY folding spectroscope you can attach to your phone. Use it to check out the light from a neon sign or from a street light!

For a wonderful description of the elements, including stories of how they were first discovered, read John Emsley’s Nature’s Building Blocks.


Sikhote-Alin meteorite from Vatican Observatory's
On July 13th, the Earth Science Women’s Network is hosting Science-A-Thon, in which participating scientists are chronicling a day in their life on Twitter and Instagram (follow #dayofscience and #scienceathon).

Join me for the day!  I'll be posting a photo every hour on a day when I'll be working from the Vatican Observatory in Albano Laziale, outside of Rome.  Starting with my early morning stop at the espresso bar through a full day of science behind the walls of Vatican City.   There might even be aliens from other worlds (of the inorganic variety). The observatory might seem focused on anything-but-earth science, but the meteorites that the earth sweeps up as she moves through the heavens are clues not only to the otherworldly, but to our own planet's history.

Participants are listed by country — so far I'm the only one under "Vatican City"!

This is a first-ever fund raiser for the Earth Science Women’s Network, so if you are inclined to support them, you can donate here.

Hidden figures: 2.303, slide rules and classrooms mired in the last century

A five -place table of logarithms from my dad's CRC Handbook of 
Mathematics (why is that set of values circled?) and a circa 
1958 Hemmi 257 slide rule designed for chemical calculations.  

 Wonder why random values of 2.303 are "hidden" in formulae? To make them easier to use with a slide rule.

A slide rule?  The last slide rule slid out the door of Keuffel & Esser in 1975 (they sent their engraving equipment to the Smithsonian).  You can still find them, used and even new - still packaged up to sell to engineers and scientists.  The Oughtred Society has a online museum, as well.

We still have my mother-in-law's K&E, in it's leather case with her name impressed into it.  Family history says she bought it with the money she earned tutoring Jackie Robinson in chemistry at UCLA.

I have an essay out in this month's Nature Chemistry, "It figures", about how the computational tools we use shapes what we teach and not necessarily in good ways. Given that slide rules were obsolete by the time many of my student's parents were born, why does their use still linger in general chemistry book?  (The 2.303's in texts are lowly going away. I checked texts running back about a decade.)

More critically to my mind why, several decades after  digital computing tools became ubiquitous on college campuses do many physical chemistry texts eschew any discussion of numerical techniques for solving the rate equations for a chemical reaction?  I suspect the chasm between the computational tools used in the field and those used in the classroom is a result of apathy. We teach what we learned as we learned it.  As I note in the article, I don't think it is defensible on intellectual grounds.

Don't know how to use a slide rule?  It's fun, it's geeky. No need to buy one to play, check out this simulator and the instructions at Nature Chemistry!

You can read the article here:

1.  2.303 is the natural log of 10. To change the base of logs recognize that
x = blogbx
ln(x) = ln(10log10x)
ln(x) = log10x ln(10)
ln(x) =(log10x)(2.303)
ln(x) = 2.303(log10x)

Implications of Charles law in a biological matrix: farts

See note 3 for source.
Maggie Koerth-Baker has a great piece up at the 538 blog: "How Big Is A Fart? Somewhere Between A Bottle Of Nail Polish And A Can Of Soda."  It's well researched, digging into the biomedical literature with verve.  And it's great that she gives the answer a context, it's easier to visualize a bottle of nail polish than a 17 ml blob for most people, me included.

I'm not at all surprised at what you can find in the primary literature (I tracked down papers on exploding people and deuterated dogs1 for my introductory chemistry class last spring). The piece is the first in a series Science Question From A Toddler, though I suspect that people somewhat past the target age group (5 and under) would be interested in the answer to this question, too.

In a footnote Koerth-Baker suggests that farts in the body would be smaller because the gas would be compressed inside the body.  But the pressure inside the human colon is the same as atmospheric pressure.  Farts and burps keep it that way. What's different is the temperature, higher inside the body by about 30oF (17oC).  Gases expand at higher temperatures. You can use Charles' law to figure out by how much the volume changes with changes in temperature:  V2=(T2/T1)V2

The researchers measured the volume of the farts at room temperature (I read the paper!), so the volume of a fart should be slightly larger in the body than the reported numbers by a factor of (310/293) or about 6% larger.  So how big is a fart?  Just before exit, it's about the size of a 14 ounce ketchup bottle for the largest one in the 1997 study.

The details of the experiments are fascinating.  The technique for quantitatively2 capturing flatus in the bathtub is elegant, and while a significant improvement over the method used for the studies in the 1860s3 you have to wonder how they got volunteers for either experiment.  And speaking of volunteers, the assessment of the "flatus perception threshold" was done by delivering 100 ml of an odorant mixture "from a large 250ml syringe in about 1s, 1 meter beneath the nose of the panel members, mimicking a flatus emission."

And just in case you don't think this is serious stuff: "The common tendency to treat rectal gas as a humorous topic has obscured appreciation of the complex physiology that underlies the formation of this gas." Suarez et al. American Journal of Physiology  272, G1028-G1033 (1997).

1.  The physiological effects of drinking heavy water, D2O, on dogs.  If you've ever wondered what would happen if your poured that little vial of D2O into your coffee, the answer is not much.  It's not great for the dogs as a steady diet, but a sip or two won't hurt.
2.  The fancy chemistry term for "we got all of it!"
3.  See the figure, from Tangerman, "Measurement and biological significance of the volatile sulfur compounds hydrogen sulfide, methanethiol and dimethyl sulfide in various biological matrices" Journal of Chromatography B, 877,  3366-3377 (2009).

Chemists: Strangers to fiction

That Mars habitat?
"The basement corridor is dim, I can hear pumps chugging, hoods noisily venting, and the solid-state physicist down the hall swearing. 'Welcome to Mars!' says the cheery sign outside my colleague’s door. Perhaps it is the pile of grading on my desk or the endless round of meetings on my calendar that is fuelling my escapist fantasy, but every time I pass Selby’s office, I imagine the door is a portal and if I were to walk through, I’d  find myself in a habitat on Mars, its pumps working hard to compress the thin atmosphere." from "Strangers to Fiction" in Nature Chemistry8, 636-637 (2016).

I've been a sci-fi fan for going on five decades, imagining myself in labs on Mars, mining comets, and exploring strange new worlds. I don't read it for the chemistry, which is a good thing, because there isn't much fiction in which chemistry plays a key role.

My latest Nature Chemistry Thesis column looks at chemistry and fiction, suggesting that there are good reasons to both read SF, particularly for young chemists, and for chemists to encourage the writing of chemistry-inflected science fiction.  And if you have the talent for it (which I do not!) perhaps even give the writing of it a fly.

You can read the whole thing here.  My list of fictional chemistry is here.

A matter of degrees: when low temperatures were hot

Diagram of a thermometer similar to
the one describe by Leurechon, c. 1638.
Note that  hotter temperatures have 
smaller magnitudes degrees associated with 
them. Image from Wellcome collection, 
used under CC license.
We say the mercury is rising to mean it's getting hot out, despite the fact that most home thermometers have no mercury in them anymore.  Regardless of the liquid they contain, the level rises with increasing temperature in the iconic liquid thermometer.  But this was not always the case.

The word thermometer was first coined (in French) in a book of mathematical recreations written in 1626 by Jean Leurechon, SJ (writing as Hendrik van Etten).  In his description he notes the thermometer you can construct from a glass tube and small container of water (or other non-viscous liquid) can be used to quantify temperature by placing marks on the glass, associating each with some fraction of the classical four (or eight) degrees of hotness.  Such thermometers, he suggests, can be used to adjust the temperature of a room or a furnace, to record (and predict) the weather and to measure fevers in the ill.

But Leurechon's thermometer (and similar designs) were constructed such that as the temperature increased, the water level in the tube fell.  Increases in temperature caused the air trapped in the ball at the top of the tube to increase in volume, pushing the liquid down in the tube.  (These are air thermometers, in contrast to the familiar liquid thermometers in widespread use today.) A reading of 9 degrees on the thermometer shown in the sketch accompanying Leurechon's thermometer problem was colder than that of 2 degrees (see also the one in Robert Fludd's diagram, in the figure.)

A century later, Anders Celsius constructed a temperature scale based on water's phase changes which ran in the same direction.  Water on Celsius' scale boiled at 0 degrees and froze at 100 degrees. This reverse run didn't last long, two years later Carl Linnaeus (of taxonomic fame) used the scale to describe conditions in a greenhouse, but flipped it to the form in which we know it today, where 100 is the boiling point of water.

It is tempting to think that Celsius' scale ran in the direction it did because it mimicked the earliest marked thermometers. But Fahrenheit's scale, which preceded Celsius' by two decades, runs in the modern direction, things get hotter in the positive direction. This also parallels the classic notions of degrees of heat in play during the medieval period. There were four (or eight or six, depending on the source) degrees of heat, the first being more or less physiological temperature, the fourth being a blazing hot furnace.

The word degree has its roots in the Latin degradum, a down step.  This matches Leurechon and Celsius' use - 9 degrees is eight steps lower (colder) than 1 degree.

Chemical fiction

Topi Barr's Antithiotimoline is in this vintage Analog
Seven years ago, Andy Mitchinson, an editor at Nature, wrote at The Sceptical Chymist (Episodes II and III) about the dearth of science fiction that involved the science of chemistry in a substantive way.  Why isn't there more of it?

He pointed to a list put together by Connie Willis, an award winning SF author, and an article by Philip Ball in Chemistry World.

I'm working on a column for Nature Chemistry about the ways in which chemistry and science fiction play off each other.  Is science fiction more than escapist entertainment?  Should chemists care that there's not more chemistry inflected fiction out there?  Should we deliberately expose students to science fiction? Should we encourage them to write it?

To go alone with the piece, I'm trying to create a periodic table of chemical fiction (not including articles called out by Retraction Watch).  Are there pieces on my list you particularly love?  Something I'm missing?  I'd love to hear in the comments!

For a full set of periodic science fiction short stories, I encourage you to browse Michael Swanwick's Periodic Table of Science Fiction.  What really happened to the Hindenburg?

Author Work
As Asimov, Isaac Whiff of Death, The Endochronic Properties of Resublimated Thiotimoline, Thiotimoline to the Stars, Pate de Fois Gras
Pb Ball, Philip The Sun and Moon Corrupted
Ba Barr, Topi “Antithiotimoline”
B Bujold, Lois McMaster Vorkosigan series
Ac Christie, Agatha "The Blue Geranium” in The Thirteen Problems
Cl Clements, Hal Phases in Chaos
Co Conan Doyle, Arthur Holmes
Md Dewar, Michael “Temporal Chirality:  The Burgenstock Communication”
F Foster Wallace, David Infinite Jest
Ag Goodman, Allegra Intuition
He Heinlein, Robert Glory Road, Have Spacesuit will Travel
Hf Hoffman, Roald Oxygen
Li King, Laurie Russell & Holmes series
U Le Guin, Ursula “Schrödinger’s Cat”
Sn Lem, Stanislaw “Uranium Earpieces” in Mortal Engines
P Levi, Primo The Monkey’s Wrench
Am McCaffrey, Anne Pern series
H Piper, H Beam Omnilingual
Kr Robinson, Kim Stanley Mars series
O Sachs, Oliver Uncle Tungsten
Dy Sayer, Dorothy The Documents in the Case
Sm Smith, Edward Elmer “Doc”  “Tedric,” “Lord Tedric" in The Best of E. E. “Doc” Smith
Ne Stephenson, Neal Anathem
Br Stoker, Bram Dracula
Fr Vance, Jack “Potters of Firsk”
K Vonnegut, Kurt Cat’s Cradle
V Vourvoulias, Sabrina INK
Hg Well, H.G. “The Diamond Maker” in The Stolen Bacillus and Other Incidents
C Willis, Connie The Sidon in the Mirror

A day in pchem lecture: NMR, lululemon yoga pants and tattoos

By lululemon athletica
(Flickr: Yoga Journal Conference)
 [CC BY 2.0], via Wikimedia Commons
It's the end of term, two more 90 minute lectures left in my introductory quantum chemistry and spectroscopy course.  We've done the basics of wave functions and expectation values, we've looked at linear variation theory and written code to do Hückel MO calculations, we've covered rotational and vibrational and rotational-vibrational spectroscopy.  So what to do with these last few days?  The quantum mechanics of NMR.

I kicked off today's lecture by looking at magnetic field strengths, what's the earth's magnetic field (5 μT) or of a refrigerator magnet (5 mT), compared to the superconducting magnets used in NMR, which are on the order of 10T. (1T is one tesla.)  This led to a quick review of the risks in MRI, which aren't about the energy of the radiation used (which is billions of times lower than X-rays), but more about the interactions of the high magnetic fields, the radiofrequencies and metals.

A hand shot up and student who is an EMT describes a patient whose tattoo started burning during an MRI.  I pointed out this is a known phenomenon, and while most inks don't pose an issue, it should discourage you from DIY tattooing.  Then a student asked, "Is it true you can't wear lululemon pants when you have an MRI?"

I admitted this was out of my zone, but promised to follow up.

I can now report that yes, wearing lululemon pants — or any clothing with metallic microfibers, such as those great antimicrobial t-shirts — in an MRI can lead to serious burns, particularly in patients that have been sedated or are otherwise unconscious and unable to signal their discomfort.  Even non-ferromagnetic materials presents problems in the MRI as eddy currents can develop around them, creating little induction heaters.  Loops of all sorts, even skin to skin contact between a patient's own body parts can lead to heating and subsequent burns.  And tattoos with large loops in them?  They can heat as well.

Other things I learned this afternoon.  You can levitate a frog with a 16T field (thank you Wikipedia), and neutron stars have magnetic fields on the megaTesla scale.