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.

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.

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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.

Polysemy and Polyphony: Listening to Messiah

Last spring I wrote a piece for Nature Chemistry on polysemy — the phenomenon where words take on quite different meanings in different contexts. The iconic chemistry example might be mole (the quantity versus the animal versus the verb¹), but there's a long list.

So you might think that when I ran into a homograph² on Twitter the other day, I'd be alert to the possibility. My first thought when the conversation between two chemists about the insights they find in Messiah showed up in my feed they were talking about the classic quantum mechanics text by French physicist Albert Messiah.  Actually, not.  Handel's Messiah was the text under discussion.  Polyphony crashes into polysemy.  And evidence I really am a science geek first and foremost.

The text is still in print, though Albert Messiah died in 2013 at aged 92.   I used Messiah's text when I took a year long course in quantum physics as a graduate student (from the physics department, have exhausted the chemistry offerings as an undergrad). We pronounced his name "mess-ee-uh" rather than "mess-eye-uh," making this technically a homograph (though not a capitonym³).  I wondered today how he might have pronounced his name, is it really a homograph, or did my professor simply choose to pronounce it this way to avoid sounding like an evangelical preacher when he assigned reading?  I dove into the interwebs to see if I could uncover any clues.  I discovered Messiah had been part of the French Resistance in World War II (joining at age 19, the age my youngest son is now), worked at the Institute for Advanced Study in Princeton with Niels Bohr and eventually returned to France to teach and write this text.

I also listened to a few minutes of a presentation Messiah gave in 2009 at Le Ecole Polytechnique.  It was oddly moving to hear the voice of someone whose written words I had spent so much time wrestling with almost forty years ago.  And at the end of the questions, I learned how he pronounced his name.

And, on the Sceptical Chymist, Reuben Hudson has a post responding to my column on a different kind of doubling-up in chemical language.

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1.  Yes, mole is a verb, to mole a garden is to remove the moles.
2.  Homographs are words that have the same spelling, but different pronunciation (lead and lead).
3.  Capitonyms are homographs with different capitalization.  DEFT and deft.

Shedding some light on chemistry: Mole Day and the Year of Light

[If you want to participate in some science about science blogs, see the bottom of this post!]

It's October and there is lots of science to celebrate.  Chemists in the US and elsewhere are celebrating Mole Day on Friday (October 23 at 6:02 pm) to honor Avogadro's number (6.02 x 10²³ items are in a mole -- it's the chemist's version of a dozen).  It's also the International Year of Light, and while you might think that light is the purview of physicists, it's an element of chemistry as well.  I suggested in a recent essay that one might want to celebrate the year of light on the 10th October at 3 in the afternoon (3 x 10¹⁰  is the speed of light in cm/sec)

I've written two pieces on the relationship between chemistry and light for the celebration.  The first for Nature Chemistry, The Enlightenment of Chemistry, looks at the two-way relationship between chemistry and light.  Light is  not just an energy source for doing chemistry, but the production of light in various ways has pushed chemistry forward.  The full text is here.

The second, for the UN's Year of Light blog celebrates the October 27th anniversary of Bunsen's and Kirchhoff's publication on the spectroscope and atomic emission spectra — and the role the spectroscope played in not only filling out the periodic table, but in confirming the periodicity of the table.

"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.)" — read the rest here.

Want to participate in some science to celebrate?

Help us do science! I’ve teamed up with researcher Paige Brown Jarreau to create a survey of the Culture of Chemistry's readers. By participating, you’ll be helping me improve the blog and contributing to SCIENCE on blog readership. You will also get science art from Paige's Photography for participating, as well as a chance to win a t-shirt, a $50 Amazon gift card and other perks!   It should only take 10-15 minutes to complete. You can find the survey here: http://bit.ly/mysciblogreaders. 

Changing exams

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?)

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

Bunsen and quantum mechanics

Today's Google doodle honors the 200th birthday of Robert Bunsen, the inventor (or not?)of the eponymous burner. The doodle is great, click on it and it bubbles and whirs.

"It is known that several substances have the property of producing certain bright lines when brought into the flame. A method of qualitative analysis can be based on these lines, whereby the field of chemical reactions is greatly widened and hitherto inaccessible problems are solved. We limit ourselves here to developing the method for alkali and earth-alkali metals and demonstrating its value by some examples.

The lines show up the more distinctly the higher the temperature and the lower the luminescence of the flame itself. The gas burner described by one of us (Bunsen, these Ann. 100, p. 85) has a flame of very high temperature and little luminescence and is, therefore, particularly suitable for experiments on the bright lines that are characteristic for these substances." Opening to Gustav Kirchhoff and Robert Bunsen, Annalen der Physik und der Chemie 110 (1860), 161-189.

Bunsen is not a name typically associated with the development of quantum mechanics, yet I might argue he is one of the key figures. The observation of line spectra and the realization that the lines are characteristic of particular elements is a significant step toward the development of quantum mechanics. It's one of the observations that Bohr was trying to explain in his model of the atom. General chemistry texts boast figures of line spectra to demonstrate the point - I showed several in my lecture last week. This apparatus developed by Kirchoff and Bunsen made possible the routine observation of such lines. I have a beautiful brass example in my office.

This paper goes on to note that sodium, even at very low concentrations produces quite bright lines. It reminds me of the many happy hours I spent playing with my mom's gas stove and making flame tests on anything I could scrounge up (most of which contained sodium). Is this the formative experience that impelled me toward quantum mechanics? Who knows! I do still think of sodium and line spectra every time the pasta boils over and the flame on my stove flares that characteristic sodium yellow-orange.

Happy birthday, Bunsen, I might not have a job without you!

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There is more on Bunsen beyond the burner at The Sceptical Chymist.

Cooling trends

In a previous piece about density and the insulating power of whipped cream, I mentioned that trapped "nothingness" was even better at insulating things than the trapped air in the whipped cream.

A hot object can transfer heat to its surroundings by conduction and convection, in other words by having molecules (or atoms) in the surroundings collide with the surface of the object, pick up some of its energy and move off. Imagine a little bucket brigade, stepping up to the object, grabbing a piece of heat (energy, really) and heading off to dump it elsewhere. Air doesn't conduct heat very well, and trapping it reduces convection (the bulk movement of the air - air currents), hence trapped air is an effective insulator.

If you halt the bucket brigade - prevent any molecules from walking off with a chunk of energy, all the energy stays in the object and voila my cocoa stays hot. So if I could envelop my cocoa in a bubble of nothingsness — a vacuum — I could keep it from cooling via conduction and convection. (Of course, I'd have to put a vacuum tight lid on the thing, lest it instantly boil off — but that's another post…) A convenient way to do this is to use a vacuum flask, where a 'layer' of vacuum is enclosed between two walls. Originally glass walls were used. If you're my age, you might remember dropping your lunch box, and then opening your thermos at lunch to find shard of glass inside. These days stainless steel or plastic walls make lugging your milk to school a less risky proposition.

Chemists still use the glass walled version of a thermos. We call it a Dewar (for James Dewar, who invented the contraption in 1892).

But an object untouched by other molecules can still lose heat by radiation, by emitting infrared photons - light at wavelengths longer than visible light. Reflective wrappings, like metalized Mylar, keep the light - and the warmth - inside.

A friend recently wondered why clear weather was cold weather. The earth radiates some of its heat back out into the universe as infrared radiation. Clouds act as insulating wraps for the earth, the water molecules grabbing the heat before it radiates out into space and re-emitting some of it back toward the planet's surface. (This is the greenhouse effect — it's not up for debate.) On a clear night, no clouds, so less heat is retained. Suddenly the temperatures are well below freezing.

(Other molecules besides water help trap infrared radiation within the atmosphere, including carbon dioxide and methane.)

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Photo is from Wikimedia.

Oprah's take on quantum mechanics - and mine

I was checking my blog stats (read seriously procrastinating folding the laundry) and noticed that one of the search terms that was sending surfers to my other blog was "Oprah's take on quantum mechanics". She has one?

I promptly popped it into Google to see what would come up. I had to know.

I found out. The Law of Attraction. Think and you can change what happens. Proven by quantum mechanics. The Quantum Cleanse. (Don't ask - you don't want to know.)

Somehow the word "quantum" manages to sound simultaneously mysterious and scientific, and so people attach it to things that they want to sound simultaneously mysterious and scientific. Like diets and the power of positive thinking, or even theology.

I named my personal blog "Quantum Theology" as a play on the two fields I'm trained in: quantum mechanics and theology. Recently a friend of almost forty years wondered just exactly what was quantum mechanics - just what do I do for a living? Repair broken quantums?

When I say something is quantized, I don't mean it's mysterious, I mean that only certain values are allowed, and nothing in between. A good everyday example is your shoe size. You are a 5 or a 5 1/2, but never a 5 1/6. Off the rack shoes (are there any other kind these days?) are quantized.

To a physicist or physical chemist, a quantum is a fixed portion of energy. (The word was coined by Max Planck in 1900.) Quantum mechanics considers the interaction of energy and matter on the atomic level. What happens when light hits an atom? Why is it that only certain amounts of energy can be absorbed? How is it that matter can behave as a particle, and as a wave? Evidence that matter could behave like a wave suggested to Erwin Schrodinger that he could write an equation to find a mathematical description of this behavior.

So what is it I actually do? I use quantum mechanics — specifically solving Schrodinger's handy little equation — to predict the structures of molecules and their energy, then use that information to think about what molecules might exist, or how hard it would be for them to react and what products are likely to form. Right now I'm exploring molecules that are uncomfortably twisted - and topologically "interesting" (Moebius strip molecules).

What do flamingos, Cheetos and Quantum Chemistry have in common?

The vibrant colors of flamingos and Cheetos Cheez Whiz™ are both the result of related carotenoid dyes. Carotenoids (named for the vegetable in which they were first found!) are based on a linear conjugated diene skeleton, and provide nature with many colorful accents. Canthaxanthin, for example, is fed to captive flamingos to produce their characteristic pink color (a similar pigment found in brine shrimp does the same favor for wild flamingos). Astaxanthin is responsible for the characterstic color of lobsters. Canaries, whose signature color is a greenish yellow, can be turned red if they are fed paprika during their molt. The new feathers will grow in orange-red.

If you're tired of only changing the color of your hair, you can try for a pumpkin look for fall. The compound that gives this class of vegetable pigments its name - β-carotene - when consumed in large quantities by humans, will turn them orange. [Really, but don't try this at home! It was observed clinically in Britain during WW II when food shortages led some people to include large amounts of carrots in their diets.]

If you thought the bright color of Cheez-Whiz and Cheetos was artificial -- it's not. Bixin or annatto, a natural pigment used for centuries, is the source of that unforgettable orange. Researchers have recently elucidated the biochemical pathway for the synthesis of bixin and are pursuing genetic engineering approaches to its bulk synthesis in tomatoes [Florence Bouvier in Science, 300:2089-2091, June 27, 2003].

What does this all have to do with quantum chemistry? A very simple quantum mechanical model, the particle in a one-dimensional box, can be used to predict the color of conjugated dyes.

UPDATE: Bixin isn't used to color Cheetos, but is used in Cheez Whiz.