Thursday, June 02, 2011

Of a hundred all is made

kw: book reviews, nonfiction, chemistry, elements

If you were to touch molten aluminum, you'd lose some skin, at the very least. But there is an aluminum-like metal, about twice its density, but less dense than pewter or tin, that you can hold in your hand, where it will gradually melt into a shiny puddle. It isn't mercury, but gallium, and it melts at just under 30°C, or 85.6°F. This makes it amenable to a chemist's practical joke. Cast an ounce or two of it into a teaspoon, and bring it out with a serving of tea on a cool day. When your friend attempts to stir in some sugar or cream, the spoon will vanish into the tea! Do stand by, and prevent your friend from drinking the tea at that point; gallium is not strongly toxic, but ingestion is a rather bad idea.

Thus the title of the book, 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. While early sections of the book introduce the elements and the periodic table and its history in a pretty regularized way, most of the book is a cross between rummaging through the intellectual attic and a topical, eclectic story-fest that manages to introduce us to every element by the end. That includes the newest members, Roentgenium (Rg; 111) and Copernicium (Cn; 112), and some others that, though long known, have never been seen, such as Francium (Fr; 87). In the latter case, a visible amount of the element's most stable isotope, 87Fr223 with a half life of but 22 minutes, would probably kill you within the first few seconds of viewing. It is 38 million times as radioactive as Radium.

Fortunately, most elements are more prosaic, frequently with lively or even colorful chemistry, but with little tendency to strike you dead upon sight. The regularities in the elements' chemical nature underpin the structure of the periodic table. Ever wonder why it is called "Periodic"? Even many who have taken a chemistry course or two may not be clear on this. Here is a blank table, such as those used for dreaded quizzes, because the clues are in the table's shape:

The term "periodic" stems from the regular repetition of chemical properties as you advance through the elements. There are eighteen columns to the main table, plus fourteen in the added double row below. If those were inserted where the little gap is near the bottom, the table would have 32 columns. But it is the way in which each pair of rows lengthens that reveals something that is only explained by quantum electrodynamics—and by "explained", I mean the "what" of it, not the "how" or "why", which remain wholly mysterious.

In all the natural elements, and in all fabricated elements up to number 120, electrons are nested in shells and sub-categorized into four kinds of sub-shells called orbitals. The rightmost column in the periodic table contains those elements which have completely filled shells, while the first column contains those with completed shells, plus one electron. This lone electron gives these elements a strong similarity; they can rip the oxygen right out of water, and as you go down the column, they get more violent about it. But the elements with filled electron shells will have none of it. Only a few of them, under extreme duress, have been persuaded to form fragile compounds. Ordinarily they are aloof from chemistry, and are called the noble gases as a result.

The odd shape of the table is due to the way orbitals enter the picture; each has a letter with historical import, and a capacity for a finite number of electrons. They are designated s (2 electrons), p (6 electrons), d (10 electrons) and f (14 electrons). These letters are based on an old spectroscopists' convention, in which spectral lines caused by electron jumps that terminate in the various orbitals are denoted "sharp", "principal", "diffuse", and "fundamental". Highly-excited gaseous atoms can have even higher orbitals, which are just called g, h and so forth. But no known elements fill such orbitals in a stable state (yet).

From the second row onward, the s and p orbitals together form the "valence shell" that controls the principal chemical nature of the element, though the presence of other orbitals in deeper shells modifies it. Thus, row 1 contains just two elements, hydrogen (H; 1) and helium (He; 2), with only s electrons. The next eight elements fill first the s orbital then the p orbitals (there are three, each holding just two electrons, to total the six mentioned above). The following eight elements repeat the pattern.

After elements 19 and 20, however, the pattern changes. For reasons unknown, but mathematically explained by quantum mechanical equations, the five d orbitals "want" to be filled first, before the p orbitals. This pattern occupies the fourth and fifth rows. Then in the sixth row, we find that the f orbitals (seven of them) get filled first, then the d orbitals, then finally the p orbitals. The seventh row is the same. All those elements that occupy the middle ten columns characterized by d orbital filling are called "transition metals", because their chemistry makes a transition from oxidizing at one end to reducing at the other. The lower section, two part-rows of fourteen columns, are called first "rare earths" above then "rare earth-like" below, and have such similar chemistry that obtaining any of them as the pure element is nearly impossible (a mass spectrometer is the most effective method).

The table shown has room for 118 elements. To date, 112 have been unambiguously found or produced. Scientists, and the author, have speculated about finding quasi-stable elements with atomic numbers (proton number, in other words) greater than 118. What happens in the eighth row, assuming enough of an element could be accumulated long enough to scope out its chemistry? Here is a point that I've never seen discussed. Element 119 will have a single s electron in its eighth shell, and 120 will have two. Then what? Here, quantum mechanics predicts the intervention of a new orbital, which we can call g, still following the spectroscopists. Assuming the nine possible g orbitals all fill first, before the f orbitals get any electrons, there will be the need for another "moat" of length 18, below what is already shown. Were the table to be assembled with all elements in order, it would at that point be fifty columns wide.

How far can it go? As the author explains, the velocity function of the electron is described by a "fine structure constant", which has a value just a whisker greater than 1/137. Once the number of electrons surrounding an atom reach this number, the electrons in the innermost s orbital will need to exceed the speed of light! Maybe so, maybe not. I suspect people now living will be around when elements with atomic numbers approaching 137 are synthesized. Whether such elements are possible or not, the attempt will cause rewriting of our physics books.

That is quite a digression from a few opening chapters, plus a little speculation on my part. Most of the book is the stories of how all this was put together, by Mendeleev and others, and even more how various elements have played a role in world events, or declined to do so. Consider Ruthenium (Ru; 44). A white metal very similar to its neighbor silver, (Ag, 47), though harder, it was considered useless until the Parker Pen Company decided to produce a really luxury fountain pen, the Parker 51, with a gold nib. Gold is soft, so the very tip was an alloy of 96% Ru and 4% Iridium (Ir; 77), which hardened it further. So far as I know, its only other uses are as a minor alloying element, and it is seldom considered essential. While the Parker 51 didn't change the world, it was "the" accessory of every powerful person from 1944 until fountain pens were superseded by ball point pens a generation later.

The elements that changed the world the most, uranium (U; 92) and plutonium (Pu; 94) are known primarily for their damaging potential, as the stuff of bombs. Yet they (nearly all U) quietly drive a couple of hundred power generation plants around the world, keeping the air conditioners of millions of people running.

The topical sections of the book cover the gamut of human history and endeavors. While the "heavy metals" such as lead are in the news these days (can you believe we used to burn a lead compound with our gasoline?!?), there is a "poisoner's corner" from arsenic (As; 33) down and to the right, that contains the really bad actors. Arsenic sits below nitrogen (N; 7) and phosphorus (P; 15), such that its chemistry allows it to substitute for phosphorus when it gets into us, gumming up the energy-producing machinery that keeps us alive. This is because the As ion is larger than the P ion, so things get stuck. Interestingly, below arsenic we find antimony (Sb; 51), which is reasonably safe to handle, and used to be used in type metal. It and gallium (Ga; 31) are the only two metals that freeze into a solid less dense than the liquid. The Sb atom is too big to substitute for P, so it doesn't poison us the way As does. Next down we find bismuth (Bi; 83), so innocuous it is safe to ingest (there is a lot of it in Pepto-Bismol!). By the way, the most stable isotope of Bismuth, 83Bi209, was recently found to be very slightly radioactive, but with such a very long half life—between ten and twenty billion billion years—that a gram of it will experience only about one decay per day.

The next column over is just the reverse. Oxygen (O; 8) and sulfur (S; 16), which are required for life, sit just above selenium (Se; 34). A tiny bit of selenium is an essential nutrient, at least for most mammals. But a lot of it damages the brain, leading to the designation of Se-bearing plants as "locoweed". But cattle love it, preferring the great high it gives, even as it kills off the brain. Ingested selenium also smells bad, so that it is sometimes called "stinkelenium". Next below Se we find tellurium (Te; 52). Not too toxic, but just a little of it will give you such stinky breath and bad BO that it will be years before your social life recovers. Then, next down, there's polonium (Po; 84). Always radioactive, with a very short half life and thus very strong activity, it is easier to make this stuff these days than to extract it from uranium ore. It was named in honor of Poland by Marie Curie, but the world simply yawned. Poland has usually not even been there, though these days it looks like it'll endure for a while. But polonium became famous when it was used to poison a former Soviet spy. A few micrograms sprinkled on some sushi were all it took to do him in.

I guess I simply have to wrap up. This is one book I could go on reading, if only it were longer (It is only 376 pages, endnotes included). I have a really geeky desire to repeat all the stories, but it is for the author to shine, not me. I had to slow myself down so as not to miss things. It is a real page-turner. The introductory chemistry that starts the book eases any reader into the meat of the book, so I am sure it is accessible to an audience much wider than chemistry addicts like myself. More than any popular book about chemistry, it shows how chemistry is geometrical in nature (remember that bit about "big" arsenic messing up the phosphorus works). This is even more true of organic chemistry, which is only slightly touched on in this book (hint to Sam Kean: You could do book after book about interesting families of organic compounds!). Now I regretfully lay it aside to take up the next book on my nightstand.

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