Chemical & Chemical Engineering News (80th Anniversary Issue), Vol. 81, No. 36, 2003, Sept. Edited by X. Lu Introduction

RUDY M. BAUM , C&EN WASHINGTON

It is among the most seductive of elements. Quicksilver. The only metal that is liquid at room temperature. Indeed, mercury's chemical symbol, Hg, derives from the Latin name for the element, hydrargyrum--literally, liquid silver.

It is cool, dense, shiny, and slippery. It is irresistible.

Mercury is one of the handful of elements that have been known since antiquity. Alchemists were beguiled by it, convinced that it held the key to transmuting base metals into gold.

And like so many seductive things, mercury is dangerous. The ancients also knew that mercury was highly toxic. Compounds of mercury have been used in medicine and as disinfectants for centuries. Calomel--Hg₂Cl₂--was long the only treatment for syphilis, although the treatment was nearly as awful as the disease.

DROPPING IN Droplets of mercury on a green surface. The surface tension of mercury is so high that the smaller droplets form almost perfect spheres.

Even almost 40 years ago, when I discovered a small plastic bottle filled with mercury in the storage locker in my seventh-grade science classroom, I knew that mercury was poisonous. That it should not be handled. That it should not be trifled with.

That knowledge didn't stop me and a few of my scientifically oriented friends, of course. The attraction of quicksilver was far too great.

"You can put a penny in it and it will coat it like that!" one said. "Can't get it off, either."

So in went the penny, and sure enough, out it came shining like a newly minted dime. But more than that, the penny had that unmistakable slipperiness you associate with mercury. And you couldn't rub it off. What had happened? Did the mercury react somehow with the copper of the penny? We didn't know where to turn, and we certainly weren't going to ask our science teacher, because we knew we weren't supposed to be performing this particular experiment.

We wanted to do something with the mercury. We thought we might be able to make a barometer, but the problems associated with filling a glass tube with mercury and somehow sealing one end of it were beyond us technically. We proved it was conductive with a battery, wire, and light bulb. We ran it through plastic tubing, fascinated by the lightning-quick flow of the metal. Although we were reasonably careful, we probably contaminated the classroom with more mercury than some government agency deems acceptable.

I mentioned the mercury to my father over dinner, and he recalled his service in Europe in World War II. He and his brother Walt were both in the infantry, and they met twice while they were crossing Europe and serving in the occupation forces after the war ended. On one of those occasions, Dad met his brother at a small railroad freight yard in Germany. Walt was among a contingent of Army engineers accompanying a freight train transporting a shipment of mercury. Dad recalled Walt opening up the doors of one of the boxcars to show him a single layer of flasks of mercury spaced evenly across the floor of the railcar. They were so heavy they couldn't be packed any more densely. Dad didn't remember where the train was coming from or where it was going, but he did remember that it carried that load of mercury.

I can't say that my early exposure to mercury affected my much later decision to become a chemist. Those seventh-grade experiences were soon a distant memory. Mercury isn't a particularly interesting element chemically, and it's so poisonous that it doesn't play a role in chemistry sets, even those of my boyhood. (I went to college expecting to major in biology; my shift to chemistry had much more to do with the Duke University chemistry faculty than my early experiences with a chemistry set.)

But I've always retained a mild fascination with mercury. I still find it unfortunate that mercury has been banished from thermometers for safety reasons. A mercury thermometer has more gravitas than one filled with a red liquid. And a frisson of fear, as well: I remember clearly C&EN reporting in 1997 the death of Dartmouth University chemistry professor Karen E. Wetterhahn after she was exposed to methyl mercury while preparing a nuclear magnetic resonance standard.

When C&EN Editor-in-chief Madeleine Jacobs mentioned more than a year ago that we should celebrate C&EN's 80th anniversary in 2003 and asked whether I had any ideas for a special issue of the magazine, my eyes went to the periodic table hanging on my office wall. What's element 80? Why, it's my old friend mercury. Other anniversaries are designated silver, gold, and diamond, I thought, so why not call this our "mercurial" anniversary? It was a silly idea, but it started the discussions that led to this special issue of C&EN. We hope you enjoy reading it.

Rudy M. Baum is C&EN's deputy editor-in-chief. He has covered a variety of topics in chemistry, the interface of chemistry and biology, and science and society for the magazine for more than 20 years.

MERCURY AT A GLANCE

Name: Named after the planet Mercury, which was named for the Roman god of eloquence, skills, and commerce. The symbol comes from the Latin for liquid silver, hydrargyrum.

Atomic mass: 200.59.

History: Known since ancient times.

Occurrence: Rare in Earth's crust. It is primarily found in cinnabar ore.

Appearance: Silvery white, liquid metal.

Behavior: The only metal that is liquid at room temperature. It alloys easily with most metals and is very volatile. The element vapors are toxic, as are all mercury compounds. It is a cumulative poison that affects the central nervous system and the mouth, gums, and teeth.

Uses: Used in thermometers, barometers, diffusion pumps, and other instruments. It is also used for making batteries, switches and other electrical apparatus, some pesticides, and antifouling paint. Mercury is the basis of dental amalgams and preparations. Gaseous mercury is used in mercury-vapor lamps and advertising signs.

BORON

E. J. COREY,HARVARD UNIVERSITY

Nature and evolution appear to have bypassed the element boron in forming the living world, even though boron is next to carbon in the periodic table. Boron escapes mention in textbooks on biochemistry and at best can claim trace nutrient status. Yet it is an element of great versatility and individuality. Long in the domain of inorganic chemistry, boron has increasingly acquired an organic face. That is why I succumbed to the invitation of C&EN to write this essay, even though my favorite molecules are carbogens, members of the family of carbon compounds (see "The Logic of Chemical Synthesis," John Wiley & Sons, 1989).

There are a few things about boron that one does not forget. Boron nitride is as hard as diamond and similar in structure. Boric acid is cheap and useful. When dissolved in alcohol, it forms ethyl borate, which burns with a beautiful green flame (light emission from an electronically excited state of BO₂?). Boron-oxygen bonds are very strong. Compounds of boron with B–H and/or B–C bonds are highly reactive and easily oxidized. Whereas hydrocarbons such as C₂H₆ and C₄H₁₀ are stable at 25 ºC in air, the boron equivalents B₂H₆ and B₄H₁₀ burn spontaneously, an obvious reason for nature to avoid such boron compounds in living systems.

The understanding of the chemistry of boron that we now possess did not come easily. Preparing boron equivalents of the hydrocarbons was technically difficult, and their structures turned out to be surprisingly complex compared to their carbon cousins. The pioneering work of Alfred Stock, Hermann I. Schlesinger, E. Wiberg, and Anton B. Burg, together with structural studies by Hugh C. Longuet-Higgins and William N. Lipscomb, laid the foundations of modern boron hydride chemistry, which is dominated by the fact that boron has more bonding orbitals (4) than electrons (3). Schlesinger's simplified syntheses of NaBH₄ and B₂H₆ led to extensive application of these compounds as selective reagents for organic synthesis, especially in the work of Herbert C. Brown; many analogous boron compounds are now indispensable general reagents.

The boron halides, especially BF₃, are widely used as acidic catalysts in synthesis. The available p orbital of BF₃ (electron deficient and isoelectronic with CF₃⁺) is responsible for the strong Lewis acidity and the massive industrial use of this reagent. Electron deficiency in BX₃, BH₃, and higher boron compounds such as B₁₀H₁₄ has enormous chemical implications, including the existence of large numbers of polyhedral cluster compounds, for instance, polyhedral boranes, carboranes, metalloboranes, boron nitrides, and heteroboranes. The understanding of the complex three-dimensional architecture of these compounds represents a triumph of modern molecular orbital theory (Lipscomb, Kenneth Wade). Electron delocalization can even be seen in the dimer of BH₃, which is held together by three-center two-electron B–H–B bonds, leading to a cyclic structure (pictured), in stark contrast to the acyclic structure of ethane, H₃C–CH₃.

What intrigues me most about boron are the possibilities for catalytic asymmetric (enantioselective) synthesis when it is embedded in a suitable chiral molecular environment. For instance, we found that a simple (S)-proline-derived chiral B–N compound (1, pictured) combines at nitrogen with BH₃ (for example, from Me₂S–BH₃). This complex can then bind to and reduce an achiral ketone, as a second substrate, to form enantioselectively a chiral secondary alcohol. I like to think of 1 as a molecular robot that picks up one achiral reactant, then a second, and snaps the two together in one preferred three-dimensional arrangement to form a chiral product, which the robot then releases. In fact, 1 is even more versatile. When protonated at nitrogen by a strong acid (CF₃SO₃H), it becomes a chiral super Lewis acid that can serve as an effective catalyst for a wide range of enantioselective Diels-Alder reactions [Angew. Chem., Int. Ed. Engl., 41, 1650 (2002); J. Am. Chem. Soc., 125, 6388 (2003)].

Sometimes, boron outsmarts us. During the 1950s, the U.S. government had a program to develop boron hydrides as rocket fuels based on the calculated energy release for combustion in O₂ to form (HO)₃B and H₂O. Unfortunately, much less energy is released, possibly because the actual product is HOBO at flame temperatures. For many years, it was hoped that organoboron compounds could be delivered into the brain for neutron beam therapy of brain tumors, taking advantage of the high neutron absorption cross section of ¹⁰B. This goal has so far not been realized. It is safe to say that boron chemistry is still underdeveloped.

E. J. Corey, who won the Nobel Prize in Chemistry in 1990, is the Sheldon Emory Research Professor at Harvard University. He is the recipient of more than 60 honorary degrees and awards.

BORON AT A GLANCE

Name: From the Arabic buraq, borax, its most important ore.

Atomic mass: 10.81.

History: Boron compounds have been known for years, but the pure form was not isolated until 1808 by British chemist Sir Humphry Davy and independently by French chemists Joseph-Louis Gay-Lussac and Louis-Jacques Thénard.

Occurrence: Makes up 0.0003% of Earth's crust by mass. It is never found in its pure form in nature, but occurs as orthoboric acid in certain volcanic spring waters and as borates in colemantie. Important sources of boron are the ores rasorite (kernite) and borax (tincal), which can be found in the Mohave Desert. Extensive borax deposits are also found in Turkey.

Appearance: Black-brown, solid metalloid.

Behavior: Crystalline boron is chemically inert. Excessive amounts of boric acid and borates are poisonous, although they were once used in medicines.

Uses: An essential micronutrient for plants. Boric acid is used to make borosilicate glass, or Pyrex. Boron is used in fission-reactor control rods to capture neutrons and regulate the power produced. It is also used with silicon in making p-type semiconductors and eye disinfectant and in pyrotechnics to make a green color.

ALUMINUM