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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.


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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.

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