6 Boron Compounds
Kaufmann, D. E., Science of Synthesis, (2005) 6, 1.
General Introduction
This volume describes the organometallic and organic chemistry of boron. Despite covering just a single element, an extremely broad spectrum of chemistry is discussed within. As for other volumes of Science of Synthesis, the authors are experts in their respective fields and have covered the most important developments selectively and critically. Applications of the compounds synthesized are emphasized and, wherever appropriate, experimental procedures are provided in the text.
Old manuscripts indicate that the Arabs and Persians knew of the mineral borax (sodium tetraborate decahydrate), the first and still the most important natural boron source, over 2000 years ago. Boron oxide has been detected in Chinese enamels from the 4th century B.C. Elemental boron was discovered as a gray powder independently in 1808 by Gay-Lussac and Thenard in France, and by Davy in England as the main reduction product of boric acid by means of potassium. Davy suggested the name “boron”, formed from the first syllable from borax and the second from carbon, recognizing important similarities with the latter. The green flame of burning triethylborane, the first representative of the group of organoboron compounds, was reported by Frankland and Duppa in 1860. It then took 50 more years before Stock,[1] working on this task during 1912–36, and later Lipscomb (Nobel Prize 1976)[2] investigated the synthesis, structures, and bonding properties of the novel class of hydroboranes. Subsequently, hydridopolyborates were discovered, followed by heteroboranes, with carboranes being the most important subclass of the latter. The important field of organic synthesis via organoboranes was, broadly, developed first by Brown (Nobel Prize 1979)[3] and co-workers since the middle of the 20th century, strongly accelerated by the discovery of the hydroboration reaction of alkenes and alkynes in 1956.[4] In 1959, the first diastereoselective,[5] and 2 years later a highly enantioselective,[6] hydroboration reaction by means of chiral, terpene-based hydroboranes developed by Brown and Zweifel marked the beginning of practical asymmetric synthesis.
Boron reagents have become very important standard tools for the synthetic chemist[7,8] because boron compounds combine a unique mixture of interesting properties such as Lewis acidity and facile though highly stereoselective oxidizability.[9] The boron atom is slightly larger than carbon. All trivalent boron compounds are planar, whereas tetravalent boron compounds assume tetrahedral or nearly tetrahedral geometry. The B—C bond (∼323 kJ·mol−1)[10,11] is not much weaker than a C—C bond (∼358 kJ·mol−1); this situation is also reflected by similar bond lengths (B—C ∼156 pm). These two classes of compounds are closely related; neutral tricoordinate boron compounds such as trimethylborane are isoelectronic with carbenium ions such as the corresponding tert-butyl cation, whereas boronic acids [R1B(OH)2] are isoelectronic with carboxylic acids, and negatively charged tetracoordinate borates such as the borohydride anion are isoelectronic with neutral hydrocarbons such as methane.[12]
Boron forms even stronger bonds with nitrogen, oxygen, fluorine, and chlorine. The B—H bond (∼375 kJ·mol−1) is also slightly weaker than the C—H bond. By means of three-center bonds, monoorganohydroboranes and diorganohydroboranes form dimers, which always display hydride bridges rather than alkyl bridges, as illustrated in Scheme 1.[13–15]
Scheme 1 Hydride Bridging in Monoorganohydroboranes and Diorganohydroboranes[13–15]
In contrast to hydroboranes, trialkylboranes are monomeric, which may be traced to hyperconjugation with the alkyl substituents. In organoboranes, the B—C bond has largely single σ-bond character. Electron-rich π-systems such as vinyl or aryl groups provide the adjacent B—C bond with partial double-bond character. In small-ring boron heterocycles such as the 2π-systems, 1H-borirenes (e.g., 1), structural and spectroscopic data point to extensive π-electron delocalization.[16] The four-membered dihydrodiboretes exist in two isomers, a planar 1,2-dihydro-1,2-borete 2 with two localized π-electrons and a “butterfly” shaped 1,3-dihydro-1,3-borete 3 with delocalized π-electrons.[17] The 4π-antiaromatic 1H-borole is labile, even in perarylated form 4 (Scheme 2).[18]
Scheme 2 Boron Heterocycles[16–18]
Substituents with electron lone pairs act similarly while the strength of the boron—heteroatom π-bonds decreases in the order nitrogen > fluorine > oxygen > sulfur > chlorine. All mesomeric interactions lead to a distinct stabilization of the particular organoboranes.
Boron generally functions as a Lewis acid. In particular, the stable complexes of boron trifluoride with diethyl ether and of the parent hydroborane with amines, ethers, and dimethyl sulfide are of high preparative value. Complexation of less-stable organoboron intermediates with Lewis bases is a proven method of stabilization and storage. Boronic acids also react as Lewis and not as protic acids.
The feasibility of both the highly stereoselective introduction of a boryl group (e.g., by means of a hydroboration reaction) and its subsequent substitution by either a proton or a carbon or heteroatom electrophile is an extremely important prerequisite for its manifold successful applications in stereodirected syntheses. In particular, the ready oxidizability of boryl groups brings about a high preparative potential. It also means that, in general, uncomplexed organoboron compounds of a low oxidation state have to be handled strictly under inert gas (nitrogen or argon) using Schlenk techniques or a glove box.[19,20] Boronic acids and their esters are usually quite stable in air and can therefore be handled as ordinary organic compounds without any special precautions. Frequently, the purification of cyclic boronates by column chromatography is also feasible. Because of the low polarity of the B—C bond, many organoboranes are stable in deoxygenated water.
The most important industrial source of boron compounds is still borax. Acidification with carbon dioxide yields boric acid which, together with its anhydride, is the key compound for most of the commercially available boron compounds important for organic synthesis, such as hydroborane–amine, –ether, and –sulfide adducts, metal borohydrides, the haloboranes, boric acid esters, boronic acids and esters, trialkylboranes, and diboron(4) compounds. The main routes to organoboron compounds are transmetalation reactions, boration reactions of unsaturated compounds, and ligand-exchange reactions.[21]
The most important analytical tool for organoboron compounds is NMR spectroscopy. Of the two boron isotopes, 10B (20%, I = 3) and 11B (80%, I = 3/2), the latter possesses superior NMR properties. The 11B NMR chemical shifts cover a broad range of about 250 ppm; they depend on the charge, the coordination number, and the substituents at boron.[22–25]
Vibrational spectroscopy is especially valuable for the structural elucidation of hydroboranes with strongly differing B—H stretching frequencies for bridging and terminal positions (1500–2600 cm−1).[12,23]
Molecular ions of organoboron compounds are frequently of low abundance or absent in EI mass spectra. Thermally unstable compounds such as boronic acids tend to form decomposition or condensation products under the evaporation conditions. ESI mass spectrometry has proven to be a mild and, therefore, valuable modern tool in many cases.
The effects of boron compounds on biological systems have been reviewed extensively.[26] Diborane is a toxic hazard (industrial exposure limit 0.1 ppm) as it is not oxidized in air immediately.[27] The highly volatile lower trialkylboranes (BR13, R1 = Me, Et, Pr) have noxious and lachrymatory properties and they tend to ignite spontaneously. The higher trialkylboranes should also be handled in a closed system. Parallel to their higher oxidation state, boronic acids and esters can be handled as typical organic compounds. During boron neutron capture therapy (BNCT) studies, indications were gathered that water soluble boronic acids tend to have low toxicity while fat-soluble boronic acids are moderately toxic.[28] In addition, it should be noted that some boronic acids and related compounds are specific enzyme inhibitors.[29] Up to 350 ppm, boric acid is tolerated well in the diet of rats and dogs; toxic effects are apparent at higher doses.[30] The lethal dose of boric acid for human adults is estimated to be 15–20 g.[27] Natural boron-containing compounds occur in plants, algae, and microorganisms.[31] Boric acid is known to be an essential micronutrient for plants; there is some evidence that traces of boron may also be important for animals.[32]
References
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