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Borane
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| Names | |||
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| IUPAC names
Borane[1]
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| Systematic IUPAC name
Borane (substitutive) Trihydridoboron (additive) | |||
Other names
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| Identifiers | |||
3D model (JSmol)
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| ChEBI | |||
| ChemSpider | |||
| 44 | |||
PubChem CID
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| Properties | |||
| BH3 | |||
| Molar mass | 13.83 g·mol−1 | ||
| Appearance | colourless gas | ||
| Conjugate acid | Boronium | ||
| Thermochemistry | |||
Std molar
entropy (S⦵298) |
187.88 J mol−1 K−1 | ||
Std enthalpy of
formation (ΔfH⦵298) |
106.69 kJ mol−1 | ||
| Structure | |||
| D3h | |||
| trigonal planar | |||
| 0 D | |||
| Related compounds | |||
Other cations
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Alumane Gallane Indigane Thallane | ||
Related compounds
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Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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Borane is an inorganic compound with the chemical formula BH
3. Because it tends to dimerize or form adducts, borane is very rarely observed. It normally dimerizes to diborane in the absence of other chemicals.[2] It can be observed directly as a continuously produced, transitory, product in a flow system or from the reaction of laser ablated atomic boron with hydrogen.[3]
Structure and properties
[edit]BH3 is a trigonal planar molecule with D3h symmetry. The experimentally determined B–H bond length is 119 pm.[4]
In the absence of other bases, it dimerizes to form diborane. Thus, it is an intermediate in the preparation of diborane according to the reaction:[5]
- BX3 +BH4− → HBX3− + (BH3) (X=F, Cl, Br, I)
- 2 BH3 → B2H6
The standard enthalpy of dimerization of BH3 is estimated to be −170 kJ mol−1.[6] The boron atom in BH3 has 6 valence electrons. Consequently, it is a strong Lewis acid and reacts with any Lewis base ('L' in equation below) to form an adduct:[7]
- BH3 + L → L—BH3
in which the base donates its lone pair, forming a dative covalent bond. Such compounds are thermodynamically stable, but may be easily oxidised in air. Solutions containing borane dimethylsulfide and borane–tetrahydrofuran are commercially available; in tetrahydrofuran a stabilising agent is added to prevent the THF from oxidising the borane.[8] A stability sequence for several common adducts of borane, estimated from spectroscopic and thermochemical data, is as follows:
BH3 has some soft acid characteristics as sulfur donors form more stable complexes than do oxygen donors.[5] In water, it readily hydrolyzes into boric acid:[9][10]
- BH
3 + 3H2O → B(OH)
3 + 3 H
2
Reactions
[edit]Molecular species BH3 is a very strong Lewis acid. It can be isolated in the form of various adducts, such as borane carbonyl, BH3(CO).[11]
Molecular BH3 is believed to be a reaction intermediate in the pyrolysis of diborane to produce higher boranes:[5]
- B2H6 ⇌ 2BH3
- BH3 +B2H6 → B3H7 +H2 (rate determining step)
- BH3 + B3H7 ⇌ B4H10
- B2H6 + B3H7 → BH3 + B4H10
- ⇌ B5H11 + H2
Further steps give rise to successively higher boranes, with B10H14 as the most stable end product contaminated with polymeric materials, and a little B20H26.
Borane ammoniate, which is produced by a displacement reaction of other borane adducts, eliminates elemental hydrogen on heating to give borazine (HBNH)3.[12]
Borane adducts are widely used in organic synthesis for hydroboration, where BH3 adds across the C=C bond in alkenes to give trialkylboranes:[13]
- (THF)BH3 + 3 CH2=CHR → B(CH2CH2R)3 + THF
This reaction is regioselective.[14] Other borane derivatives can be used to give even higher regioselectivity.[15] The product trialkylboranes can be converted to useful organic derivatives. With bulky alkenes one can prepare species such as [HBR2]2, which are also useful reagents in more specialised applications. Borane dimethylsulfide which is more stable than borane–tetrahydrofuran may also be used.[16][15]
Hydroboration can be coupled with oxidation to give the hydroboration-oxidation reaction. In this reaction, the boryl group in the generated organoborane is substituted with a hydroxyl group.[17]
As a Lewis acid
[edit]Phosphine-boranes, with the formula R3−nHnPBH3, are adducts of organophosphines and borane. Borane adducts with amines are more widely used.[18] Borane makes a strong adduct with triethylamine; using this adduct requires harsher conditions in hydroboration. This can be advantageous for cases such as hydroborating trienes to avoid polymerization. More sterically hindered tertiary and silyl amines can deliver borane to alkenes at room temperature.
Borane(5) is the dihydrogen complex of borane. Its molecular formula is BH5 or possibly BH3(η2-H2).[19] It is only stable at very low temperatures and its existence is confirmed in very low temperature.[20][21] Borane(5) and methanium (CH5+) are isoelectronic.[22] Its conjugate base is the borohydride anion.
See also
[edit]References
[edit]- ^ "Borane".
- ^ Carey, Francis A.; Sundberg, Richard J. (2007). Advanced Organic Chemistry: Part B: Reactions and Synthesis (5th ed.). New York: Springer. p. 337. ISBN 978-0387683546.
- ^ Tague, Thomas J.; Andrews, Lester (1994). "Reactions of Pulsed-Laser Evaporated Boron Atoms with Hydrogen. Infrared Spectra of Boron Hydride Intermediate Species in Solid Argon". Journal of the American Chemical Society. 116 (11): 4970–4976. doi:10.1021/ja00090a048. ISSN 0002-7863.
- ^ Kawaguchi, Kentarou (1992). "Fourier transform infrared spectroscopy of the BH3 ν3 band". The Journal of Chemical Physics. 96 (5): 3411–3415. Bibcode:1992JChPh..96.3411K. doi:10.1063/1.461942. ISSN 0021-9606.
- ^ a b c Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. doi:10.1016/C2009-0-30414-6. ISBN 978-0-08-037941-8.
- ^ Page, M.; Adams, G.F.; Binkley, J.S.; Melius, C.F. (1987). "Dimerization energy of borane". J. Phys. Chem. 91 (11): 2675–2678. doi:10.1021/j100295a001.
- ^ Carey, Francis A.; Sundberg, Richard J. (2007). Advanced Organic Chemistry: Part B: Reactions and Synthesis (5th ed.). New York: Springer. p. 337. ISBN 978-0387683546.
- ^ Hydrocarbon Chemistry, George A. Olah, Arpad Molner, 2d edition, 2003, Wiley-Blackwell ISBN 978-0471417828
- ^ Finn, Patricia; Jolly, William L. (August 1972). "Asymmetric cleavage of diborane by water. The structure of diborane dihydrate". Inorganic Chemistry. 11 (8): 1941–1944. doi:10.1021/ic50114a043.
- ^ D'Ulivo, Alessandro (May 2010). "Mechanism of generation of volatile species by aqueous boranes". Spectrochimica Acta Part B: Atomic Spectroscopy. 65 (5): 360–375. doi:10.1016/j.sab.2010.04.010.
- ^ Burg, Anton B.; Schlesinger, H. I. (May 1937). "Hydrides of boron. VII. Evidence of the transitory existence of borine (BH
3): Borine carbonyl and borine trimethylammine". Journal of the American Chemical Society. 59 (5): 780–787. doi:10.1021/ja01284a002. - ^ Housecroft, C. E.; Sharpe, A. G. (2008). "Chapter 13: The Group 13 Elements". Inorganic Chemistry (3rd ed.). Pearson. p. 336. ISBN 978-0-13-175553-6.
- ^ Carey, Francis A.; Sundberg, Richard J. (2007). Advanced Organic Chemistry: Part B: Reactions and Synthesis (5th ed.). New York: Springer. p. 337. ISBN 978-0387683546.
- ^ Carey, Francis A.; Sundberg, Richard J. (2007). Advanced Organic Chemistry: Part B: Reactions and Synthesis (5th ed.). New York: Springer. p. 338. ISBN 978-0387683546.
- ^ a b Burkhardt, Elizabeth R.; Matos, Karl (July 2006). "Boron reagents in process chemistry: Excellent tools for selective reductions". Chemical Reviews. 106 (7): 2617–2650. doi:10.1021/cr0406918. PMID 16836295.
- ^ Kollonitisch, J. (1961). "Reductive Ring Cleavage of Tetrahydrofurans by Diborane". J. Am. Chem. Soc. 83 (6): 1515. doi:10.1021/ja01467a056.
- ^ Carey, Francis A.; Sundberg, Richard J. (2007). Advanced Organic Chemistry: Part B: Reactions and Synthesis (5th ed.). New York: Springer. p. 344. ISBN 978-0387683546.
- ^ Carboni, B.; Mounier, L. (1999). "Recent developments in the chemistry of amine- and phosphine-boranes". Tetrahedron. 55 (5): 1197. doi:10.1016/S0040-4020(98)01103-X.
- ^ Szieberth, Dénes; Szpisjak, Tamás; Turczel, Gábor; Könczöl, László (19 August 2014). "The stability of η2-H2 borane complexes – a theoretical investigation". Dalton Transactions. 43 (36): 13571–13577. doi:10.1039/C4DT00019F. PMID 25092548.
- ^ Tague, Thomas J.; Andrews, Lester (1 June 1994). "Reactions of Pulsed-Laser Evaporated Boron Atoms with Hydrogen. Infrared Spectra of Boron Hydride Intermediate Species in Solid Argon". Journal of the American Chemical Society. 116 (11): 4970–4976. doi:10.1021/ja00090a048.
- ^ Schreiner, Peter R.; Schaefer III, Henry F.; Schleyer, Paul von Ragué (1 June 1994). "The structure and stability of BH5. Does correlation make it a stable molecule? Qualitative changes at high levels of theory". The Journal of Chemical Physics. 101 (9): 7625. Bibcode:1994JChPh.101.7625S. doi:10.1063/1.468496.
- ^ A Life of Magic Chemistry: Autobiographical Reflections Including Post-Nobel Prize Years and the Methanol Economy, 159p
Borane
View on Grokipediafrom Grokipedia
Overview and Nomenclature
Definition and Classification
Boranes are a class of neutral chemical compounds composed exclusively of boron and hydrogen atoms, with the general formula BₙHₘ where n ≥ 1 and m ≥ 3.[5] These compounds are distinct from ionic borohydrides, which are anionic species such as [BH₄]⁻, and represent covalent boron hydrides that exhibit electron deficiency due to insufficient valence electrons for conventional two-center bonding.[6] Boranes are broadly classified into classical and non-classical types based on their structural motifs. Classical boranes feature open-chain or simple bridged structures, exemplified by monomeric borane (BH₃) and diborane (B₂H₆), which rely on localized bonds including three-center two-electron bonds.[7] In contrast, non-classical boranes encompass cluster or polyhedral structures, often derived from precursors like the closo-borane dianions [BₙHₙ]²⁻, where bonding involves delocalized electrons across a cage framework.[5] To distinguish boranes from related species, they exclude carboranes, which incorporate carbon atoms into the boron-hydrogen framework (e.g., C₂B₁₀H₁₂), and borohydrides, which are charged and typically metal salts like Na[BH₄]. Simple boranes such as BH₃ and B₂H₆ represent the foundational members with low boron counts, while higher boranes like pentaborane (B₅H₉) and decaborane (B₁₀H₁₄) illustrate more complex neutral clusters.[2] As electron-deficient compounds, boranes—particularly the polyhedral clusters—adhere to Wade's rules for structural classification, which predict geometries based on skeletal electron pairs: closo structures for 2n+2 pairs (closed polyhedra), nido for 2n+4 pairs (one vertex removed), and arachno for 2n+6 pairs (two vertices removed).[8] This electron-counting framework, developed by Kenneth Wade in 1971, provides a systematic basis for understanding the diversity of borane architectures beyond simple chains.[9]Naming Conventions
The nomenclature of boranes follows systematic rules established by the International Union of Pure and Applied Chemistry (IUPAC) to distinguish them from other boron compounds, emphasizing the number of boron atoms, hydrogen content, and structural features.[10] For simple, neutral molecular hydrides, the parent hydride is named "borane" preceded by a numerical prefix indicating the number of boron atoms, with the total number of hydrogen atoms specified in parentheses, such as borane for $ \mathrm{BH_3} $, diborane(6) for $ \mathrm{B_2H_6} $, and pentaborane(9) for $ \mathrm{B_5H_9} $.[10] This stoichiometric approach ensures clarity in composition without implying structure, and it applies similarly to anionic species named as "hydridoborates," for example, decahydridodecaborate(2−) for $ [\mathrm{B_{12}H_{12}}]^{2-} $.[10] For cluster boranes, structural nomenclature incorporates descriptors derived from polyhedral boron hydride classifications—such as closo for closed deltahedral structures, nido for those with one missing vertex, and arachno for two missing vertices—to reflect the skeletal arrangement, as briefly referenced in borane classification systems.[10] These descriptors precede the boron count and parent hydride name, yielding names like closo-dodecaborane(12) for $ \mathrm{B_{12}H_{12}} $, nido-pentaborane(9) for $ \mathrm{B_5H_9} $, and arachno-decaborane(14) for $ \mathrm{B_{10}H_{14}} $.[10] Locants for hydrogen positions further specify bridging (μ), exo (radial), or endo (tangential) orientations, as in 1,2-μH₂-diborane(6) for the bridged structure of $ \mathrm{B_2H_6} $.[10] Substitutive nomenclature is used for boranes where hydrogen atoms are replaced by substituents, treating the parent borane as the base and adding prefixes for the groups, while retaining the hydrogen count in parentheses; for instance, dimethylborane for $ (\mathrm{CH_3})2\mathrm{BH} $ and 1-methyl-nido-pentaborane(9) for a methylated $ \mathrm{B_5H_8(CH_3)} $.[10] In contrast, additive nomenclature applies to coordination compounds where two-electron donor ligands are attached, often altering the effective structure, and uses ligand names with locants, such as 6,9-bis(dimethylsulfane)-arachno-decaborane(12) for $ \mathrm{B{10}H_{12}(S(CH_3)_2)_2} $.[10] The distinction ensures substitutive names for one-electron replacements maintain the original hydrogen framework, while additive names account for ligand coordination without implying substitution.[10] Historically, earlier naming conventions, such as those from the 1950s, relied more on empirical formulas or ad hoc terms like "diborane" without the (6) specifier, but modern IUPAC rules from 2019 standardize the parenthetical hydrogen count and replace outdated prefixes like "hydro" with "hydrido" in additive contexts to align with broader inorganic nomenclature.[10] Acronyms are generally avoided in formal naming, though common retained names like diborane persist for $ \mathrm{B_2H_6} $ despite the preferred diborane(6).[10] These conventions facilitate precise communication in scientific literature, particularly for distinguishing borane clusters from carboranes or metalloboranes.[10]History
Discovery
The discovery of boranes traces back to 1912, when German chemist Alfred Stock first isolated diborane (B₂H₆) through the reaction of magnesium boride with hydrochloric acid, marking the initial identification of these volatile, reactive compounds.[11] Early investigations highlighted significant challenges, as the monomeric borane (BH₃) proved highly unstable and spontaneously dimerized to form the more persistent B₂H₆, shifting Stock's focus to this dimer and prompting his extensive systematic exploration of higher boranes throughout the 1920s and 1930s using innovative high-vacuum techniques.[12] A pivotal advancement occurred in 1936 with Alfred Stock's isolation of pentaborane (B₅H₉) through pyrolysis of diborane, which broadened the family of known boranes and facilitated further studies on their structural diversity.[13] By the 1950s, the concept of electron-deficient bonding in boranes gained recognition, building on Kenneth S. Pitzer's 1945 proposal of multicenter bonds to explain their unconventional structures, thereby establishing boranes as early exemplars of non-classical bonding that predated fullerene chemistry.[14]Key Developments
In the mid-20th century, borane chemistry advanced significantly through experimental innovations in organic synthesis. Herbert C. Brown pioneered hydroboration in the 1950s, demonstrating the addition of borane (BH₃) to alkenes to form organoboranes, a method that provided anti-Markovnikov regioselectivity and syn stereochemistry unavailable with traditional reagents. This work, initiated with the first reported hydroboration in 1956, transformed synthetic methodology and earned Brown the Nobel Prize in Chemistry in 1979 for developing boron-containing compounds in organic reactions.[15] Concurrently, the preparation of stable borane adducts, such as BH₃·THF, emerged in the early 1960s, enabling safer handling and broader application of reactive borane species beyond the volatile diborane. The 1970s marked a theoretical milestone with Kenneth Wade's formulation of electron-counting rules for polyhedral boranes, published in 1971, which correlated the number of skeletal electron pairs with cluster geometries—closo for n+1 pairs, nido for n+2, and arachno for n+3, where n is the number of vertices. These rules, derived from structural analyses of neutral and anionic boranes, provided a predictive framework for electron-deficient clusters and extended to carboranes and metallaboranes, fundamentally shaping understanding of three-dimensional bonding in main-group chemistry. From the 1980s to the 2000s, experimental and computational progress deepened insights into borane bonding and stability. Computational modeling, leveraging density functional theory, revealed the multicenter bonding in bridging hydrides and polyhedral frameworks, confirming Wade's predictions and explaining reactivity trends. Stable adducts like BH₃·THF became staples in synthesis. During the 1940s and 1950s, interest in boranes as high-energy rocket fuels in the U.S. led to scaled-up production and further structural investigations.[16] Post-2010 developments have emphasized supramolecular and applied aspects of boranes. Supramolecular borane complexes, exploiting their Lewis acidity for host-guest interactions, have enabled catalytic applications in hydrogen activation and small-molecule transformations, as reviewed in studies on boron-ligand cooperation mechanisms. In materials science, 2020s research on B₁₀H₁₄ (decaborane) derivatives has focused on high-temperature pyrolysis to generate boron-rich materials for aerospace, yielding thermally stable microcrystals suitable for propulsion and thermal protection systems. Overall, boranes have profoundly impacted cluster chemistry by modeling electron delocalization and materials science via incorporation into polymers, enhancing thermal and chemical resilience in advanced composites.[17][18]Structure and Bonding
Monomeric Borane
Monomeric borane, BH₃, exhibits a trigonal planar geometry with D_{3h} symmetry, characterized by three equivalent B-H bonds. The experimentally determined B-H bond length is 118.5 pm, as obtained from rovibrational spectroscopy.[19] The electronic structure of BH₃ arises from sp² hybridization of the boron atom, forming three σ bonds with the hydrogen atoms using its 2s and 2p_x, 2p_y orbitals, while the unoccupied 2p_z orbital remains empty. This configuration results in only six valence electrons in the valence shell of boron, creating an electron-deficient species that manifests strong Lewis acidity.[20] Due to its inherent instability, BH₃ spontaneously dimerizes to form diborane, B₂H₆, with a dimerization enthalpy of approximately -160 kJ/mol at 300 K. The monomer persists only under specialized conditions, such as in the gas phase or low-temperature matrices, or when stabilized as adducts with Lewis bases, including ammonia-borane (BH₃·NH₃) and tetrahydrofuran-borane (BH₃·THF).[19][1] The trigonal planar structure of gaseous BH₃ has been confirmed through infrared spectroscopy, which reveals vibrational modes (such as the ν₃ asymmetric stretch near 2600 cm⁻¹) consistent with D_{3h} symmetry. Additionally, ¹¹B NMR spectroscopy of the adducts displays characteristic chemical shifts (typically around -10 to 0 ppm for BH₃·THF), indicative of the tetrahedral coordination at boron and supporting the presence of intact monomeric BH₃ units.[21]Diborane and Bridging Hydrides
Diborane (B₂H₆) features a dimeric structure composed of two BH₂ units connected by two bridging hydrogen atoms, resulting in four terminal B-H bonds and two three-center-two-electron (3c-2e) B-H-B bridges that stabilize the electron-deficient system.[22] This arrangement adopts D_{2h} symmetry, with the boron atoms forming a rectangular core alongside the bridges; the terminal B-H bonds are conventional two-center-two-electron (2c-2e) bonds, while the bridges involve delocalized electron density shared among the boron atoms and the bridging hydrogen.[23] Unlike the monomeric BH₃, which dimerizes spontaneously due to its instability, the bridging hydrides in diborane provide the necessary stabilization for the overall framework.[24] The bonding in the bridges is characterized by banana-shaped molecular orbitals, where two electrons occupy a delocalized orbital spanning the B-H-B unit, often described as 3c-2e bonds.[22] Molecular orbital theory further elucidates this by showing that the bridge bonds arise from the overlap of sp³ hybrid orbitals on each boron with the 1s orbital of the bridging hydrogen, leading to σ-type bonding with partial multicentered character; the terminal bonds, in contrast, are localized σ bonds from similar hybrid orbitals.[25] Bond angles reflect this asymmetry: the H_terminal-B-H_terminal angle is approximately 121°, approaching trigonal planar geometry around the terminal hydrogens, while the bridge angles, such as ∠H_bridge-B-H_bridge, are about 97°, and the B-H_bridge-B angle is roughly 90°, indicative of the compressed tetrahedral environment at boron.[26] Substituted derivatives, such as methyldiborane (CH₃B₂H₅), maintain a similar bridged architecture, where one terminal hydrogen is replaced by an alkyl group, preserving the 3c-2e bonds but introducing asymmetry in the BH₂ unit bearing the substituent. Microwave spectroscopy confirms this structure, with the methyl group adopting a position that minimally disrupts the core dimer framework, and bond lengths adjusted slightly due to the steric influence of the alkyl moiety. In comparison, the aluminum analog, dialane (Al₂H₆), exhibits a parallel dimeric form with two 3c-2e Al-H-Al bridges and four terminal Al-H bonds, but the larger atomic radius of aluminum results in wider bridge angles (around 80° for Al-H-Al) and longer overall distances, highlighting periodic trends in group 13 hydride stability. The stability of diborane's pyramidal configuration at each boron atom is underscored by a high barrier to pyramidal inversion of the terminal BH₂ groups, estimated at over 30 kcal/mol from computational models, preventing facile rearrangement to a monomeric form under ambient conditions. Raman spectroscopy provides insight into this stability through its vibrational modes, revealing characteristic bands including the symmetric stretch of the bridging B-H-B units at approximately 1170 cm⁻¹ and terminal B-H stretches near 2520 cm⁻¹, which collectively confirm the bridged structure and its dynamic integrity.[26] These modes, analyzed via isotopic substitution, align with the predicted 18 vibrational degrees of freedom for D_{2h} symmetry, emphasizing the role of the bridges in distributing electron deficiency across the molecule.[26]Polyhedral Boranes
Polyhedral boranes represent a class of larger boron hydride clusters characterized by deltahedral geometries, where boron atoms occupy vertices of polyhedra, and hydrogen atoms serve as terminal or bridging ligands. These structures arise from the electron-deficient nature of boron, leading to multicenter bonding that delocalizes electrons over the cluster framework. The geometries and stability of polyhedral boranes are rationalized by the Wade-Mingos rules, which correlate the number of skeletal electron pairs with the cluster's shape.[27] The bonding in polyhedral boranes involves delocalized multicenter molecular orbitals, often described as three-center two-electron (3c-2e) bonds, that form the skeletal framework. Each boron atom contributes three valence electrons, but after accounting for terminal B-H bonds (one electron pair per B-H), the remaining skeletal electrons are counted in pairs. For closo clusters, which feature closed deltahedral polyhedra with n vertices, the formula is $ \mathrm{B_n H_n^{2-}} $, requiring $ n + 1 $ skeletal electron pairs or $ 2n + 2 $ electrons for stability. This electron count supports a closed cage structure, as exemplified by the octahedral $ [\mathrm{B_6 H_6}]^{2-} $, where 14 skeletal electrons fill seven bonding molecular orbitals derived from 18 atomic orbitals.[9] Cluster types are classified based on their relation to a hypothetical closo parent polyhedron and the number of skeletal electron pairs. Closo clusters have $ n + 1 $ pairs and a complete deltahedron. Nido clusters, derived by removing one vertex from a closo structure and adding a proton, possess $ n + 2 $ pairs and feature an open face, as in $ \mathrm{B_5 H_9} [\mathrm{B_6 H_6}]^{2-} $. Arachno clusters remove two vertices, yielding $ n + 3 $ pairs and more open structures like butterfly-shaped $ \mathrm{B_4 H_{10}} \mathrm{B_{10} H_{14}} n = 10 $), exhibiting a structure resembling an incomplete dodecahedron with an open pentagonal face; four boron atoms near this face bear bridging hydrogens, while the rest have terminal hydrides. This geometry was confirmed by X-ray crystallography, revealing bond lengths consistent with delocalized bonding, and further supported by $ ^{11}\mathrm{B} $ NMR spectroscopy, which displays five distinct signals corresponding to the unique boron environments.[9] For larger or more irregular clusters, the Wade-Mingos rules extend to hypho and klado types. Hypho clusters, with $ n + 4 $ skeletal electron pairs, correspond to structures missing three vertices from a closo parent, resulting in net-like open frameworks. Klado clusters, possessing $ n + 5 $ pairs, involve four missing vertices and exhibit branched or irregular shapes. These extensions apply primarily to derivatives rather than simple boranes, accommodating increased hydrogen content and structural diversity in higher polyhedra.[28][9]Synthesis
Laboratory Preparation
Laboratory preparation of boranes requires strict inert atmosphere conditions due to their high reactivity with air and moisture. The simplest stable borane, diborane (B₂H₆), is commonly synthesized on a small scale by the reaction of sodium borohydride (NaBH₄) with boron trifluoride diethyl etherate (BF₃·OEt₂) in diglyme solvent at room temperature. The ionic representation of the reaction is BH₄⁻ + 4 BF₃ → B₂H₆ + 4 BF₄⁻, with the balanced molecular equation being 3 NaBH₄ + 4 BF₃ → 2 B₂H₆ + 3 NaBF₄. This method proceeds over 2-4 hours, yielding 70-80% diborane based on NaBH₄, and the product is isolated by distillation under vacuum.[29] Borane adducts, such as BH₃ coordinated to ethers or sulfides, are prepared by reacting sodium borohydride (NaBH₄) with boron trifluoride (BF₃) in THF or by oxidation of NaBH₄ with iodine in THF. For instance, NaBH₄ + BF₃ in THF generates the BH₃·THF adduct along with NaF precipitate, which is filtered under inert conditions.[2] These reactions are typically conducted at low temperatures (0-25°C) to prevent decomposition, affording stable adducts suitable for further use in synthesis without the need for gaseous diborane. Higher boranes, including pentaborane (B₅H₉) and decaborane (B₁₀H₁₄), are obtained via thermal pyrolysis of diborane in a flow system or static reactor at temperatures around 200°C. The decomposition follows pathways such as 5 B₂H₆ → 2 B₅H₉ + 6 H₂ for B₅H₉ formation, with yields up to 60% based on 90% diborane conversion. Products are separated by fractional condensation using vacuum lines, often at -60°C for B₅H₉ and higher for B₁₀H₁₄. Since monomeric borane (BH₃) is unstable and dimerizes readily, these methods focus on diborane as the precursor.[30] All borane preparations demand specialized equipment like Schlenk lines or gloveboxes to maintain an inert atmosphere, as these compounds are pyrophoric and ignite spontaneously in air. Yields and conditions must be optimized to minimize side reactions, with typical overall efficiencies of 70-80% for diborane-based syntheses; operators must employ fire-resistant PPE, extinguishers suitable for metal fires, and protocols to quench spills with dry sand or inert powders.[31]Industrial Production
Diborane, the primary borane compound produced industrially, is manufactured through the continuous hydrogenation of boron trichloride with hydrogen gas. The process involves reacting BCl₃ and H₂ at elevated temperatures around 500°C in a silver-lined reactor under pressure, yielding diborane and HCl as a byproduct, with rapid quenching to maximize selectivity.[32] This method enables scalable production by recycling unreacted gases and intermediate dichloroborane, which disproportionates to form additional diborane.[32] An alternative commercial route for diborane synthesis employs the reduction of boron trifluoride with sodium hydride or lithium hydride at controlled temperatures to avoid side reactions.[33] These processes address the compound's reactivity and toxicity through specialized handling in closed systems. Borane adducts, such as BH₃·SMe₂ and BH₃·THF, are key commercial forms of borane, produced by complexing diborane with dimethyl sulfide or tetrahydrofuran, or via intermediates from sodium borohydride and boron halides. These stable, liquid reagents are manufactured in ton-scale quantities annually for use in fine chemical synthesis, offering safer alternatives to gaseous diborane.[34] Higher boranes, including polyhedral species like B₅H₉ and B₁₀H₁₄, are synthesized in limited quantities using plasma pyrolysis of lower boranes. These methods face significant challenges, including high energy consumption, equipment corrosion, and the inherent toxicity of boranes requiring stringent safety protocols.[35]Properties
Physical Properties
Boranes display diverse physical states influenced by molecular size and structure. Diborane (B₂H₆), the parent borane, exists as a colorless, toxic gas at room temperature, characterized by a repulsive sweet odor and high volatility, with a melting point of -165 °C and a boiling point of -92.5 °C.[36] Its vapor density is approximately 1.23 g/L at standard temperature and pressure, reflecting its molecular weight of 27.67 g/mol, and it is less dense than air (relative density 0.96).[37] Higher boranes transition to more stable forms: pentaborane(9) (B₅H₉) is a colorless, volatile liquid, while decaborane(14) (B₁₀H₁₄) appears as white crystalline needles or solids with a density of 0.94 g/cm³ at 25 °C, a melting point of 99.6 °C, and a boiling point of 213 °C.[38] Solubility profiles of boranes favor nonpolar organic solvents over aqueous media. Diborane exhibits high solubility in ethers like diethyl ether and diglyme (e.g., up to 0.4 M in diglyme at 0 °C) but decomposes rapidly in water via hydrolysis, precluding true solubility measurements.[39] Similarly, higher boranes such as decaborane show slight solubility in cold water (with hydrolysis in hot water) but good solubility in benzene, ethyl acetate, alcohols, and carbon disulfide.[38] Spectroscopic properties provide key insights into borane structures. In ¹¹B NMR, diborane shows a single broad signal at δ ≈ 16.6 ppm at room temperature due to rapid hydrogen exchange between terminal and bridging positions. At lower temperatures, separate signals for the terminal and bridging boron environments can be resolved.[40] Infrared spectroscopy reveals characteristic B-H stretching vibrations in the 2350-2630 cm⁻¹ range, with terminal B-H modes around 2500-2600 cm⁻¹ and bridging B-H-B modes lower, near 2100-2300 cm⁻¹, enabling identification of cluster types.[2] Thermodynamic parameters underscore the instability and reactivity of boranes. Gaseous diborane has a standard molar heat capacity of 57.57 J mol⁻¹ K⁻¹ at 298 K and an entropy of 232.49 J mol⁻¹ K⁻¹, with vapor pressures increasing rapidly (e.g., 224 mmHg at -112 °C, reaching 1 atm at the boiling point of -92.5 °C).[41][36] Flammability is extreme: diborane ignites spontaneously in air at 40-50 °C and forms explosive mixtures over a broad range of 0.8-88 vol% in air, necessitating inert handling conditions.[42] Higher boranes share pyrophoric tendencies, with decaborane flammable above 80 °C.[38]Chemical Stability
Boranes, particularly diborane (B₂H₆), exhibit limited thermal stability due to their tendency to undergo decomposition at elevated temperatures. Diborane begins to decompose above 0°C, forming higher boranes and hydrogen gas, with the rate increasing exponentially with temperature and concentration in the gas phase.[36][43] In contrast, borane adducts such as BH₃ coordinated to amines demonstrate significantly enhanced thermal stability, remaining intact up to approximately 100°C before decomposition initiates, making them preferable for practical handling.[44] Hydrolysis represents a primary degradation pathway for boranes, occurring rapidly upon contact with water. For diborane, the reaction proceeds vigorously and exothermically, yielding boric acid and hydrogen gas as shown in the equation:
This process is kinetically fast in neutral or basic conditions, with the rate influenced by pH; acidic environments accelerate hydrolysis by promoting protonation and bond cleavage, while the overall reaction is essentially instantaneous at room temperature without catalysts.[45][46]
Boranes are highly air-sensitive, with diborane igniting spontaneously in moist air through concurrent oxidation and hydrolysis, forming explosive mixtures over a wide concentration range (0.8–88% in air).[3] Among borane adducts, air stability follows the donor strength of the Lewis base: amine adducts > sulfide adducts > ether adducts, where stronger coordination in amine complexes provides greater protection against atmospheric moisture and oxygen.
In terms of oxidative stability, boranes show resistance to dry oxygen when stored under inert atmospheres such as nitrogen or argon, preventing spontaneous combustion; however, trace moisture dramatically reduces this tolerance. Storage half-lives for diborane in dilute mixtures reach up to 6 months at temperatures below 5°C, with degradation primarily attributed to thermal rather than oxidative pathways under these conditions.[43][47]
Reactions
Lewis Acid Behavior
Borane (BH₃) functions as a Lewis acid owing to the electron deficiency of its central boron atom, which adopts a trigonal planar geometry with an empty p-orbital available to accept an electron pair from a Lewis base, forming a dative bond. This reactivity arises from boron's incomplete octet in the monomeric form, enabling coordination without significant π-backdonation from typical donor ligands.[48] The resulting adducts are often stable under ambient conditions, though their formation constants vary with the base's nucleophilicity and steric profile; for instance, the ammonia-borane complex (BH₃·NH₃) exhibits a favorable equilibrium in solution, reflecting borane's strong affinity for nitrogen donors.[49] Common adducts include amine-boranes such as BH₃·NH₃ and phosphine-boranes like trimethylphosphine-borane (BH₃·PMe₃), which serve as convenient storage forms for reactive borane due to their enhanced thermal stability compared to free BH₃. Another notable example is the carbonyl complex (H₃B·CO), a volatile gas that can be isolated at room temperature but exists in equilibrium with diborane and carbon monoxide, with the dissociation becoming faster at higher temperatures.[50] Phosphine-boranes, in particular, are valued in synthetic applications for protecting phosphorus centers and delivering borane equivalents without the hazards of diborane.[51] The Lewis acidity of BH₃ surpasses that of alkyl-substituted analogs like trimethylborane (BMe₃), primarily because the smaller hydrogen substituents impose minimal steric repulsion, facilitating closer approach and stronger binding of Lewis bases.[52] In contrast, the bulkier methyl groups in BMe₃ hinder coordination, reducing its acid strength despite similar electronic profiles. Exotic variants like borane(5) (BH₅), a hypervalent species isoelectronic with methanium (CH₅⁺), which is theoretically metastable at very low temperatures.[53] In coordination chemistry, borane's Lewis acidity enables the formation of ate complexes, such as lithium tetrahydroborate (Li[BH₄]), where BH₃ accepts a hydride (H⁻) ligand to achieve tetrahedral coordination around boron, stabilizing the anion in ionic lattices.[54] These complexes highlight borane's versatility in expanding its valence shell beyond four electrons, influencing properties like hydrogen storage in metal borohydrides.[55]Hydroboration Reactions
Hydroboration reactions involve the addition of borane (BH₃) or its derivatives across the double or triple bonds of alkenes and alkynes, respectively, providing a versatile method for carbon-boron bond formation in organic synthesis. Discovered by Herbert C. Brown and B. C. Subba Rao in 1956 during studies on alkene reductions, this process proceeds under mild conditions, typically at room temperature, and contrasts with traditional electrophilic additions by following anti-Markovnikov regiochemistry, where the boron atom attaches to the less substituted carbon.[56] The reaction's stereospecificity results in syn addition of the B-H bond, delivering both boron and hydrogen from the same face of the unsaturated substrate.[56] The mechanism begins with the concerted insertion of the alkene into the B-H bond of BH₃, forming a four-center transition state that favors the anti-Markovnikov orientation due to the partial positive charge on boron and steric factors directing it to the less hindered position. For example, borane reacts with three equivalents of a terminal alkene such as 1-hexene (RCH=CH₂, where R = butyl) to yield a trialkylborane:
This trialkylborane can then undergo oxidation with hydrogen peroxide and hydroxide (H₂O₂/OH⁻) to produce the corresponding primary alcohol in high yield, often exceeding 90% for terminal alkenes.[56] The overall hydroboration-oxidation sequence thus achieves stereospecific anti-Markovnikov hydration of alkenes, a transformation not readily accessible by other methods.[56]
To enhance regioselectivity and enable selective monoaddition, dialkylboranes such as disiamylborane ((sia)₂BH) and 9-borabicyclo[3.3.1]nonane (9-BBN) are employed. Disiamylborane, prepared from BH₃ and 2-methyl-2-butene, exhibits high selectivity for less hindered terminal alkenes, minimizing isomer formation in sterically demanding substrates.[56] Similarly, 9-BBN provides exceptional regioselectivity, achieving over 98% purity in the hydroboration of terminal alkenes for applications like pheromone synthesis.[56]
The scope extends to alkynes, where borane derivatives like 9-BBN or dibromoborane-dimethyl sulfide enable selective monohydroboration to form vinylboranes, which can be further functionalized.[56] Asymmetric variants utilize chiral boranes, such as diisopinocampheylborane derived from α-pinene, to induce enantioselectivity in the addition, facilitating the synthesis of enantioenriched alcohols upon oxidation.[56]