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Diborane(6), generally known as diborane is the chemical compound consisting of boron and hydrogen with the formula B₂H₆. It is a colorless, pyrophoric gas with a repulsively sweet odor. Synonyms include boroethane, boron hydride, and diboron hexahydride. Diborane is a key boron compound with a variety of applications. It has attracted wide attention for its electronic structure. Its derivatives are useful reagents.

Structure and bonding

Bonding diagram of diborane (B₂H₆) showing with curved lines a pair of three-center two-electron bonds, each of which consists of a pair of electrons bonding three atoms, two boron atoms and a hydrogen atom in the middle.

Diborane adopts a D_2h structure containing four terminal and two bridging hydrogen atoms. The model determined by molecular orbital theory indicates that the bonds between boron and the terminal hydrogen atoms are conventional 2-center, 2-electron covalent bonds. The bonding between the boron atoms and the bridging hydrogen atoms is, however, different from that in molecules such as hydrocarbons. Having used two electrons in bonding to the terminal hydrogen atoms, each boron has one valence electron remaining for additional bonding. The bridging hydrogen atoms provide one electron each. Thus the B₂H₂ ring is held together by four electrons, an example of 3-center 2-electron bonding. This type of bond is sometimes called a 'banana bond'. The lengths of the B-H_bridge bonds and the B-H_terminal bonds are 1.33 and 1.19 Å respectively, and this difference in the lengths of these bonds reflects the difference in their strengths, the B-H_bridge bonds being relatively weaker. The weakness of the B-H_bridge vs B-H_terminal bonds is indicated by their vibrational signatures in the infrared spectrum, being ~2100 and 2500 cm⁻¹, respectively.^[6] The structure is isoelectronic with C₂H₆²⁺, which would arise from the diprotonation of the planar molecule ethene.^[7] Diborane is one of many compounds with such unusual bonding.^[8]

Of the other elements in Group IIIA, gallium is known to form a similar compound, digallane, Ga₂H₆. Aluminium forms a polymeric hydride, (AlH₃)_n, although unstable Al₂H₆ has been isolated in solid hydrogen and is isostructural with diborane.^[9]

Production and synthesis

Extensive studies of diborane have led to the development of multiple syntheses. Most preparations entail reactions of hydride donors with boron halides or alkoxides. The industrial synthesis of diborane involves the reduction of BF₃ by sodium hydride, lithium hydride or lithium aluminium hydride:^[10]

8 BF₃ + 6 LiH → B₂H₆ + 6 LiBF₄

Two laboratory methods start from boron trichloride with lithium aluminium hydride or from boron trifluoride ether solution with sodium borohydride. Both methods result in as much as 30% yield:

4 BCl₃ + 3 LiAlH₄ → 2 B₂H₆ + 3 LiAlCl₄

4 BF₃ + 3 NaBH₄ → 2 B₂H₆ + 3 NaBF₄

Older methods entail the direct reaction of borohydride salts with a non-oxidizing acid, such as phosphoric acid or dilute sulfuric acid.

2 BH₄⁻ + 2 H⁺ → 2 H₂ + B₂H₆

Similarly, oxidation of borohydride salts has been demonstrated and remains convenient for small scale preparations. For example, using iodine as an oxidizer:

2 NaBH
₄ + I
₂ → 2 NaI + B
₂H
₆ + H
₂

Another small-scale synthesis uses potassium hydroborate and phosphoric acid as starting materials.^[11]

Reactions

Diborane is a highly reactive and versatile reagent that has numerous applications.^[12] Its dominating reaction pattern involves formation of adducts with Lewis bases. Often such initial adducts proceed rapidly to give other products. It reacts with ammonia to form the diammoniate of diborane, DADB, with lesser quantities of ammonia borane depending on the conditions used. Diborane also reacts readily with alkynes to form substituted alkene products which will readily undergo further addition reactions.

As a pyrophoric substance, diborane reacts exothermically with oxygen to form boron trioxide and water, so much that it was considered as a possible rocket or ramjet propellant^[13][14][15] but discarded because back then it was too expensive and dangerous to handle:

2 B₂H₆ + 6 O₂ → 2 B₂O₃ + 6 H₂O (ΔH_r = –2035 kJ/mol = –73.47 kJ/g)

Diborane also reacts violently with water to form hydrogen and boric acid:

B₂H₆ + 6 H₂O → 2 B(OH)₃ + 6 H₂ (ΔH_r = –466 kJ/mol = –16.82 kJ/g)

Diborane also reacts with methanol to give hydrogen and trimethoxyborate ester:^[16]

B₂H₆ + 6 MeOH → 2 B(OMe)₃ + 6 H₂

Treating diborane with sodium amalgam gives NaBH₄ and Na[B₃H₈]^[16]
When diborane is treated with lithium hydride in diethyl ether, Lithium borohydride is formed:^[16]

B₂H₆ + 2 LiH → 2 LiBH₄

Diborane reacts with anhydrous hydrogen chloride or hydrogen bromide gas to give a boron halohydride:^[16]

B₂H₆ + HX → B₂H₅X + H₂ (X = Cl, Br)

Treating diborane with carbon monoxide at 470 K and 20 bar gives H₃BCO.^[16]

Reagent in organic synthesis

Diborane and its variants are central organic synthesis reagents for hydroboration, whereby alkenes add across the B-H bonds to give trialkylboranes. Diborane is used as a reducing agent roughly complementary to the reactivity of lithium aluminium hydride. The compound readily reduces carboxylic acids to the corresponding alcohols, whereas ketones react only sluggishly.

History

Diborane was first synthesised in the 19th century by hydrolysis of metal borides, but it was never analysed. From 1912 to 1936, the major pioneer in the chemistry of boron hydrides, Alfred Stock, undertook his research that led to the methods for the synthesis and handling of the highly reactive, volatile, and often toxic boron hydrides. He proposed the first ethane-like structure of diborane.^[17]Electron diffraction measurements by S. H. Bauer initially appeared to support his proposed structure.^[18][19]

Because of a personal communication with L. Pauling (who supported the ethane-like structure), H. I. Schlessinger did not specifically discuss 3-center-2-electron bonding in his then classic review in the early 1940s.^[20] The review does, however, discuss the C_2v structure in some depth, "It is to be recognized that this formulation easily accounts for many of the chemical properties of diborane..."

In 1943 an undergraduate student at Balliol College, Oxford, H. Christopher Longuet-Higgins, published the currently accepted structure together with R. P. Bell.^[21] This structure had already been described in 1921 by Dilthey.^[22] The years following the Longuet-Higgins/Bell proposal witnessed a colorful discussion about the correct structure. The debate ended with the electron diffraction measurement in 1951 by K. Hedberg and V. Schomaker, with the confirmation of the structure shown in the schemes on this page.^[23]

William Nunn Lipscomb, Jr. further confirmed the molecular structure of boranes using X-ray crystallography in the 1950s, and developed theories to explain its bonding. Later, he applied the same methods to related problems, including the structure of carboranes on which he directed the research of future Nobel Prize winner Roald Hoffmann. Lipscomb himself received the Nobel Prize in Chemistry in 1976 for his efforts.

Other uses

Diborane has been tested as a rocket propellant ^[24]. Complete combustion is strongly exothermic. However, combustion is not complete in the rocket engine, as some boron monoxide, BO, is produced. This mirrors the incomplete combustion of hydrocarbons, to produce carbon monoxide, CO.

Diborane has been used as a rubber vulcaniser, as a catalyst for hydrocarbon polymerisation, as a flame-speed accelerator, and as a doping agent for the production of semiconductors. It is also an intermediate in the production of highly pure boron for semiconductor production. It is also used to coat the walls of tokamaks to reduce the amount of heavy metal impurities in the core plasma.

Safety

The toxic effects of diborane are primarily due to its irritant properties. Short-term exposure to diborane can cause a sensation of tightness of the chest, shortness of breath, cough, and wheezing. These signs and symptoms can occur immediately or be delayed for up to 24 hours. Skin and eye irritation can also occur. Studies in animals have shown that diborane causes the same type of effects observed in humans.^{[citation needed]}

People exposed for a long time to low amounts of diborane have experienced respiratory irritation, seizures, fatigue, drowsiness, confusion, and occasional transient tremors.

References

Cited sourced

• 
  Haynes, William M., ed. (2011). CRC Handbook of Chemistry and Physics (92nd ed.). CRC Press. ISBN 978-1439855119.
  

External links

• International Chemical Safety Card 0432


• National Pollutant Inventory – Boron and compounds


• NIOSH Pocket Guide to Chemical Hazards


• U.S. EPA Acute Exposure Guideline Levels

Diborane

Names
IUPAC name Diborane(6)
Other names Boroethane Boron hydride Diboron hexahydride
Identifiers
CAS Number	19287-45-7 Y
3D model (JSmol)	Interactive image
ChEBI	CHEBI:33590 Y
ChemSpider	17215804 Y
ECHA InfoCard	100.039.021
EC Number	242-940-6
PubChem CID	6328200 wrong structure
RTECS number	HQ9275000
UNII	BS9K982N24
UN number	1911
InChI InChI=1S/B2H6/c1-3-2-4-1/h1-2H2 Y Key: KLDBIFITUCWVCC-UHFFFAOYSA-N Y InChI=1/B2H6/c1-3-2-4-1/h1-2H2 Key: KLDBIFITUCWVCC-UHFFFAOYAF
SMILES [BH2]1[H][BH2][H]1
Properties
Chemical formula	B2H6
Molar mass	27.67 g·mol−1
Appearance	Colorless gas
Odor	repulsive and sweet
Density	1.131 g/L[1]
Melting point	−164.85 °C (−264.73 °F; 108.30 K)[1]
Boiling point	−92.49 °C (−134.48 °F; 180.66 K)[1]
Solubility in water	Reacts[2]
Vapor pressure	39.5 atm (16.6°C)[2]
Structure
Coordination geometry	Tetrahedral (for boron)
Molecular shape	see text
Dipole moment	0 D
Thermochemistry
Heat capacity (C)	56.7 J/(mol·K)[3]
Std molar entropy (So298)	232.1 J/(mol·K)[3]
Std enthalpy of formation (ΔfHo298)	36.4 kJ/mol[3]
Hazards
Main hazards	highly flammable, reacts with water
Safety data sheet	See: data page
GHS pictograms
GHS signal word	Danger
GHS hazard statements	H220, H280, H314, H318, H330, H370, H372
GHS precautionary statements	P210, P260, P264, P270, P271, P280, P284, P301+330+331, P303+361+353, P304+340, P305+351+338, P307+311, P310, P314, P320, P321, P363, P377, P381, P403, P403+233, P405, P410+403, P501
NFPA 704	[5]
Autoignition temperature	38 °C (100 °F; 311 K)
Explosive limits	0.8%–88%[2]
Lethal dose or concentration (LD, LC):
LC50 (median concentration)	40 ppm (rat, 4 hr) 29 ppm (mouse, 4 hr) 40–80 ppm (rat, 4 hr) 159–181 ppm (rat, 15 min)[4]
LCLo (lowest published)	125 ppm (dog, 2 hr) 50 ppm (hamster, 8 hr)[4]
US health exposure limits (NIOSH):
PEL (Permissible)	TWA 0.1 ppm (0.1 mg/m3)[2]
REL (Recommended)	TWA 0.1 ppm (0.1 mg/m3)[2]
IDLH (Immediate danger)	15 ppm[2]
Related compounds
Related boron compounds	Decaborane BF3
Supplementary data page
Structure and properties	Refractive index (n), Dielectric constant (εr), etc.
Thermodynamic data	Phase behaviour solid–liquid–gas
Spectral data	UV, IR, NMR, MS
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
Y verify (what is YN ?)
Infobox references