My watch list  


CAS number 19287-45-7
RTECS number HQ9275000
Molecular formula B2H6
Molar mass 27.67 g/mol
Appearance colorless gas
Density 1.18 g/l, gas (15 °C)
Melting point

−165 °C (108.15 K)

Boiling point

−92.5 °C (180.65 K)

Solubility in water Reacts
Tetrahedral (for boron)
Molecular shape see text
Dipole moment 0 D
MSDS External MSDS
EU classification not listed
NFPA 704
Flash point flammable gas
38 °C
Related Compounds
Related boron compounds Decaborane
Supplementary data page
Structure and
n, εr, etc.
Phase behaviour
Solid, liquid, gas
Spectral data UV, IR, NMR, MS
Except where noted otherwise, data are given for
materials in their standard state
(at 25 °C, 100 kPa)

Infobox disclaimer and references

Diborane is the chemical compound with the formula B2H6. It is a colorless gas at room temperature with a repulsively sweet odor. B2H6 mixes well with air, easily forming explosive mixtures. Diborane will ignite spontaneously in moist air at room temperature. Synonyms include boroethane, boron hydride, and diboron hexahydride.

B2H6 is a key boron compound with a variety of applications. The compound is endothermic as indicated by a positive heat of formation, ΔH°f of 36 kJ/mol. Despite its intrinsic instability, B2H6 is kinetically quite stable and possesses an extensive chemistry.


Structure and bonding

Diborane adopts a D2h structure containing four terminal and two bridging hydrides. The molecular orbital model indicates that the bonding between boron and the terminal hydrides is fairly conventional, not unlike the bonding between carbon and hydride in ethane. The bonding between the boron atoms and the bridging hydrogen atoms is, however, distinctive relative to ordinary hydrocarbons. Formed two 2-center, 2-electron bonds with the terminal hydrogen atoms, each boron "has" one electron remaining for additional bonding. The bridging hydrogen atoms each provide one electron each. Thus the B2H2 ring is held together with four electrons. The lengths of the B-Hbridge bonds and the B-Hterminal bonds are 1.33 and 1.19 Å, respectively. The difference in the lengths of these bonds reflects the difference in their strengths: the B-Hbridge bonds are relatively weaker. The bonding between the boron atoms and two bridging hydrides of diborane is an example of 3-center-2-electron bonding. The structure is equivalent to that for isoelectronic, C2H62+, which would arise from the diprotonation of the planar molecule ethene. B2H6 is one of many compounds with such unusual bonding. This 3-center-2-electron bond is sometimes called a 'banana bond'.[1]

Of the other elements in Group 13, only gallium appears to form a similar compound, digallane, Ga2H6. Aluminium forms a polymeric hydride. No hydrides of indium and thallium have yet been found.[2]

Production and syntheses

Diborane is so central and has been studied so often that many syntheses exist. Most preparations entail reactions of hydride donors with boron halides or alkoxides. The industrial synthesis involves the reduction of BF3:

2 BF3 + 6 NaH → B2H6 + 6 NaF

Two laboratory methods start from boron trichloride with lithium aluminium hydride or from boron trifluoride ether solution with sodium borohydride. Both methods yield in up to 30% of diborane:

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 sulphuric acid.

2 BH4- + 2 H+ → 2H2 + B2H6

Similarly, oxidation of borohydride salts has been demonstrated and remains convenient for small scale preparations:

2 NaBH4 + I2 → 2 NaI + B2H6 + H2


Diborane is a highly reactive and versatile reagent that has a large number of applications.[3] 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 borazine. Diborane also reacts readily with alkynes to form substituted alkene products which will readily undergo further addition reactions.

Diborane reacts with water to form hydrogen and boric acid.

The compound forms complexes with Lewis bases. Notable are the complexes with THF and dimethyl sulfide, both liquid compounds are popular reducing agents in organic chemistry. In these 1:1 complexes, boron assumes a tetrahedral geometry, being bound to three hydrides and the Lewis base (THF or Me2S). The THF adduct is usually prepared as a 1:5 solution in THF. The latter is indefinitely stable when stored under nitrogen at room temperature.

Reagent in organic synthesis

Diborane is the central organic synthesis reagent for hydroboration, whereby alkenes add across the B-H bonds to give trialkylboranes:

(THF)BH3 + 3 CH2=CHR → B(CH2CH2R)3 + THF

This reaction is regioselective, and 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 specialized applications.

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.


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.[4] Electron diffraction from S. H. Bauer appeared to initially supported his proposed structure.[5][6]

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 1940's.[7] The review does, however, discuss the C2v 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, H. Christopher Longuet-Higgins, in Oxford published the currently accepted structure together with R. P. Bell.[8] This structure had already been described in 1921.[9][10][11] 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.[12]

William Nunn Lipscomb, Jr. further confirmed the molecular structure of boranes using X-ray crystallography in the 1950's, 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 received The Nobel Prize in Chemistry in 1976 for his efforts.

Other uses

Diborane is used in rocket propellants, 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.


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.

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


  1. ^ Laslo P (2000). "A Diborane Story". Angewandte Chemie International Edition 39: 2071-2072. doi:<2071::AID-ANIE2071>3.0.CO;2-C 10.1002/1521-3773(20000616)39:12<2071::AID-ANIE2071>3.0.CO;2-C. abstract
  2. ^ Downs, Anthony J.; Colin R. Pulham (1994). "The hydrides of aluminium, gallium, indium and thallium: A re-evaluation". Chemical Society Reviews 23 (3): 175 - 184. Cambridge: Royal Society of Chemistry. ISSN 0306-0012.
  3. ^ Mikhailov BM (1962). "The Chemistry of Diborane". Russian Chemical Review (31): 207-224.
  4. ^ Stock A. (1933). The Hydrides of Boron and Silicon. New York: Cornell University Press. 
  5. ^ Bauer S.H. (1937). "The Structure of Diborane". Journal of the American Chemical Society 59: 1096. doi:10.1021/ja01285a041.
  6. ^ Bauer S.H. (1942). "Structures and Physical Properties of the Hydrides of Boron and of their Derivatives". Chemical Reviews 31: 43-75. doi:10.1021/cr60098a001.
  7. ^ Schlesinger H.I., Burg A.B. (1942). "Recent Developments in the Chemistry of the Boron Hydrides". Chemical Reviews 31: 1-41. doi:10.1039/JR9430000250.
  8. ^ Longuet-Higgins, H.C., Bell, R.P. (1943). "The structure of the boron hydrides". Journal of the Chemical Society: 250-255.
  9. ^ Dilthey W. (1921). "". Zeitschriffte fuer Angewandte Chemie 34: 594.
  10. ^ Nekrassov BV (1940). "". J Gen Chem USSR 10: 1021.
  11. ^ Nekrassov BV (1940). "". J Gen Chem USSR 10: 1056.
  12. ^ Hedberg K, Schomaker V (1951). "A Reinvestigation of the Structures of Diborane and Ethane by Electron Diffraction". Journal of the American Chemical Society 73: 1482-1487. doi:10.1021/ja01148a022.

Further reading

H. C. Brown "Organic Synthesis via Boranes" John Wiley, New York, 1975. ISBN 0-471-11280-1.

This article is licensed under the GNU Free Documentation License. It uses material from the Wikipedia article "Diborane". A list of authors is available in Wikipedia.
Your browser is not current. Microsoft Internet Explorer 6.0 does not support some functions on Chemie.DE