View source for Diborane

{{See also |diborane(2)|diborane(4)}}
{{Chembox
| Watchedfields  = changed
| verifiedrevid  = 455184432
| ImageFile      = Diborane-2D-dimensions.svg
| ImageClass = skin-invert-image
| ImageFile_Ref  = {{chemboximage|correct|??}}
| ImageSize      = 244
| ImageName      = Stereo skeletal formula of diborane with all explicit hydrogens added and assorted measurements
| ImageFile1     = Diborane-3D-balls-A.png
| ImageClass1 = bg-transparent
| ImageFile1_Ref = {{chemboximage|correct|??}}
| ImageSize1     = 121
| ImageName1     = Ball and stick model of diborane
| IUPACName      = Diborane(6)
| OtherNames     = 
| Section1       = {{Chembox Identifiers
| CASNo = 19287-45-7
| CASNo_Ref = {{cascite|correct|CAS}}
| ChemSpiderID = 17215804
| ChemSpiderID_Ref = {{chemspidercite|correct|chemspider}}
| EINECS = 242-940-6
| RTECS = HQ9275000
| ChEBI_Ref = {{ebicite|correct|EBI}}
| ChEBI = 33590
| UNNumber = 1911
| UNII = BS9K982N24
| PubChem = 12544638
| SMILES = [BH2]1[H][BH2][H]1
| StdInChI = 1S/B2H6/c1-3-2-4-1/h1-2H2
| StdInChI_Ref = {{stdinchicite|correct|chemspider}}
| InChI = 1/B2H6/c1-3-2-4-1/h1-2H2
| StdInChIKey = KLDBIFITUCWVCC-UHFFFAOYSA-N
| StdInChIKey_Ref = {{stdinchicite|correct|chemspider}}
| InChIKey = KLDBIFITUCWVCC-UHFFFAOYAF
}}
| Section2       = {{Chembox Properties
| H=6 | B=2
| Appearance = Colorless gas
| Odor = repulsive and sweet
| Density = 1.131 g/L<ref name=r2/>
| MeltingPtC = −164.85
| MeltingPt_ref=<ref name=r2>Haynes, p. 4.52.</ref>
| BoilingPtC = −92.49
| BoilingPt_ref=<ref name=r2/>
| Solubility = Reacts<ref name=PGCH/>
| SolubleOther = [[Diglyme]] and [[Diethyl Ether]],<ref>Yerazunis, S., et al. "Solubility of Diborane in the Dimethyl Ether and Diethylene Glycol, in Mixtures of Sodium Borohydride and Dimethyl Ether of Diethylene Glycol, and in Ditertiary Butyl Sulfide". ''Journal of Chemical & Engineering Data'', vol. 7, no. 3, July 1962, pp. 337–39, {{doi|10.1021/je60014a004}}.</ref>
| Solvent = other solvents
| VaporPressure = 39.5 atm (16.6&nbsp;°C)<ref name=PGCH/>
}}
| Section3       = {{Chembox Structure
| MolShape = ''see text''
| Coordination = [[Tetrahedron|Tetrahedral]] (for boron)
| Dipole = 0 [[Debye|D]]
}}
| Section4       = {{Chembox Thermochemistry
| DeltaHf = 36.4 kJ/mol<ref name=r1>Haynes, p. 5.6.</ref>
| DeltaHc = 
| Entropy = 232.1{{nbsp}}J/(mol·K)<ref name=r1/>
| HeatCapacity = 56.7{{nbsp}}J/(mol·K)<ref name=r1/>
}}
| Section5       = {{Chembox Hazards
| NFPA-H = 4
| NFPA-F = 4
| NFPA-R = 3
| NFPA-S = W
| NFPA_ref = <ref name=ICSC>{{cite web |title=DIBORANE – CAMEO Chemicals - Chemical Datasheet - Database of Hazardous Materials – NOAA |url=https://cameochemicals.noaa.gov/chemical/4943|access-date=2022-10-26}}</ref>
| AutoignitionPtC = 38
| AutoignitionPt_notes = 
| PEL = TWA 0.1 ppm (0.1&nbsp;mg/m<sup>3</sup>)<ref name=PGCH>{{PGCH|0183}}</ref>
| ExploLimits = 0.8–88%<ref name=PGCH/>
| IDLH = 15 ppm<ref name=PGCH/>
| REL = TWA 0.1 ppm (0.1&nbsp;mg/m<sup>3</sup>)<ref name=PGCH/>
| MainHazards = toxic, highly flammable, reacts with water
| LC50 = 40 ppm (rat, 4&nbsp;[[hour|h]])<br/>29 ppm (mouse, 4&nbsp;h)<br/>40–80 ppm (rat, 4&nbsp;h)<br/>159–181 ppm (rat, 15&nbsp;[[minute|min]])<ref name=IDLH>{{IDLH|19287457|Diborane}}</ref>
| LCLo = 125 ppm (dog, 2&nbsp;h)<br/>50 ppm (hamster, 8&nbsp;h)<ref name=IDLH/>
| GHSPictograms = {{GHS02}}{{GHS06}}{{GHS08}}
| GHSSignalWord = Danger
| HPhrases = {{H-phrases|220|314|330|370|372}}
| PPhrases = {{P-phrases|210|260|264|270|271|280|284|301+330+331|303+361+353|304+340|305+351+338|307+311|310|314|320|321|363|377|381|403|403+233|405|410+403|501}}
| ExternalSDS = [https://www.inchem.org/documents/icsc/icsc/eics0432.htm ICSC 0432]
}}
| Section6       = {{Chembox Related
| OtherFunction_label = boron compounds
| OtherFunction = [[Decaborane]]<br />{{chem2|BF3|link=Boron trifluoride}}
}}
}}

'''Diborane(6)''', commonly known as '''diborane''', is the [[inorganic compound]] with the formula {{chem2|B2H6}}. It is a highly  [[toxic]], colorless, and [[pyrophoric]] gas with a repulsively sweet odor. Given its simple formula, diborane is a fundamental [[boron]] compound. It has attracted wide attention for its unique electronic structure. Several of its derivatives are useful [[Reagent|reagents]].

==Structure and bonding==
[[Image:Diborane 02.svg|thumb|left|Bonding diagram of diborane ({{chem2|B2H6}}) showing with curved lines a pair of [[three-center two-electron bond]]s, each of which consists of a pair of electrons bonding three atoms; two boron atoms and a hydrogen atom in the middle]]
The structure of diborane has [[Molecular symmetry#Common point groups|D<sub>2h</sub>]] symmetry. Four hydrides are terminal, while two bridge between the boron centers. The lengths of the B–H<sub>bridge</sub> bonds and the B–H<sub>terminal</sub> bonds are 1.33 and 1.19&nbsp;Å respectively. This difference in bond lengths reflects the difference in their strengths, the B–H<sub>bridge</sub> bonds being relatively weaker. The weakness of the B–H<sub>bridge</sub> compared to B–H<sub>terminal</sub> bonds is indicated by their vibrational signatures in the [[infrared spectrum]], being ≈2100 and 2500&nbsp;cm<sup>−1</sup> respectively.<ref>{{cite book |author=Cooper, C. B., III |author2=Shriver, D. F. |author3=Onaka, S. |chapter=Ch. 17. Vibrational spectroscopy of hydride-bridged transition metal compounds |series=Advances in Chemistry |year=1978 |title=Transition Metal Hydrides |pages=[https://archive.org/details/transitionmetalh0000unse_j8w9/page/232 232–247] |doi=10.1021/ba-1978-0167.ch017 |volume=167 |isbn=9780841203907 |chapter-url=https://archive.org/details/transitionmetalh0000unse_j8w9/page/232}}</ref>

The model determined by [[molecular orbital theory]] describes the bonds between boron and the terminal hydrogen atoms as conventional 2-center 2-electron [[covalent bond]]s. The bonding between the boron atoms and the bridging hydrogen atoms is, however, different from that in molecules such as hydrocarbons. Each boron uses two electrons in bonding to the terminal hydrogen atoms and has one [[valence electron]] remaining for additional bonding. The bridging hydrogen atoms provide one electron each. The {{chem2|B2H2}} ring is held together by four electrons forming two [[three-center two-electron bond|3-center 2-electron bonds]]. This type of bond is sometimes called a "[[banana bond]]".

{{chem2|B2H6}} is [[isoelectronic]] with {{chem2|C2H6(2+)}}, which would arise from the [[protonation|diprotonation]] of the planar molecule [[ethylene]].<ref>{{ cite journal | journal = [[Journal of Physical Chemistry A]] | year = 2005 | volume = 109 | issue = 5 | pages = 798–801 | author1 = Rasul, G. | author2 = Prakash, G. K. S. | author3 = Olah, G. A. | author-link3 = George Andrew Olah | title = Comparative ''ab initio'' Study of the Structures and Stabilities of the Ethane Dication C<sub>2</sub>H<sub>6</sub><sup>2+</sup> and Its Silicon Analogues Si<sub>2</sub>H<sub>6</sub><sup>2+</sup> and CSiH<sub>6</sub><sup>2+</sup> | pmid = 16838949 | doi = 10.1021/jp0404652 | bibcode = 2005JPCA..109..798R }}</ref> Diborane is one of many compounds with such unusual bonding.<ref>{{ cite journal | author = Laszlo, P. | title = A Diborane Story | journal= Angewandte Chemie International Edition | year = 2000  | volume = 39 | issue = 12 | pages = 2071–2072 | doi = 10.1002/1521-3773(20000616)39:12<2071::AID-ANIE2071>3.0.CO;2-C | pmid = 10941018 }}</ref>

Of the other elements in the [[boron group]], gallium is known to form a similar compound [[digallane]], {{chem2|Ga2H6}}. Aluminium forms a polymeric hydride, ({{chem2|(AlH3)_{''n''}|link=aluminium hydride}}; although unstable, {{chem2|Al2H6}} has been isolated in solid hydrogen and is isostructural with diborane.<ref>{{ cite journal | title = The Infrared Spectrum of Al<sub>2</sub>H<sub>6</sub> in Solid Hydrogen |author1=Andrews, L. |author2=Wang, X. | journal = Science | year = 2003 | volume = 299 | issue = 5615 | pages = 2049–2052 | doi = 10.1126/science.1082456 | pmid = 12663923 | bibcode = 2003Sci...299.2049A |s2cid=45856199 }}</ref>

==Production and synthesis==
Extensive studies of diborane have led to the development of multiple synthesis routes. Most preparations entail reactions of [[hydride]] donors with boron [[halide]]s or [[alkoxide]]s. The industrial synthesis of diborane involves the reduction of {{chem2|BF3|link=Boron trifluoride}} by [[sodium hydride]] (NaH), [[lithium hydride]] (LiH) or [[lithium aluminium hydride]] ({{chem2|LiAlH4}}):<ref name="Georg">{{cite book |last=Brauer |first=Georg |title=Handbook of Preparative Inorganic Chemistry |volume=1 |edition=2nd |date=1963 |publisher=Academic Press |location=New York |isbn=978-0121266011 |page=773 |url=https://books.google.com/books?id=TLYatwAACAAJ&q=Handbook+of+Preparative+Inorganic+Chemistry}}</ref>
:{{chem2|8 BF3 + 6 LiH -> B2H6  + 6 LiBF4}}

[[Lithium hydride]] used for this purpose must be very finely powdered to avoid the formation of a [[passivation (chem)|passivating]] [[lithium tetrafluoroborate]] layer on the reactant.  Alternatively, a small amount of diborane product can be added to form [[lithium borohydride]], which will react with the BF<sub>3</sub> to produce more diborane, making the reaction [[autocatalytic]].<ref>{{cite book |last=DeQuasie |first=Andrew |title=The Green Flame |at=Chapter 3}}  Excerpted in 2&nbsp;parts and archived at the [[WayBack Machine]]: [https://web.archive.org/web/20020414035653/http://www.dequasiebooks.com/green.html], [https://web.archive.org/web/20020604175527/http://www.dequasiebooks.com:80/green.html]</ref>

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:
:{{chem2|4 BCl3 + 3 LiAlH4 -> 2 B2H6 + 3 LiAlCl4}}
:{{chem2|4 BF3 + 3 NaBH4 -> 2 B2H6 + 3 NaBF4}}

When heated with {{chem2|NaBH4}}, [[tin(II) chloride]] is reduced to elemental tin, forming diborane in the process:
:{{chem2|SnCl2 + 2NaBH4 -> 2NaCl + Sn + B2H6 + H2}}

Older methods entail the direct reaction of borohydride salts with a [[Oxidizing acid|non-oxidizing acid]], such as [[phosphoric acid]] or dilute [[sulfuric acid]]:
:{{chem2|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:<ref>{{Kirk-Othmer|title=Boron hydrides, heteroboranes, and their <em>metalla</em> derivatives|first=David&nbsp;M.|last=Schuster|doi=10.1002/0471238961.0215181519030821.a01.pub2|p=184}}</ref>
:{{chem2|2 NaBH4 + I2 -> 2 NaI + B2H6 + H2}}

Another small-scale synthesis uses [[potassium borohydride]] and phosphoric acid as starting materials.<ref>{{ Cite book |author1=Norman, A. D. |author2=Jolly, W. L. |author3=Saturnino, D. |author4=Shore, S. G.|chapter=Diborane | title = Inorganic Syntheses | year = 1968 | volume = 11 |pages = 15–19| doi =10.1002/9780470132425.ch4 |isbn=9780470132425 }}</ref>

==Reactions==
[[File:BH3SMe2improve.svg|thumb|left|122px|[[Borane dimethylsulfide]] generally functions equivalently to diborane and is easier to use.<ref>{{cite journal | doi = 10.1080/00304948109356130 | title = Utility and Applications of Borane Dimethylsulfide in Organic Synthesis. A Review | year = 1981 | last1 = Hutchins | first1 = Robert O. | last2 = Cistone | first2 = Frank | journal = [[Organic Preparations and Procedures International]] | volume = 13 | issue = 3–4 | pages = 225}}</ref>]]
Diborane is a highly reactive and versatile reagent.<ref name=Lane>{{ cite journal | author = Mikhailov, B. M. | title = The Chemistry of Diborane | journal = Russian Chemical Reviews | year = 1962 | volume = 31 | issue = 4 | pages = 207–224 | doi = 10.1070/RC1962v031n04ABEH001281 | bibcode = 1962RuCRv..31..207M | s2cid = 250909492 }}</ref>

===Air, water, oxygen===
As a [[pyrophoric]] substance, diborane reacts [[Exothermic reaction|exothermically]] with [[oxygen]] to form [[boron trioxide]] and water:
:{{chem2|2 B2H6 + 6 O2 -> 2 B2O3 + 6 H2O}} ; [[Enthalpy|Δ''H''<sub>r</sub>]] = −2035 k[[Joule|J]]/[[Mole (unit)|mol]] = −73.47 kJ/[[gram|g]]

Diborane reacts violently with water to form hydrogen and [[boric acid]]:
:{{chem2|B2H6 + 6 H2O -> 2 B(OH)3 + 6 H2}} ; Δ''H''<sub>r</sub> = −466 kJ/mol = −16.82 kJ/[[gram|g]])

Diborane also reacts with alcohols similarly. For example, the reaction with methanol gives hydrogen and [[trimethylborate]]:<ref name="InorgChem">{{ cite book | author1 = Housecroft, C. E. | author2 = Sharpe, A. G. | chapter = Chapter 13: The Group 13 Elements | title = Inorganic Chemistry | url = https://archive.org/details/inorganicchemist00hous_159 | url-access = limited | edition = 3rd | year = 2008 | page = [https://archive.org/details/inorganicchemist00hous_159/page/n374 336] | publisher = Pearson | isbn = 978-0-13-175553-6 }}</ref>
:{{chem2|B2H6 + 6 MeOH -> 2 B(OMe)3 + 6 H2}}

===Lewis acidity===
One dominating reaction pattern involves formation of adducts with [[Lewis bases]]. Often such initial adducts proceed rapidly to give other products. For example, borane-tetrahydrofuran, which often behaves equivalently to diborane, degrades to borate esters. Its adduct with dimethyl sulfide is an important reagent in [[organic synthesis]]. With [[ammonia]] diborane forms the diammoniate of diborane, DADB with small quantities of [[ammonia borane]] as byproduct. The ratio depends on the conditions.

===Hydroboration===
In the [[hydroboration]] reaction, diborane also reacts readily with [[alkene]]s to form tri[[alkylborane]]s.  This reaction pattern is rather general and the resulting alkyl borates can be readily derivatized, e.g. to alcohols.  Although early work on hydroboration relied on diborane, it has been replaced by borane dimethylsulfide, which is more safely handled.

===Other===
Pyrolysis of diborane gives hydrogen and diverse boron hydride clusters.  For example, [[pentaborane]] was first prepared by [[pyrolysis]] of diborane at about 200&nbsp;°C.<ref>{{ cite book | last = Stock | first = A. |year = 1933 | title = The Hydrides of Boron and Silicon | publisher = Cornell University Press | location = New York | isbn = 0-8014-0412-6 }}</ref><ref>{{ Cite book |author1=Miller, V. R. |author2=Ryschkewitsch, G. E. |chapter=Pentaborane(9) (B <sub>5</sub> H <sub>9</sub> ) | title = Inorganic Syntheses | year = 1974 | volume = 15 | pages = 118–122 | doi = 10.1002/9780470132463.ch26 |isbn=9780470132463 }}</ref> Although this pyrolysis route is rarely employed, it ushered in a large research theme of [[borane cluster]] chemistry.

Treating diborane with sodium [[Amalgam (chemistry)|amalgam]] gives {{chem2|NaBH4}} and {{chem2|Na[B3H8]}}<ref name="InorgChem" />
When diborane is treated with [[lithium hydride]] in [[diethyl ether]], lithium borohydride is formed:<ref name="InorgChem" />
:{{chem2|B2H6 + 2 LiH -> 2 LiBH4}}
Diborane reacts with anhydrous [[hydrogen chloride]] or [[hydrogen bromide]] gas to give a boron halohydride:<ref name="InorgChem" />
:{{chem2|B2H6 + HCl -> B2H5Cl + H2}}
Treating diborane with [[carbon monoxide]] at 470&nbsp;K and 20&nbsp;bar gives {{chem2|H3BCO|link=Borane carbonyl}}.<ref name="InorgChem" />

==Reagent in organic synthesis==
Diborane and its variants are central [[organic synthesis]] reagents for [[hydroboration]]. Alkenes add across the B–H bonds to give trialkylboranes, which can be further elaborated.<ref>{{cite journal |doi=10.1021/cr60304a005|title=Reduction of organic compounds with diborane|year=1976|last1=Lane|first1=Clinton F.|journal=Chemical Reviews|volume=76|issue=6|pages=773–799}}</ref> Diborane is used as a [[reducing agent]] roughly complementary to the reactivity of [[lithium aluminium hydride]]. The compound readily reduces [[carboxylic acid]]s to the corresponding [[Alcohol (chemistry)|alcohol]]s, whereas [[ketone]]s 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, [[Alfred Stock]], the major pioneer in the chemistry of boron hydrides, 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.<ref>{{ cite book | author = Stock, A. | year = 1933 | title = The Hydrides of Boron and Silicon | publisher = Cornell University Press | location = New York }}</ref> [[Electron diffraction]] measurements by S.&nbsp;H. Bauer initially appeared to support his proposed structure.<ref>{{ cite journal | author = Bauer, S. H. | title = The Structure of Diborane | journal = [[Journal of the American Chemical Society]] | year = 1937 | volume = 59 | issue = 6 | pages = 1096–1103 | doi = 10.1021/ja01285a041 | bibcode = 1937JAChS..59.1096B }}</ref><ref>{{ cite journal | author = Bauer, S. H. | title = Structures and Physical Properties of the Hydrides of Boron and of their Derivatives | journal = Chemical Reviews | year = 1942 | volume = 31 | issue = 1 | pages = 43–75 | doi = 10.1021/cr60098a002 }}</ref>

Because of a personal communication with [[Linus Pauling|L.&nbsp;Pauling]] (who supported the ethane-like structure), [[Hermann Irving Schlesinger|H.&nbsp;I. Schlessinger]] and A.&nbsp;B. Burg did not specifically discuss [[three-center two-electron bond|3-center 2-electron bonding]] in their then classic review in the early 1940s.<ref>{{ cite journal |author1=Schlesinger, H. I. |author2=Burg, A. B. | title = Recent Developments in the Chemistry of the Boron Hydrides | journal = Chemical Reviews | year = 1942 | volume = 31 | issue = 1 | pages = 1–41 | doi = 10.1021/cr60098a001 }}</ref> The review does, however, discuss the bridged D<sub>2h</sub> structure in some depth: "It is to be recognized that this formulation easily accounts for many of the chemical properties of diborane..."

In 1943, [[H.&nbsp;Christopher Longuet-Higgins]], while still an undergraduate at Oxford, was the first to explain the structure and bonding of the boron hydrides. The article reporting the work, written with his tutor [[Ronnie Bell (chemist)|R.&nbsp;P. Bell]],<ref>{{Cite journal | last1 = Longuet-Higgins | first1 = H. C. | last2 = Bell| first2 = R. P. | author-link = H. Christopher Longuet-Higgins | title = 64. The Structure of the Boron Hydrides | journal = Journal of the Chemical Society (Resumed) | year = 1943 | volume = 1943 | pages = 250–255 | doi = 10.1039/JR9430000250 }}</ref> also reviews the history of the subject beginning with the work of Dilthey.<ref>{{ cite journal | author = Dilthey, W. | title = Über die Konstitution des Wassers | journal = Angewandte Chemie | year = 1921 | volume = 34 | issue = 95 | pages = 596 | doi = 10.1002/ange.19210349509 | doi-access = free }}</ref> Shortly afterwards, the theoretical work of Longuet-Higgins was confirmed in an infrared study of diborane by Price.<ref>{{cite journal | last1 = Price | first1 = W. C. | year = 1948 | title = The absorption spectrum of diborane | journal = J. Chem. Phys. | volume = 16 | issue = 9| page = 894 | doi=10.1063/1.1747028 | bibcode = 1948JChPh..16..894P }}</ref>
The structure was re-confirmed by electron-diffraction measurement in 1951 by K.&nbsp;Hedberg and V.&nbsp;Schomaker, with the confirmation of the structure shown in the schemes on this page.<ref>{{ cite journal |author1=Hedberg, K. |author2=Schomaker, V. | title = A Reinvestigation of the Structures of Diborane and Ethane by Electron Diffraction | journal = [[Journal of the American Chemical Society]] | year = 1951 | volume = 73 | issue = 4 | pages = 1482–1487 | doi = 10.1021/ja01148a022 |bibcode=1951JAChS..73.1482H }}</ref>

[[William Nunn Lipscomb]]&nbsp;Jr. further confirmed the molecular structure of boranes using [[X-ray crystallography]] in the 1950s and developed theories to explain their bonding. Later, he applied the same methods to related problems, including the structure of carboranes, on which he directed the research of future 1981 [[Nobel Prize]] winner [[Roald Hoffmann]]. The 1976 [[Nobel Prize in Chemistry]] was awarded to Lipscomb "for his studies on the structure of boranes illuminating problems of chemical bonding".<ref>{{cite web |url=http://nobelprize.org/nobel_prizes/chemistry/laureates/1976/ |title=The Nobel Prize in Chemistry 1976 |publisher=Nobelprize.org |access-date=2012-02-01}}</ref>

Traditionally, diborane has often been described as [[Electron deficiency|electron-deficient]], because the 12 valence electrons can only form 6 conventional 2-centre 2-electron bonds, which are insufficient to join all 8 atoms.<ref>{{cite journal |last1=Longuet-Higgins |first1=H. C. |title=The structures of electron-deficient molecules |journal=Quarterly Reviews, Chemical Society |date=1957 |volume=11 |issue=2 |pages=121–133 |doi=10.1039/qr9571100121 |url=https://pubs.rsc.org/en/content/articlelanding/1957/qr/qr9571100121 |access-date=15 July 2020|url-access=subscription }}</ref><ref>{{cite book |last1=Murrell |first1=J. N. |author-link=John Murrell (chemist) |last2=Kettle |first2=S. F. A. |last3=Tedder |first3=J. M. |title=Valence theory |date=1965 |publisher=John Wiley and Sons |page=243}}</ref> However, the more correct description using 3-centre bonds shows that diborane is really electron-precise, since there are just enough valence electrons to fill the 6 [[bonding molecular orbital]]s.<ref>{{cite web |last1=Lipscomb |first1=William N. |title=The Boranes and their relatives (Nobel lecture) |url=https://www.nobelprize.org/uploads/2018/06/lipscomb-lecture.pdf |website=nobelprize.org |publisher=Nobel Foundation |access-date=16 July 2020 |pages=224–245 |date=11 December 1976 |quote=One of the simple consequences of these studies was that electron deficient molecules, defined as having more valence orbitals than electrons, are not really electron deficient.}}</ref> Nevertheless, some leading textbooks still use the term "electron-deficient".<ref>{{cite book |last1=Housecroft |first1=Catherine E. |last2=Sharpe |first2=Alan G. |title=Inorganic Chemistry |date=2005 |publisher=Pearson Prentice-Hall |isbn=0130-39913-2 |page=326 |edition=2nd |quote=An electron-deficient species possesses fewer valence electrons than are required for a localized bonding scheme.}}</ref>

==Other uses==
Because of the exothermicity of its reaction with oxygen, diborane has been tested as a [[rocket propellant]].<ref>{{cite web |last1=Bilstein |first1=Roger |title=Stages to Saturn |url=http://www.history.nasa.gov/SP-4206/sp4206.htm |publisher=NASA Public Affairs Office |access-date=14 November 2015 |location=chapter 5 |page=133 |archive-date=25 December 2017 |archive-url=https://web.archive.org/web/20171225230901/https://history.nasa.gov/SP-4206/sp4206.htm |url-status=dead }}</ref> Complete combustion is strongly exothermic. However, combustion is not complete in the rocket engine, as some [[boron monoxide]], {{chem2|B2O}}, is produced. This conversion mirrors the incomplete combustion of [[hydrocarbon]]s, to produce [[carbon monoxide]] (CO). Diborane also proved difficult to handle.<ref>{{cite book |title=A preliminary experimental and analytical evaluation of diborane as a ram-jet fuel |url=https://ntrs.nasa.gov/archive/nasa/casi.ntrs.nasa.gov/19930086375.pdf |author1=Gammon, Benson E. |author2=Genco, Russell S. |author3=Gerstein, Melvin |publisher=National Advisory Committee for Aeronautics |year=1950}}</ref><ref>{{cite book |publisher=National Advisory Committee for Aeronautics |url=http://naca.central.cranfield.ac.uk/reports/1958/naca-report-1362.pdf |title=Theoretical Combustion Performance of Several High-Energy Fuels for Ramjet Engines |author1=Tower, Leonard K. |author2=Breitwieser, Roland |author3=Gammon, Benson E. |year=1958 }}</ref><ref>{{cite web |url=https://history.nasa.gov/SP-4404/ch5-1.htm |title=LIQUID HYDROGEN AS A PROPULSION FUEL, 1945–1959. Part II: 1950–1957. Chapter 5. NACA Research on High-Energy Propellants |website=history.nasa.gov}}</ref>

Diborane has been investigated as a precursor to metal boride films<ref>{{Ullmann|doi=10.1002/14356007.a04_309|chapter=Boron Compounds|year=2000|last1=Brotherton|first1=Robert J.|last2=Weber|first2=C. Joseph|last3=Guibert|first3=Clarence R.|last4=Little|first4=John L.|isbn=3527306730}}</ref> and for the p-doping of silicon semiconductors.<ref>{{cite journal |doi=10.1149/1.1864452|title=A Kinetic Model for Boron and Phosphorus Doping in Silicon Epitaxy by CVD|year=2005|last1=Mehta|first1=Bhavesh|last2=Tao|first2=Meng|journal=Journal of the Electrochemical Society|volume=152|issue=4|pages=G309|bibcode=2005JElS..152G.309M|doi-access=free}}</ref>

==Safety==
Diborane is a pyrophoric gas. Commercially available [[adduct]]s are typically used instead, at least for applications in organic chemistry.  These adducts include [[borane-tetrahydrofuran]] (borane-THF) and [[borane-dimethylsulfide]].<ref name=Lane/>
The toxic effects of diborane are mitigated because the compound is so unstable in air.  The toxicity toward laboratory rats has been investigated.<ref>{{cite journal |doi=10.1006/taap.1996.0100|title=Evaluation of the Subacute Pulmonary and Testicular Inhalation Toxicity of Diborane in Rats|year=1996|last1=Nomiyama|first1=Tetsuo|last2=Omae|first2=Kazuyuki|last3=Ishizuka|first3=Chizuru|last4=Hosoda|first4=Kanae|last5=Yamano|first5=Yuko|last6=Nakashima|first6=Hiroshi|last7=Uemura|first7=Takamoto|last8=Sakurai|first8=Haruhiko|journal=Toxicology and Applied Pharmacology|volume=138|issue=1|pages=77–83|pmid=8658516|bibcode=1996ToxAP.138...77N }}</ref>

==References==
{{Reflist}}

==Cited sources==
*{{cite book | editor= Haynes, William M. | year = 2011 | title = CRC Handbook of Chemistry and Physics | edition = 92nd | publisher = [[CRC Press]] | isbn = 978-1439855119| title-link = CRC Handbook of Chemistry and Physics }}
*Yerazunis, S., et al. “Solubility of Diborane in the Dimethyl Ether and Diethylene Glycol, in Mixtures of Sodium Borohydride and Dimethyl Ether of Diethylene Glycol, and in Ditertiary Butyl Sulfide.” ''Journal of Chemical & Engineering Data'', vol. 7, no. 3, July 1962, pp. 337–39, doi:10.1021/je60014a004.

==Further reading==
*{{Cite book |last1=Constantine |first1=M. T. |last2=Weber |first2=J. Q. |last3=Youel |first3=K. J. |date=1971 |url=https://ntrs.nasa.gov/citations/19700078681 |title=Diborane Handbook |publisher=Rocketdyne |id=NASA-CR-111126}}
*{{Cite encyclopedia |title=Boranes  |chapter=Diborane |last=Schmidt |first=Eckart W. |date=2022 |encyclopedia=Encyclopedia of Liquid Fuels |publisher=De Gruyter |pages=978–1013 |doi=10.1515/9783110750287-012 |isbn=978-3-11-075028-7}}

==External links==
* [https://www.inchem.org/documents/icsc/icsc/eics0432.htm International Chemical Safety Card 0432]
* [https://web.archive.org/web/20060209040519/http://www.npi.gov.au/database/substance-info/profiles/15.html National Pollutant Inventory – Boron and compounds]
* [https://www.cdc.gov/niosh/npg/npgd0183.html NIOSH Pocket Guide to Chemical Hazards]
* [http://www.epa.gov/oppt/aegl/results65.htm U.S. EPA Acute Exposure Guideline Levels]{{dead link|date=July 2025|bot=medic}}{{cbignore|bot=medic}}

{{Boron compounds}}
{{Hydrides by group}}
{{Authority control}}

[[Category:Boranes]]
[[Category:Rocket fuels]]
[[Category:Reducing agents]]
[[Category:Industrial gases]]
[[Category:Sweet-smelling chemicals]]
[[Category:Pyrophoric materials]]