Diborano: diferenças entre revisões

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| verifiedrevid = 443635822
[[Ficheiro:Diborane 02.svg|thumb|150px|left]]
A [[fórmula estrutural]] do diborano (B<sub>2</sub>H<sub>6</sub>) mostra com linhas curvas um par de [[ligação três centros dois elétrons|ligações três centros dois elétrons]], cada uma das quais consiste de um par de elétrons ligando três átomos, dois átomos de boro e um átomo de hidrogênio no meio.
Diborane adopts a D<sub>2h</sub> 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 bond]]s. 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 B<sub>2</sub>H<sub>2</sub> ring is held together with four electrons, an example of [[3-center-2-electron bond]]ing. This type of bond is sometimes called a 'banana bond'. The lengths of the B-H<sub>bridge</sub> bonds and the B-H<sub>terminal</sub> 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<sub>bridge</sub> bonds being relatively weaker. The structure is [[isoelectronic]] with C<sub>2</sub>H<sub>6</sub><sup>2+</sup>, which would arise from the diprotonation of the planar molecule [[ethene]].<ref>{{ cite journal | journal = [[Journal of Physical Chemistry A|J. Phys. Chem. A]] | year = 2005 | volume = 109 | issue = 5 | pages = 798–801 | author = G. Rasul, G. K. S. Prakash, [[George Andrew Olah|G. A. Olah]] | doi = 10.1021/jp0404652 | 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 }}</ref> Diborane is one of many compounds with such unusual bonding.<ref>{{cite journal | author= Laslo P | title= A Diborane Story | journal= Angewandte Chemie International Edition | year= 2000 | volume= 39 | pages= 2071–2072| doi= 10.1002/1521-3773(20000616)39:12<2071::AID-ANIE2071>3.0.CO;2-C | pmid=10941018 | issue= 12}} [http://www3.interscience.wiley.com/cgi-bin/abstract/72507084/ABSTRACT abstract]</ref>
Of the other elements in Group 13, gallium is known to form a similar compound, [[digallane]], Ga<sub>2</sub>H<sub>6</sub>. Aluminium forms a polymeric hydride, [[aluminium hydride|(AlH<sub>3</sub>)<sub>''n''</sub>]], although unstable Al<sub>2</sub>H<sub>6</sub> 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
| author = Andrews, Lester; Wang, Xuefeng
| journal = Science
| year = 2003
| volume = 299
| issue = 5615
| pages = 2049–2052
| doi = 10.1126/science.1082456
| pmid = 12663923
}}</ref> No hydrides of indium and thallium have yet been found.<ref>{{cite journal
| last = Downs | first = Anthony J. | authorlink =
| coauthors = Colin R. Pulham
| title = The hydrides of aluminium, gallium, indium and thallium: A re-evaluation
| journal = Chemical Society Reviews | volume = 23 | issue = 3 | pages = 175–184
| publisher = [[Royal Society of Chemistry]] | location = Cambridge | year = 1994 | accessdate =
| doi = 10.1039/cs9942300175 }}</ref>
== Production and synthesis ==
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 BF<sub>3</sub> by [[sodium hydride]]:
:2 BF<sub>3</sub> + 6 NaH → B<sub>2</sub>H<sub>6</sub> + 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 BCl<sub>3</sub> + 3 LiAlH<sub>4</sub> → 2 B<sub>2</sub>H<sub>6</sub> + 3 LiAlCl<sub>4</sub>
:4 BF<sub>3</sub> + 3 NaBH<sub>4</sub> → 2 B<sub>2</sub>H<sub>6</sub> + 3 NaBF<sub>4</sub>
Older methods entail the direct reaction of borohydride salts with a [[Oxidizing acid|non-oxidizing acid]], such as [[phosphoric acid]] or dilute [[sulfuric acid]].
:2 BH<sub>4</sub><sup>−</sup> + 2 H<sup>+</sup> → 2 H<sub>2</sub> + B<sub>2</sub>H<sub>6</sub>
Similarly, oxidation of borohydride salts has been demonstrated and remains convenient for small scale preparations. For example, using [[iodine]] as oxidizer:
:2 {{chem|NaBH|4}} + {{chem|I|2}} → 2 NaI + {{chem|B|2|H|6}} + {{chem|H|2}}
== Reactions ==
Diborane is a highly reactive and versatile reagent that has a large number of applications.<ref>{{cite journal| title= The Chemistry of Diborane|author= Mikhailov BM|journal= Russian Chemical Review|issue= 31 |pages= 207–224|year= 1962 | volume= 31| doi= 10.1070/RC1962v031n04ABEH001281}}</ref> 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 [[ammonia borane]] or the diammoniate of diborane, DADB, depending on the conditions used. Diborane also reacts readily with [[alkyne]]s to form substituted [[alkene]] products which will readily undergo further [[addition reaction]]s.
Diborane reacts with water to form hydrogen and [[boric acid]].
The compound forms complexes with [[Lewis base]]s. Notable are the complexes with [[tetrahydrofuran|THF]] and [[dimethyl sulfide]], both liquid compounds that 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 Me<sub>2</sub>S). The THF adduct is usually prepared as a 1:5 solution in [[tetrahydrofuran|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)BH<sub>3</sub> + 3 CH<sub>2</sub>=CHR → B(CH<sub>2</sub>CH<sub>2</sub>R)<sub>3</sub> + 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 [HBR<sub>2</sub>]<sub>2</sub>, 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 acid]]s to the corresponding [[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, 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.<ref>{{cite book |author= Stock A.|year = 1933|title = The Hydrides of Boron and Silicon| edition = |publisher = Cornell University Press|location = New York| id = | page= }}</ref> [[Electron diffraction]] measurements by S. 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 | pages=1096| volume=59| doi=10.1021/ja01285a041| issue=6 }}</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 | pages=43–75| volume=31| doi=10.1021/cr60098a001| last2=Burg| first2=Anton B. }}</ref>
Because of a personal communication with [[Linus Pauling|L. Pauling]] (who supported the ethane-like structure), [[Hermann Irving Schlesinger|H. I. Schlessinger]] did not specifically discuss [[3-center-2-electron bond]]ing in his then classic review in the early 1940s.<ref>{{cite journal| author=Schlesinger H.I., Burg A.B. | title=Recent Developments in the Chemistry of the Boron Hydrides| journal=Chemical Reviews| year=1942 | pages=1–41| volume=31| doi=10.1039/JR9430000250}}</ref> The review does, however, discuss the C<sub>2v</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 an undergraduate student at [[Balliol College, Oxford]], [[H. Christopher Longuet-Higgins]], published the currently accepted structure together with [[Ronnie Bell|R. P. Bell]].<ref>{{cite doi|10.1039/JR9430000250 }}</ref> This structure had already been described in 1921.<ref>{{cite journal| author=Dilthey W. | title=Drehbrenner mit fester Gaszuführung| journal=Zeitschriffte fuer Angewandte Chemie| year= 1921| pages=594| volume=34 | doi=10.1002/ange.19210349504| issue=95}}</ref><ref>{{cite journal| author=Nekrassov BV| title=| journal=J Gen Chem USSR| year=1940 | pages=1021| volume=10}}</ref><ref>{{cite journal| author=Nekrassov BV| title=| journal=J Gen Chem USSR| year=1940 | pages=1056| volume=10}}</ref> 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.<ref>{{cite journal| author=Hedberg K, Schomaker V | title=A Reinvestigation of the Structures of Diborane and Ethane by Electron Diffraction| journal=[[Journal of the American Chemical Society]]| year=1951 | pages=1482–1487| volume=73|doi= 10.1021/ja01148a022| issue=4}}</ref>
[[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 is used in [[rocket propellant]]s, as a [[rubber]] [[vulcanization|vulcaniser]], as a [[catalyst]] for [[hydrocarbon]] [[polymerisation]], as a flame-speed accelerator, and as a [[Doping (Semiconductors)|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 [[tokamak]]s to reduce the amount of heavy metal impurities in the 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|date=September 2009}}
People exposed for a long time to low amounts of diborane have experienced respiratory irritation, [[seizures]], fatigue, [[drowsiness]], confusion, and occasional transient tremors.-->