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Sodium borohydride
The structure of sodium borohydride
IUPAC name
Identifiers
CAS number 16940-66-2 Yes check.svgY
PubChem 22959485
UN number 1426
RTECS number ED3325000
Properties
Molecular formula NaBH4
Molar mass 37.83 g/mol
Appearance white crystals
hygroscopic
Density 1.0740 g/cm3
Melting point

400 °C[1]

Boiling point

500 °C (dec.)[1]

Solubility in water not soluble, reacts with water
Solubility soluble in liquid ammonia, amines, pyridine
Hazards
MSDS ICSC 1670
NFPA 704
NFPA 704.svg
1
2
2
W
Flash point 70 °C
Autoignition
temperature
ca. 220 °C
LD50 160 mg/kg
Related compounds
Other anions Sodium cyanoborohydride
Sodium hydride
Sodium borate
Borax
Other cations Lithium borohydride
Related compounds Lithium aluminium hydride
 Yes check.svgY (what is this?)  (verify)
Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa)
Infobox references

Sodium borohydride, also known as sodium tetrahydridoborate, is an inorganic compound with the formula NaBH4. This white solid, usually encountered as a powder, is a versatile reducing agent that finds wide application in chemistry, both in the laboratory and on a technical scale. Large amounts are used for bleaching wood pulp. The compound is insoluble in ether, and soluble in glyme solvents, methanol and water, but reacts with the latter two in the absence of base.[2]

The compound was discovered in the 1940s by H. I. Schlesinger, who led a team that developed metal borohydrides for wartime applications.[3] Their work was classified and published only in 1953.

Contents

Physical properties

Sodium borohydride is an odorless white to gray-white microcrystalline powder which often forms lumps. It is soluble in water, with which it reacts vigorously.

Structure

NaBH4 has three known polymorphs: α, β and γ. The stable phase at room temperature and pressure is α-NaBH4, which is cubic and adopts an NaCl-type structure, in the Fm3m space group. At a pressure of 6.3 GPa, the structure changes to the tetragonal β-NaBH4 (space group P421c) and at 8.9 GPa, the orthorhombic γ-NaBH4 (space group Pnma) becomes the most stable.[4][5][6]

Alpha-sodium-borohydride-xtal-2007-3D-balls.png
Beta-sodium-borohydride-xtal-2007-3D-balls.png
Gamma-sodium-borohydride-xtal-2007-3D-balls.png
α-NaBH4
β-NaBH4
γ-NaBH4

Synthesis and handling

Sodium borohydride is prepared by two routes of industrial significance. In one method, based on the original work of Schesinger, sodium hydride is treated with trimethylborate at 250-270 °C:

B(OCH3)3 + 4 NaH → NaBH4 + 3 NaOCH3

Alternatively, sodium borohydride is also produced by the action of NaH on powdered borosilicate glass. Millions of kilograms are produced annually, far exceeding the production levels of any other hydride reducing agent.[7][8]

NaBH4 can be recrystallized by dissolving in warm (50 °C) diglyme followed by cooling the solution.[9]

Reactivity

NaBH4 will reduce many organic carbonyls, depending on the precise conditions. Most typically, it is used in the laboratory for converting ketones and aldehydes to alcohols. It will reduce acyl chlorides and thiol esters. However, unlike the powerful reducing agent lithium aluminium hydride, NaBH4 typically will not reduce esters, amides, or carboxylic acids.[2]

Many other hydride reagents are more strongly reducing. These usually involve replacing hydride with alkyl groups, e.g., Lithium triethylborohydride) and L-Selectride (lithium tri-sec-butylborohydride, or replacing B with Al. Variations in the counterion also affect the reactivity of the borohydride.[10]

Oxidation of NaBH4 with iodine in tetrahydrofuran gives the BH3-THF complex, which can reduce carboxylic acids. Likewise, the NaBH4-MeOH system, formed by the addition of methanol to sodium borohydride in refluxing THF, reduces esters to the corresponding alcohols, for instance, benzyl benzoate to benzyl alcohol.[11]

Aqueous solutions of sodium borohydride are decomposed by catalytic amounts of cobalt((II) ions to yield sodium borate and hydrogen gas. Pellets of cobalt-doped sodium borohydride are commercially available for use in hydrogen generators, for applications where cylinders of hydrogen would be inconvenient.

BH4 is an excellent ligand for metal ions. Such borohydride complexes are often prepared by the action of NaBH4 (or the LiBH4) on the corresponding metal halide, e.g. Zr(BH4)4. One example is the titanocene derivative:[12]

2 (C5H5)2TiCl2 + 4 NaBH4 → 2 (C5H5)2TiBH4 + 4 NaCl + B2H6 + H2
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Combustion

Sodium borohydride is less flammable and less volatile than gasoline, but more corrosive. It is relatively environmentally friendly because of the low toxicity of borates. The hydrogen is generated for a fuel cell by catalytic decomposition of the aqueous borohydride solution:

NaBH4 + 2H2O → NaBO2 + 4H2 + heat

Applications

The principle application of sodium borohydride is the production of sodium dithionite, which is used as a bleaching agent in the for wood pulp. Sulfur dioxide reacts with the borohydride. In a related process, sodium dithionite is used in the dying industry.

Production of pharmaceuticals

Sodium borohydride reduces aldehydes and ketones into alcohols. This reaction is used in the production of various antibiotics including chloramphenicol, [[dihydrostreptomycin, and thiophenicol. Various steroids and vitamin A are prepared using sodium borohydride in at least one step.

Safety

Sodium borohydride is a source of alkali, which is corrosive, and hydrogen or diborane, which are inflammable. Spontaneous ignition can result from solution of sodium borohydride in dimethylformamide.flammable.

See also

References

  1. ^ a b MSDS data (carl roth)
  2. ^ a b Banfi, L.; Narisano, E.; Riva, R.; Stiasni, N.; Hiersemann, M. “Sodium Borohydride” in Encyclopedia of Reagents for Organic Synthesis (Ed: L. Paquette) 2004, J. Wiley & Sons, New York. doi:10.1002/047084289X.rs052.
  3. ^ Schlesinger, H. I.; Brown, H. C.; Abraham, B.; Bond, A. C.; Davidson, N.; Finholt, A. E.; Gilbreath, J. R.; Hoekstra, H.; Horvitz, L.; Hyde, E. K.; Katz, J. J.; Knight, J.; Lad, R. A.; Mayfield, D. L.; Rapp, L.; Ritter, D. M.; Schwartz, A. M.; Sheft, I.; Tuck, L. D.; Walker, A. O. (1953). "New developments in the chemistry of diborane and the borohydrides. General summary". J. Am. Chem. Soc. 75: 186–90. doi:10.1021/ja01097a049.  
  4. ^ R. S. Kumar, A. L. Cornelius (2005). Appl. Phys. Lett. 87: 261916. doi:10.1063/1.2158505.  
  5. ^ Y. Filinchuk, A. V. Talyzin, D. Chernyshov, V. Dmitriev (2007). Phys. Rev. B 76: 092104. doi:10.1103/PhysRevB.76.092104.  
  6. ^ E. Kim, R. Kumar, P. F. Weck, A. L. Cornelius, M. Nicol, S. C. Vogel, J. Zhang, M. Hartl, A. C. Stowe, L. Daemen, Y. Zhao (2007). J. Phys. Chem. B 111 (50): 13873–13876. doi:10.1021/jp709840w.  
  7. ^ Peter Rittmeyer, Ulrich Wietelmann “Hydrides” in Ullmann's Encyclopedia of Industrial Chemistry 2002, Wiley-VCH, Weinheim. doi:10.1002/14356007.a13_199
  8. ^ Schubert, F.; Lang, K.; Burger, A. “Alkali metal borohydrides” (Bayer), 1960. German patent DE 1088930 19600915 (ChemAbs: 55:120851). Supplement to . to Ger. 1,067,005 (CA 55, 11778i). From the abstract: “Alkali metal borosilicates are treated with alkali metal hydrides in approx. 1:1 ratio at >100 °C with or without H pressure”.
  9. ^ Brown, H. C. “Organic Syntheses via Boranes” John Wiley & Sons, Inc. New York: 1975. ISBN 0-471-11280-1. page 260-1.
  10. ^ Seyden-Penne, J. "Reductions by the Alumino- and Borohydrides in Organic Synthesis"; VCH–Lavoisier: Paris, 1991.
  11. ^ da Costa, Jorge C.S.; Karla C. Pais, Elisa L. Fernandes, Pedro S. M. de Oliveira, Jorge S. Mendonça, Marcus V. N. de Souza, Mônica A. Peralta, and Thatyana R.A. Vasconcelos (2006). "Simple reduction of ethyl, isopropyl and benzyl aromatic esters to alcohols using sodium borohydride-methanol system" (PDF). Arkivoc: 128–133. http://www.arkat-usa.org/ark/journal/2006/I01_General/1523/05-1523A%20as%20published%20mainmanuscript.pdf. Retrieved 2006-08-29.  
  12. ^ C. R. Lucas, “Bis(5-Cyclopentadienyl) [Tetrahydroborato(1-)]Titanium” Inorganic Syntheses, 1977, Volume 17, p.93. doi:10.1002/9780470132487.ch27

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