Sodium borohydride

Sodium borohydride
Names
Preferred IUPAC name
Sodium tetrahydridoborate (1–)
Systematic IUPAC name
Sodium boranuide
Identifiers
16940-66-2 YesY
15681-89-7 (2D4) YesY
ChEBI CHEBI:50985 YesY
ChemSpider 26189 YesY
9052313 (2D4) YesY
9312193 (3T4) YesY
EC Number 241-004-4
23167
Jmol 3D model Interactive image
MeSH Sodium+borohydride
PubChem 4311764
23673181 (2D4)
23671303 (3T4)
RTECS number ED3325000
UN number 1426
Properties
NaBH4
Molar mass 37.83 g/mol
Appearance white crystals
hygroscopic
Density 1.0740 g/cm3
Melting point 400 °C (752 °F; 673 K)[1]
Boiling point 500 °C (932 °F; 773 K) (decomposes)[1]
soluble, reacts with water
Solubility soluble in liquid ammonia, amines, pyridine
Hazards
H260, H301, H311, H314
P223, P231, P232, P280, P301+310, P370+378, P422
NFPA 704
Flammability code 1: Must be pre-heated before ignition can occur. Flash point over 93 °C (200 °F). E.g., canola oil Health code 3: Short exposure could cause serious temporary or residual injury. E.g., chlorine gas Reactivity code 2: Undergoes violent chemical change at elevated temperatures and pressures, reacts violently with water, or may form explosive mixtures with water. E.g., phosphorus Special hazard W: Reacts with water in an unusual or dangerous manner. E.g., cesium, sodiumNFPA 704 four-colored diamond
1
3
2
Flash point 70 °C (158 °F; 343 K)
ca. 220 °C (428 °F; 493 K)
Explosive limits 3%
Lethal dose or concentration (LD, LC):
160 mg/kg (Oral - Rat)
230 mg/kg (Dermal - Rabbit)
Related compounds
Other anions
 Sodium cyanoborohydride
Sodium hydride
Sodium borate
Borax
Other cations
 Lithium borohydride
Related compounds
 Lithium aluminium hydride
Sodium triacetoxyborohydride
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
YesY verify (what is YesYN ?)
Infobox references

Sodium borohydride, also known as sodium tetrahydridoborate, and sodium tetrahydroborate[2] 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. It has been tested as pretreatment for pulping of wood, but is too costly to be commercialized.[3][4] The compound is soluble in alcohols and certain ethers but reacts with water in the absence of a base.[5]

The compound was discovered in the 1940s by H. I. Schlesinger, who led a team that developed metal borohydrides for wartime applications (in particular, looking for a uranium compound more volatile than the hexafluoride to be used in isotope separation by gaseous diffusion; this line of research did not yield useful results).[6] Their work was declassified and published only in 1953.

Properties

Solvent Solubility (g/100 mL)[5]
MeOH 13
EtOH 3.16
Diglyme 5.15
Et2O insoluble

Sodium borohydride is an odorless white to gray-white microcrystalline powder which often forms lumps. It can be purified by recrystallization from warm (50 °C) diglyme.[7] Sodium borohydride is soluble in protic solvents such as water and lower alcohols; it will also react with these solvents to produce H2; however, these reactions are fairly slow. Complete decomposition in excess methanol can take nearly 90 min at 20 °C.[8] It will decompose in neutral or acidic aqueous solutions but is stable at pH 14.[5] These conditions can be exploited to allow sodium borohydride to react in a homogeneous manner, with reduced lifespan being traded against increased reactivity.

Structure

NaBH4 is a salt, consisting of the tetrahedral BH4− anion. The solid is known to exist as three 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.[9][10][11]

α-NaBH4
β-NaBH4
γ-NaBH4

Synthesis and handling

Sodium borohydride is prepared industrially following the original method of Schlesinger: sodium hydride is treated with trimethyl borate at 250–270 °C:

B (OCH3)3 + 4 NaH → NaBH4 + 3 NaOCH3

Millions of kilograms are produced annually, far exceeding the production levels of any other hydride reducing agent.[3] Sodium borohydride can also be produced by the action of NaH on powdered borosilicate glass.[12]

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, thiol esters and imines. Under typical conditions, it will not reduce esters, amides, or carboxylic acids.[5] At room temperature, the only acid derivatives it reduces are acyl chlorides, which are exceptionally electrophilic.

Many other hydride reagents are more strongly reducing. These usually involve replacing hydride with alkyl groups, such as 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.[13]

The reactivity of NaBH4 can be enhanced or augmented by a variety of compounds.[14][15] Oxidation with iodine in tetrahydrofuran gives the borane–tetrahydrofuran complex, which can reduce carboxylic acids.[16] Likewise, the NaBH4-MeOH system, formed by the addition of methanol to sodium borohydride in refluxing THF, reduces esters to the corresponding alcohols.[17] Mixing water or an alcohol with the borohydride converts some of it into unstable hydride ester, which is more efficient at reduction, but the reductant will eventually decompose spontaneously to give hydrogen gas and borates. The same reaction can run also intramolecularly: an α-ketoester converts into a diol, since the alcohol produced will attack the borohydride to produce an ester of the borohydride, which then reduces the neighboring ester.[18] The combination of NaBH4 with carboxylic acids results in the formation of acyloxyborohydride species. These can perform a variety of reductions not normally associated with borohydride chemistry, such as alcohols to hydrocarbons and nitriles to primary amines.[19]

Coordination chemistry

BH4− is a ligand for metal ions. Such borohydride complexes are often prepared by the action of NaBH4 (or the LiBH4) on the corresponding metal halide. One example is the titanocene derivative:[20]

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

Hydrogen source

In the presence of metal catalysts, sodium borohydride releases hydrogen. Exploiting this reactivity, sodium borohydride is used in prototypes of the direct borohydride fuel cell. The hydrogen is generated for a fuel cell by catalytic decomposition of the aqueous borohydride solution:

NaBH4 + 2 H2O → NaBO2 + 4 H2 (ΔH < 0)

Applications

The principal application of sodium borohydride is the production of sodium dithionite from sulfur dioxide: Sodium dithionite is used as a bleaching agent for wood pulp and in the dyeing industry.

Sodium borohydride reduces aldehydes and ketones to give the related 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.

Sodium borohydride is used as reducing agent in the synthesis of gold nanoparticles.[21]

Sodium borohydride has been considered as a solid state hydrogen storage candidate. Although practical temperatures and pressures for hydrogen storage have not been achieved, in 2012 a core–shell nanostructure of sodium borohydride was used successfully to store, release and reabsorb hydrogen under moderate conditions.[22]

Safety

Sodium borohydride is a source of basic borate salts which can be corrosive, and hydrogen or diborane, which are both flammable. Spontaneous ignition can result from solution of sodium borohydride in dimethylformamide.

See also

References

  1. 1 2 MSDS data (carl roth)
  2. ↑ Busch, D.H. (2009). Inorganic Syntheses 20. Wiley. p. 137. ISBN 9780470132869. Retrieved 20 May 2015.
  3. 1 2 Peter Rittmeyer, Ulrich Wietelmann "Hydrides" in Ullmann's Encyclopedia of Industrial Chemistry 2002, Wiley-VCH, Weinheim. doi:10.1002/14356007.a13_199
  4. ↑ Istek, A. and Gonteki, E. "Utilization of sodium borohydride (NaBH4) in kraft pulping process." Retrieved online on 25 September 2014 at http://www.jeb.co.in/journal_issues/200911_nov09/paper_05.pdf.
  5. 1 2 3 4 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.
  6. ↑ 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.
  7. ↑ Brown, H. C. “Organic Syntheses via Boranes” John Wiley & Sons, Inc. New York: 1975. ISBN 0-471-11280-1. page 260-1.
  8. ↑ Lo, Chih-ting F.; Karan, Kunal; Davis, Boyd R. "Kinetic Studies of Reaction between Sodium Borohydride and Methanol, Water, and Their Mixtures". Industrial & Engineering Chemistry Research 46 (17): 5478–5484. doi:10.1021/ie0608861.
  9. ↑ R. S. Kumar, A. L. Cornelius (2005). "Structural transitions in NaBH[sub 4] under pressure". Appl. Phys. Lett. 87 (26): 261916. doi:10.1063/1.2158505.
  10. ↑ Y. Filinchuk, A. V. Talyzin, D. Chernyshov, V. Dmitriev (2007). "High-pressure phase of NaBH4: Crystal structure from synchrotron powder diffraction data". Phys. Rev. B 76 (9): 092104. doi:10.1103/PhysRevB.76.092104.
  11. ↑ 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). "Pressure-driven phase transitions in NaBH4: theory and experiments". J. Phys. Chem. B 111 (50): 13873–13876. doi:10.1021/jp709840w. PMID 18031032.
  12. ↑ 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".
  13. ↑ Seyden-Penne, J. "Reductions by the Alumino- and Borohydrides in Organic Synthesis"; VCH–Lavoisier: Paris, 1991.
  14. ↑ Periasamy, Mariappan; Thirumalaikumar, Muniappan (2000). "Methods of enhancement of reactivity and selectivity of sodium borohydride for applications in organic synthesis". Journal of Organometallic Chemistry 609 (1–2): 137–151. doi:10.1016/S0022-328X(00)00210-2.
  15. ↑ Nora de Souza, Marcus Vinícius; Alves Vasconcelos; Thatyana Rocha (1 November 2006). "Recent methodologies mediated by sodium borohydride in the reduction of different classes of compounds". Applied Organometallic Chemistry 20 (11): 798–810. doi:10.1002/aoc.1137.
  16. ↑ Kanth, J. V. Bhaskar; Periasamy, Mariappan (1 September 1991). "Selective reduction of carboxylic acids into alcohols using sodium borohydride and iodine". The Journal of Organic Chemistry 56 (20): 5964–5965. doi:10.1021/jo00020a052.
  17. ↑ 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. Retrieved 29 August 2006. Cite uses deprecated parameter |coauthors= (help)
  18. ↑ V. Dalla, J. P. Catteau and P. Pale. Mechanistic rationale for the NaBH4 reduction of α-keto esters. Tetrahedron Letters, Volume 40, Issue 28, 9 July 1999, Pages 5193–5196.
  19. ↑ W. Gribble, Gordon (1 January 1998). "Sodium borohydride in carboxylic acid media: a phenomenal reduction system". Chemical Society Reviews 27 (6): 395. doi:10.1039/A827395Z.
  20. ↑ Lucas, C. R. (1977). "Bis(5-Cyclopentadienyl) [Tetrahydroborato(1-)]Titanium". Inorganic Syntheses 17: 93. doi:10.1002/9780470132487.ch27.
  21. ↑ Low and Bansal, 2009
  22. ↑ Stuart Gary, "Hydrogen storage no longer up in the air" in ABC Science 16 August 2012, citing Christian, Meganne; Aguey-Zinsou, Kondo François (2012). "Core–Shell Strategy Leading to High Reversible Hydrogen Storage Capacity for NaBH4". ACS Nano (American Chemical Society). doi:10.1021/nn3030018. Retrieved 20 August 2012.

External links

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