Very interesting information on NaBH4
(Rated as: excellent)
3.4. Sodium Tetrahydridoborate
Properties . (See table 2)Sodium tetrahydridoborate, sodium borohydride, sodium boranate, NaBH4 , is a white, slightly hygroscopic salt that is stable in dry air up to 300 °C. It decomposes slowly in a vacuum at 400 °C, quickly above 550 °C, with liberation of hydrogen; mp (under 1 MPa H2) 505 °C. If ignited with a naked flame it burns quietly. It can be stored in an airtight container for many years without decomposition.
Sodium borohydride absorbs water from moist air to form the dihydrate, NaBH4 · 2 H2O, which slowly decomposes to form sodium metaborate and hydrogen. In acidic media rapid hydrolysis takes place:
BH4– + H+ + 3H2O __> H3BO3+ 4 H2;
D H = – 372 kJ/mol
Sodium borohydride, in keeping with its ionic character, is readily soluble in polar solvents, especially protic compounds such as water, lower alcohols, and amines (see Table (4)).
With a few solvents it forms isolable though not very stable solvates (37). In protic solvents, solvolysis occurs, and hydrogen is evolved:
NaBH4 + 4 ROH __> NaB(OR)4 + 4 H2
Thus, sodium borohydride is 80 % decomposed after 1 h in methanol at 0 °C, 6 % decomposed after 1 h in ethanol at 60 °C, but is stable in 2-propanol and tert-butanol and at room temperature. The solubility and stability in water, the most frequently used solvent, have been extensively investigated (45) (see Fig. (4)). Protonolysis takes place in aqueous solution, the rate increasing rapidly with increasing temperature and decreasing pH (1). These solutions become more stable with increasing concentration of sodium borohydride. At 24 °C, a 10–2 M solution (pH = 9.56) decomposes to the extent of 15 % in 1 h, whereas a 0.10 M solution (pH = 10.05) undergoes only 5 % decomposition in the same time. On boiling or on acidification, however, rapid hydrolysis occurs and all the combined hydrogen is released (2.36 L/g at STP). A number of transition metal salts and finely divided metals such as nickel, cobalt, iron, and copper (1) , (46) catalyze the hydrolysis, but the addition of strong bases such as NaOH can stabilize solutions of NaBH4 (Fig. (5)).
A solution of sodium borohydride in concentrated sodium hydroxide can be stored virtually without decomposition, although concentrated solutions of sodium borohydride in dimethylformamide must not be heated because of the danger of explosion caused by autocatalytic decompositions; such solutions should not even be stored for long periods at room temperature (47). Contact with strong acid (e.g., concentrated sulfuric or phosphoric acid, or methanesulfonic acid) or Lewis acids must be avoided as these substances can liberate diborane, which is highly toxic and spontaneously flammable.
Sodium borohydride is a powerful reducing agent with wide applications in inorganic and particularly organic chemistry. The redox potential E0 of the half reaction
BH4– + 8 OH– __> B(OH)4– + 4 H2O + 8 e–
at pH 14 is – 1.24 V.
An aqueous solution of sodium borohydride will reduce a number of cations and anions, the product obtained depending on the nature of the ion and the reaction conditions. A metal ion of lower valency can be produced (e.g., iron(II) from iron(III)), a metal (e.g., all the platinum metals, cadmium, mercury, lead, silver, and gold), a metallic boride (CrB2, Co2B, Ni2B), or a hydride (GeH4, SnH4 , AsH3, BiH3). Many of these reductions are used for the recovery of noble metals, the precipitation of toxic metals from water, the preparation of f inely divided metals, and the electroless deposition of metals and metallic borides onto metallic and nonmetallic substrates (48) , (49).
Another reaction of industrial importance is the reduction of sulfur dioxide or sodium hydrogen sulf ite to sodium dithionite:
NaBH4 + 8 NaOH + 8 SO2 __> 4 Na2S2O4 + NaBO2 + 6 H2O
The reducing power of sodium borohydride towards some important functional groups is shown in Table (3). This reveals that NaBH4 is a rather milder but also more selective reducing agent than NaAlH4 , allowing the reduction of aldehydes, ketones, acid chlorides, and imines in the presence of esters, epoxides, amides, nitriles, or nitro groups. Much research has been undertaken with the object of varying the reducing power over a wide range (50) , e.g., by mixing sodium borohydride with metal salts (Pd(NO3)2 , AlCl3 , CoCl2 , NiCl2), thiols (e.g., ethanedithiol), or carboxylic acids (e.g., acetic acid, trifluoroacetic acid), by using different solvents (e.g., aprotic solvents or solvent mixtures) (51) , or by using phase transfer catalysts (52).
A solution of NaBH4 – TiCl4 (ca. 3 : 1) in 1,2-dimethoxyethane reduces carboxylic acids to alcohols and nitro compounds to amines. The NaBH4 – CoCl2 system dissolved in lower alcohols or dioxane is also effective for these reductions, as well as the reduction of sulfoxides to sulfides and the hydrogenation of alkenes. Lithium chloride and sodium borohydride dissolved in diethylene glycol dimethyl ether form the intermediate LiBH4 which, unlike NaBH4 , can reduce esters and epoxides. Sodium triacetoxyborohydride, NaBH(OAc)3 , produced in situ from NaBH4 and three equivalents of acetic acid, exhibits such a weakening of reducing power that aldehydes are selectively reduced without ketones being affected (53). Sodium borohydride and boron trifluoride in diglyme or THF yield diborane:
3 NaBH4 + 4 BF3· OEt2 __> 2 B2H6+ 3 NaBF4 + 4 Et2O
Diborane is used mainly for hydroboration, but also acts as a strong electrophilic reducing agent, reducing carboxylic acids and nitriles but not acyl chlorides or nitro groups (54).
Sodium borohydride is also an important starting material for the production of tetrahydridoborates (borohydrides) of other metals, and of boranes and organoboron compounds, their derivatives and adducts . Some examples of practical importance are:
Production. Since the discovery of sodium borohydride in 1942 by H. J. SCHLESINGER, H. C. BROWN, H. R. HOECKSTRA, and L. R. RAPP (55) , over 100 methods of preparation have been described, but few of these have achieved any practical significance. Industrial production in the United States started in 1954 using the Schlesinger process (56) , in which trimethyl borate reacts with an extremly fine dispersion of sodium hydride in a high boiling hydrocarbon oil at 250 – 280 °C and atmospheric pressure:
4 NaH + B(OCH3)3 __> NaBH4 + 3 NaOCH3
Water is then added, causing separation of the hydrocarbon oil, which is recirculated. The aqueous phase contains the NaBH4 together with caustic soda and methanol formed by hydrolysis of NaOCH3. The methanol is recovered by distillation and used again for production of trimethyl borate. The product solution contains 12 % NaBH4 and 42 % NaOH (r = 1.4 g/cm3 ; h23 = 79 mPa · s) and is sometimes marketed as Sodium Borohydride – SWS by Ventron or alkaline aqueous solution by Nokia. Solid sodium borohydride is obtained from the alkaline aqueous solution by extraction with isopropylamine; the yield is better than 90 %.
A process developed by Bayer (57) (see Fig. (6)) uses finely ground borosilicate glass, sodium, and hydrogen:
Na2B4O7 · 7 SiO2 + 16 Na + 8 H2 __> 4 NaBH4+ 7 Na2SiO3
The borosilicate Na2B4O7 · 7 SiO2 is produced by fusion of borax and silica. The borosilicate is cooled, ground, and then reacted with sodium in an atmosphere of hydrogen at 300 kPa and 400 – 500 °C in a partly heterogeneous reaction. The sodium borohydride is extracted from the borohydride – silicate mixture with liquid ammonia under pressure. The yield is in excess of 90 %. Other processes are described in Refs. (1) , (5), and (57).
Commercial sodium borohydride is in powder – lump form (bulk density 0.4 – 0.5 kg/L), the purity being typically 98 % (Bayer, Chemetall, Ventron).
Analysis. Two main methods are used for the analysis of sodium borohydride: one gas volumetric and the other titrimetric. In the f irst of these, the sodium borohydride is reacted with acid and the evolved hydrogen measured. The second method is equally straightforward and requires no special analytical equipment. The sodium borohydride is converted to metaborate by reaction with excess oxidizing agent such as sodium hypochlorite or potassium iodate:
NaBH4 + 4 NaOCl __>® NaBO2+ 4 NaCl + 2 H2O
The excess oxidizing agent is determined by iodometric back-titration (48) , (49).
Uses. The present annual consumption of sodium borohydride is estimated to be 2000 – 3000 t, mainly in the form of the 12 % solution. This means that it is by far the most important commercially available complex hydride. The largest amount is used in the paper and pulp industries, mainly in the United States and Scandinavia. The 12 % solution (Ventron, Nokia) is used for the in situ production of sodium dithionite, which is used to bleach wood pulp. Both sodium dithionite (made from sodium borohydride and sulfur dioxide) and sodium borohydride itself are used for reducing vat dyes to the leuco bases (58). A recent patent claims that the reaction with sulfur dioxide can be used for the purification of exhaust gases from power stations burning fossil fuel (59).
Solid sodium borohydride is used as a reducing agent in the organic chemical and pharmaceutical industries, and although the quantities used for this purpose are considerably smaller than the demands of the paper industry, the selective reduction of aldehydes and ketones has become of major importance (24). One of the biggest uses is in the stereoselective reduction of steroid ketones. Another is the synthesis of antibiotics such as chloramphenicol, dihydrostreptomycin, and thiophenicol, and also vitamin A. Other examples are reduction in the synthesis of prostaglandins, atropine, scopolamine, and flavorings and aromas.
Sodium borohydride is also used for cleaning and stabilization of process streams of many organic products. It chemically reduces harmful impurities such as aldehydes, ketones and peroxides, thereby improving color, odor, and the stability of the products towards heat, light, and oxidation. The technique is used in the production and subsequent treatment of epoxides, alcohols, amines, ethers, esters, carboxylic acids, alkenes, polymers (58) , and paper (60).
The reduction of metal cations already mentioned is used on a large scale for the electroless deposition of metals and metal borides onto metals and other materials such as glass, ceramics, and plastics ("Nibodur 1098" — Bayer (61) , for the recovery of mercury from wastewater from the chlor-alkali electrolytic process (Ventron process, Koertrol) of lead from lead tetraalkyl production, and of silver and cadmium from spent photographic developing solutions (62) , (63). Finely divided metals of the iron group for high quality magnetic tapes are produced by sodium borohydride reduction (58) , (64). Nickel boride and cobalt boride obtained by the reduction of the appropriate salt solutions are used as catalysts in gas- and liquid-phase reduction (58). One final application for sodium borohydride depends on its ability to generate hydrogen gas: it is used as a foaming agent for organic polymers and inorganic construction materials (58) , (65).
(37) B. D. James, M. G. H. Wallbridge: "Metal Tetrahydroborates," Prog. Inorg. Chem. 11 (1970) 99 – 231.
(45) E. H. Jensen: A Study on Sodium Borohydride, Nyt Nordisk Forlag, Arnold Busck, Copenhagen 1954.
(46) A. Levy, J. B. Brown, C. J. Lyons, Ind. Eng. Chem. 52 (1960) 211.
(47) D. A. Yeowell, R. L. Seaman, J. Mentha, Chem. Eng. News 57 (Sept. 24, 1979) 4.
(48) Bayer AG, Firmenschrift, Natriumboranat, 1975.
(49) Ventron Corp., Firmenschrift Sodium Borohydride, 1979.
(50) R. C. Wade, J. Mol. Catal. 18 (1983) 273 – 297.
(51) R. S. Varma, G. W. Kabalka, Synth. Commun. 15 (1985) 985 – 990.
(52) F. Rolla, J. Org. Chem. 46 (1981) 3909 – 3911.
(53) G. W. Gribble, C. F. Nutaitis, Org. Prep. Proced. Int. 17 (1985) 317 – 384.
(54) M. Follet, Chem. Ind. (London) 1986, 123 – 128.
(55) H. C. Brown: Boranes in Organic Chemistry, Cornell Univ. Press, Ithaka 1972.
(56) H. J. Schlesinger, H. C. Brown, J. Am. Chem. Soc. 75 (1953) 186 – 205; H. J. Schlesinger, H. C. Brown, US 2 534 533, 1945.
(57) W. Büchner, H. Niederprüm, Pure Appl. Chem. 49 (1977) 733 – 743.
(5) A. Hajos: Complex Hydrides, Elsevier, Amsterdam 1979.
(58) R. C. Wade, Spec. Publ. R. Soc. Chem. 40 (1981) 25 – 58.
(59) Gea Wiegand GmbH, DE-OS 3 525 377, 1985 (U. Hochberg, G. Klinke, P. Sauglet).
(24) A. Kleemann, J. Engel: Pharmazeutische Wirkstoffe, Thieme Verlag, Stuttgart 1982; Ergänzungsband, Thieme Verlag, Stuttgart 1987.
(60) L. C. Tang, Adv. Chem. Ser. 212 (1986) 427 – 441.
(61) R. N. Duncan, T. L. Arney, Plat. Surf. Finish. 71 (1984) 49 – 54.
(62) J. A. Ulman, Spec. Publ. R. Soc. Chem. 61 (1986) 173 – 196.
(63) M. J. Lindsay, M. E. Hackmann, R. I. Tremblay, Proc. Ind. Waste Conf. 40 (1985) 477 – 482.
(64)Mitsubishi Denki KK, DE-OS 3 704 075, 1986 (H. Ogama, Y. Fukotomi, M. Okutani).
(65)Mosc. Eng. – Cons. Inst., SU 126 167, 1984 (A. P. Merkin, B. M. Runyantsev, L. E. Vitels).
Taken from Ullman's encyclopedia of industrial chemistry
I will post the tables tommorow, they need a lot of editing to fit this forum
Physical data for alkali metal borohydrides and aluminum hydrides
Hydride CAS registry Mr Hydrogen Density, mp,
number content, wt% g/cm3 °C
LiBH4 [16949-15-8] 21.78 18.5 0.66 280
LiBH(C2H5)3 [22560-16-3] 105.94 0.95
NaBH4 [16940-66-2] 37.83 10.6 1.07 505 **
NaBH3CN [25895-60-7] 62.84 6.4 1.20 240 dec.
KBH4 [13762-51-1] 53.94 7.4 1.17 585 **
LiAlH4 [16853-85-3] 37.95 10.6 0.91 >125 dec.
NaAlH4 [13770-96-2] 54.00 7.4 1.27 178
** Under H2 atmosphere.
Table 4 (blank spaces come from the table itself,no info)
Solubility of complex borohydrides in various solvents (g/100 g solvent at 25 °C unless otherwise stated)
Solvent bp of solvent, LiBH4 NaBH4 KBH4 NaBH3CN LiHBEt3
Water 100.0 20.9 a 55 19.3 217 dec
Methanol 67.7 dec. 16.4 b 0.7 v. sol. dec.
Ethanol 78.5 44 4.0 b 0.25 sol. dec.
Isopropylamine 34.0 8 c 6.0 d fairly sol. dec.
Diethyl ether 36.0 4.3 insol. insol. insol.
THF 65.0 22.5 0.1 e insol. 36 v. sol.
Diglyme 162 5.5 insol. 17.6
Toluene 111 insol. insol. insol. insol.
Ammonia –33.3 77 104 20
DMF 153 18.0 f 15 f
a 0 °C, slow decomposition.
b 20 °C, decomposition.
c 10 °C.
d 28 °C.
e 20 °C.
f Dangerous decomposition possible at higher temperature.
Table 3 is too big to fit here.
Taken from Ullman's Encyclopedia of industrial chemistry