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| Sodium hydride | |
|---|---|
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|
| Identifiers | |
| CAS number | [7646-69-7] |
| Properties | |
| Molecular formula | NaH |
| Molar mass | 23.99 g/mol |
| Appearance | colorless to gray solid |
| Density | 1.396 g/cm³, solid |
| Melting point |
800 °C with decomp. |
| Solubility in water | Reacts |
| Hazards | |
| MSDS | External MSDS |
| EU classification | Flammable (F) |
| NFPA 704 |
 3
3
2
|
| R-phrases | R15 |
| S-phrases | (S2), S7/8, S24/25, S42 |
| Related compounds | |
| Other cations | Potassium hydride |
| Related compounds | Sodium borohydride |
| Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa) Infobox references |
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Sodium hydride is the chemical compound with the formula NaH. It is primarily used as a strong base in organic synthesis. NaH is representative of the saline hydrides, meaning it is a salt-like hydride, composed of Na+ and H− ions, in contrast to the more molecular hydrides such as borane, methane, ammonia and water. It is an ionic material that is insoluble in organic solvents (although apparently soluble in molten Na), consistent with the fact that H− remains an unknown anion in solution. Because of the insolubility of NaH, all reactions involving NaH occur at the surface of the solid.
Contents |
Basic properties and structure
NaH is produced by the direct reaction of hydrogen and liquid sodium.[1] Pure NaH is colorless, although samples generally appear gray. NaH is ca. 40% denser than Na (0.968 g/cm³).
NaH, like LiH, KH, RbH, and CsH, adopts the NaCl crystal structure. In this motif, each Na+ ion is surrounded by six H− centers in an octahedral geometry. The ionic radii of H− (146 pm in NaH) and F− (133 pm) are comparable, as judged by the Na−H and Na−F distances.[2]
Applications in organic synthesis
As a strong base
First and foremost, NaH is a base of wide scope and utility in organic chemistry.[3] It is capable of deprotonating a range of even weak Brønsted acids to give the corresponding sodium derivatives. Typical "easy" substrates contain O-H, N-H, S-H bonds, including alcohols, phenols, pyrazoles, and thiols.
NaH most notably is employed to deprotonate carbon acids such as 1,3-dicarbonyls and analogues such as malonic esters. The resulting sodium derivatives can be alkylated. NaH is widely used to promote condensation reactions of carbonyl compounds via the Dieckmann condensation, Stobbe condensation, Darzens condensation, and Claisen condensation. Other carbon acids susceptible to deprotonation by NaH include sulfonium salts and DMSO. NaH is used to make sulfur ylides, which in turn are used to convert ketones into epoxides.
As a reducing agent
NaH reduces certain main group compounds, but analogous reactivity is unknown in organic chemistry. Notably boron trifluoride reacts to give diborane:[1]
- 6 NaH + 2 BF3 → B2H6 + 6 NaF
Si-Si and S-S bonds in disilanes and disulfides are also reduced.
Drying agent
Because of its rapid and irreversible reaction with water, NaH can be used to dry some organic solvents. Other drying agents are far more widely used, such as calcium hydride.
Practical considerations
Sodium hydride is sold by many chemical suppliers such as Sigma-Aldrich and ACROS, usually as a mixture of 60% sodium hydride (w/w) in mineral oil. Such a dispersion is safer to handle and weigh than pure NaH. The pure white solid is prepared by rinsing the oil with pentane or THF, care being taken that the washings will contain traces of NaH that can ignite in air. Reactions involving NaH require an inert atmosphere, such as nitrogen gas. Typically NaH is used as a suspension in THF, a solvent that resists deprotonation but solvates many organosodium compounds.
Safety
NaH can ignite in air, especially upon contact with water to release hydrogen, which is also flammable. Hydrolysis converts NaH into sodium hydroxide (NaOH), a caustic base. In practice, most sodium hydride is dispensed as a dispersion in oil, which can be safely handled in air.
References
- ^ a b Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 2001. ISBN 0-12-352651-5.
- ^ Wells, A.F. (1984). Structural Inorganic Chemistry, Oxford: Clarendon Press
- ^ Encyclopedia of Reagents for Organic Synthesis (Ed: L. Paquette) 2004, J. Wiley & Sons, New York. DOI: 10.1002/047084289.
Wikipedia content modification information:
- This page was last modified on 2 September 2008, at 19:11.
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