Sodium sulfide
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| Sodium sulfide | |
|---|---|
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| Other names | Disodium sulfide |
| Identifiers | |
| CAS number | [], 1313-84-4 (pentahydrate) 1313-84-4 (nonahydrate) |
| PubChem | |
| EC number | |
| UN number | 1385 (anhydrous) 1849 (hydrate) |
| RTECS number | WE1905000 |
| Properties | |
| Molecular formula | Na2S |
| Molar mass | 78.0452 g/mol (anhydrous) 240.18 g/mol (nonahydrate) |
| Appearance | colorless, hygroscopic solid |
| Density | 1.856 g/cm3 (anhydrous) 1.58 g/cm3 (pentahydrate) 1.43 g/cm3 (nonohydrate) |
| Melting point |
1176 °C (anhydrous) |
| Solubility in water | 18.6 g/100 mL (20 °C) 39 g/100 mL (50 °C) |
| Solubility | insoluble in ether slightly soluble in alcohol |
| Structure | |
| Crystal structure | Antifluorite (cubic), cF12 |
| Space group | Fm3m, No. 225 |
| Coordination geometry |
Tetrahedral (Na+); cubic (S2–) |
| Hazards | |
| MSDS | ICSC 1047 |
| EU Index | 016-009-00-8 |
| EU classification | Corrosive (C) Dangerous for the environment (N) |
| R-phrases | R31, R34, R50 |
| S-phrases | (S1/2), S26, S45, S61 |
| NFPA 704 |
 1
3
1
|
| Autoignition temperature |
>480 ºC |
| Related compounds | |
| Other anions | Sodium oxide Sodium selenide Sodium telluride |
| Other cations | Lithium sulfide Potassium sulfide |
| Related compounds | Sodium hydrosulfide |
| 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 sulfide is the name used to refer to the chemical compound Na2S but more commonly its hydrate Na2S.9H2O. Both are colorless water-soluble salts that give strongly alkaline solutions. When exposed to moist air, Na2S and its hydrates emit hydrogen sulfide, which smells much like rotten eggs.
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[edit] Structure
Na2S adopts the antifluorite structure,[1][2] which means that the Na+ centers occupy sites of the fluoride in the CaF2 framework, and the larger S2− occupy the sites for Ca2+. In solution, the salt, by definition, dissociates. The dianion S2− does not, however, exist in appreciable amounts in water. Sulfide is too strong a base to coexist with water. Thus, the dissolution process can be described as follows:
- Na2S(s) + H2O(l) → 2Na+(aq) + HS− + OH−
[edit] Production
Industrially Na2S is produced by reduction of Na2SO4 with carbon, in the form of coal:[3]
- Na2SO4 + 4 C → Na2S + 4 CO
In the laboratory, the anhydrous salt can be prepared by reduction of sulfur with sodium in anhydrous ammonia. Alternatively, sulfur can be reduced by sodium in dry THF with a catalytic amount of naphthalene:[4]
- 2 Na + S → Na2S
[edit] Safety
Like sodium hydroxide, sodium sulfide is strongly alkaline and can cause skin burns. Acids react with it to rapidly produce hydrogen sulfide, which is a toxic and foul-smelling gas.
[edit] References
- ^ Zintl, E.; Harder, A.; Dauth B. (1934), "Gitterstruktur der oxyde, sulfide, selenide und telluride des lithiums, natriums und kaliums", Z. Elektrochem. Angew. Phys. Chem. 40: 588–93
- ^ Wells, A.F. (1984) Structural Inorganic Chemistry, Oxford: Clarendon Press. ISBN 0-19-855370-6.
- ^ Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 2001. ISBN 0-12-352651-5.
- ^ So, J.-H.; Boudjouk, P. (1992). "Hexamethyldisilathiane". Inorg. Synth. 29: 30. doi:.
[edit] External links
- chemicalland21.com Sodium sulfide.
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