|Jmol 3D model||Interactive image|
|UN number||1385 (anhydrous)
|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||1,176 °C (2,149 °F; 1,449 K) (anhydrous)
100 °C (pentahydrate)
50 °C (nonahydrate)
|12.4 g/100 mL (0 °C)
18.6 g/100 mL (20 °C)
39 g/100 mL (50 °C)
|Solubility||insoluble in ether
slightly soluble in alcohol
|Antifluorite (cubic), cF12|
|Fm3m, No. 225|
|Tetrahedral (Na+); cubic (S2−)|
|Safety data sheet||ICSC 1047|
EU classification (DSD)
Dangerous for the environment (N)
|R-phrases||R31, R34, R50|
|S-phrases||(S1/2), S26, S45, S61|
|> 480 °C (896 °F; 753 K)|
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
|what is ?)(|
Sodium sulfide is the chemical compound with the formula Na2S, or 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 like rotten eggs. Some commercial samples are specified as Na2S·xH2O, where a weight percentage of Na2S is specified. Commonly available grades have around 60% Na2S by weight, which means that x is around 3. Such technical grades of sodium sulfide have a yellow appearance owing to the presence of polysulfides. These grades of sodium sulfide are marketed as 'sodium sulfide flakes'. Although the solid is yellow, solutions of it are colorless.
- Na2SO4 + 2 C → Na2S + 2 CO2
In the laboratory, the 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:
- 2 Na + S → Na2S
Sodium sulfide can also be produced by the reaction with sodium cyanide and sulfuric acid, however, extremely toxic hydrogen cyanide is formed in the process.
Reactions with inorganic reagents
The dissolution process can be described as follows:
- S−2 + H2O → HS− + OH−
- 2 Na2S + 3 O2 + 2 CO2 → 2 Na2CO3 + 2 SO2
Oxidation with hydrogen peroxide gives sodium sulfate:
- Na2S + 4 H2O2 → 4 H2O + Na2SO4
Upon treatment with sulfur, polysulfides are formed:
- 2 Na2S + S8 → 2 Na2S5
It is used in water treatment as an oxygen scavenger agent and also as a metals precipitant, in chemical photography for toning black and white photographs, in textile industry as a bleaching, and as a desulfurising and as a dechlorinating agent and in leather trade for the sulfitisation of tanning extracts. It is used in chemical manufacturing as a sulfonation and sulfomethylation agent. It is used in the production of rubber chemicals, sulfur dyes and other chemical compounds. It is used in other applications including ore flotation, oil recovery, making dyes, and detergent. It is also used for leather processing in liming operation as unhairing agent.
Reagent in organic chemistry
Alkylation of sodium sulfide give thioethers:
- Na2S + 2 RX → R2S + 2 NaX
- 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; Hong, Harry H.; Weber, William P. (1992). "Hexamethyldisilathiane". Inorg. Synth. Inorganic Syntheses. 29: 30. doi:10.1002/9780470132609.ch11. ISBN 978-0-470-13260-9.
- L. Lange, W. Triebel, "Sulfides, Polysulfides, and Sulfanes" in Ullmann's Encyclopedia of Industrial Chemistry 2000, Wiley-VCH, Weinheim. doi:10.1002/14356007.a25_443
- Charles C. Price, Gardner W. Stacy "p-Aminophenyldisulfide" Org. Synth. 1948, vol. 28, 14. doi:10.15227/orgsyn.028.0014
- Hartman, W. W.; Silloway, H. L. (1955). "2-Amino-4-nitrophenol". Org. Synth. ; Coll. Vol., 3, p. 82