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Sodium sulfide

Sodium sulfide
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Other names
Disodium sulfide
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1313-82-2 7pxY

1313-84-4 (pentahydrate)
1313-84-4 (nonahydrate)
ChEBI CHEBI:76208 7pxN
ChemSpider 14120 7pxN
EC number 215-211-5
Jmol-3D images Image
PubChem Template:Chembox PubChem/format
RTECS number WE1905000
UN number 1385 (anhydrous)
1849 (hydrate)
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Na2S
Molar mass 78.0452 g/mol (anhydrous)
240.18 g/mol (nonahydrate)
Appearance colorless, hygroscopic solid
Odor rotten eggs
Density 1.856 g/cm3 (anhydrous)
1.58 g/cm3 (pentahydrate)
1.43 g/cm3 (nonohydrate)
Melting point Script error: No such module "convert". (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
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Crystal structure Antifluorite (cubic), cF12
Space group Fm3m, No. 225
Tetrahedral (Na+); cubic (S2−)
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SDS 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

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Other anions
Sodium oxide
Sodium selenide
Sodium telluride
Other cations
Lithium sulfide
Potassium sulfide
Related compounds
Sodium hydrosulfide
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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Infobox references

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 Na2xH2O, 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.

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+.

Production

Industrially Na2S is produced by carbothermic reduction of sodium sulfate often using coal:[3]

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:[4]

2 Na + S → Na2S

Sodium sulfide can also be produced by the reaction with sodium cyanide and sulfuric acid, however, toxic hydrogen cyanide is formed in the process.

Reactions with inorganic reagents

The dissolution process can be described as follows:

Na2S + H2O → NaHS + NaOH

Sodium sulfide can oxidize when heated to sodium carbonate and sulfur dioxide:

2 Na2S + 3 O2 + 2 CO2 → 2 Na2CO3 + 2 SO2

Oxidation with hydrogen peroxide gives sodium sulfate:[5]

2 Na2S + 4 H2O2 → 4 H2O + Na2SO4

Upon treatment with sulfur, polysulfides are formed:

2Na2S + S8 → 2 Na2S5

Uses

Sodium sulfide is primarily used in pulp and paper industry in the Kraft process.

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

Even aryl halides participate in this reaction.[6] Sodium sulfide reduces1,3-dinitrobenzene derivatives to the 3-nitroanilines.[7]

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 highly toxic.

References

  1. ^ 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. 
  2. ^ Wells, A.F. (1984) Structural Inorganic Chemistry, Oxford: Clarendon Press. ISBN 0-19-855370-6.
  3. ^ Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 2001. ISBN 0-12-352651-5.
  4. ^ So, J.-H; Boudjouk, P; Hong, Harry H.; Weber, William P. (1992). "Hexamethyldisilathiane". Inorg. Synth. Inorganic Syntheses 29: 30. ISBN 978-0-470-13260-9. doi:10.1002/9780470132609.ch11. 
  5. ^ 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
  6. ^ Charles C. Price, Gardner W. Stacy "p-Aminophenyldisulfide" Org. Synth. 1948, vol. 28, 14. doi:10.15227/orgsyn.028.0014
  7. ^ Hartman, W. W.; Silloway, H. L. (1955). "2-Amino-4-nitrophenol". Org. Synth. ; Coll. Vol. 3, p. 82