Sodium sulfide

Sodium sulfide
Other names Disodium sulfide
Identifiers
CAS number	1313-82-2 Y,  1313-84-4 (pentahydrate) 1313-84-4 (nonahydrate)
PubChem	237873
EC number	215-211-5
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) 100 °C (pentahydrate) 50 °C (nonhydrate)
Solubility in water	186 g/L (20 °C) 390 g/L (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
Y sulfide (verify) (what is: Y/N?) Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa)
Infobox references

Sodium sulfide is the name used to refer to the chemical compound Na₂S, but more commonly it refers to the 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 or flatus. Generally, commercially available sodium sulfide is not a unique chemical entity, but it is 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. These grades of sodium sulfide are marketed as 'sodium sulfide flakes'. Although the solid is yellow, solutions of it are colorless.

Structure

Fluorite 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⁻

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

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

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

Uses

It is primarily used in pulp and paper industry in the kraft process . It is used in water treatment as an oxygen scavenger agent, in the photographic industry to protect developer solutions from oxidation, in textile industry as a bleaching, 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, food preservative, making dyes, and detergent.

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.

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. doi:10.1002/9780470132609.ch11. ISBN 9780470132609. 

External links

• chemicalland21.com Sodium sulfide.