Sodium sulfide

Sodium sulfide is the name used to refer to the chemical compound Na2S but more commonly its hydrate Na2S.9H2O. Both are a 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.


Na2S adopts the "antifluorite" structure, 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 of a base to coexist with water. Thus, the dissolution process can be described as follows:
Na2S(s) + H2O(l) → 2Na+(aq) + SH + OH


Industrially Na2S is produced by reduction of Na2SO4 with carbon, in the form of coal:
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:

2 Na + S → Na2S


Na2S and its hydrates are dangerous and should only be handled by experts. Like lye, it is strongly alkaline and will cause skin burns. Acids react rapidly to produce hydrogen sulfide, which is a highly toxic gas.


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