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.

Structure

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⁻

Production

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

Safety

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.

References

External links

• chemicalland21.com Sodium sulfide