Sodium sulfide Names Other names Identifiers ChEBI ChemSpider ECHA InfoCard 100.013.829 EC Number RTECS number UNII UN number 1385 (anhydrous)1849 (hydrate) Properties Na2S Molar mass 78.0452 g/mol (anhydrous)240.18 g/mol (nonahydrate) Appearance colorless, hygroscopic solid Odor Foul, pungent, like that of rotten eggs (hydrolysis in atmospheric moisture to hydrogen sulfide) 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) (hydrolyses) Solubility insoluble in ether slightly soluble in alcohol[1] −39.0·10−6 cm3/mol Structure Antifluorite (cubic), cF12 Fm3m, No. 225 Tetrahedral (Na+); cubic (S2−) Hazards GHS labelling: Danger H302, H311, H314, H400 P260, P264, P270, P273, P280, P301+P312, P301+P330+P331, P302+P352, P303+P361+P353, P304+P340, P305+P351+P338, P310, P312, P321, P322, P330, P361, P363, P391, P405, P501 NFPA 704 (fire diamond) > 480 °C (896 °F; 753 K) Safety data sheet (SDS) ICSC 1047 Related compounds Sodium oxideSodium selenideSodium tellurideSodium polonide Lithium sulfidePotassium sulfideRubidium sulfideCaesium sulfide Sodium hydrosulfide
Sodium sulfide is a chemical compound with the formula Na2S. More common is its hydrate Na2S·9H2O. Both the anhydrous and the hydrated salts are colorless solids, although technical grades of sodium sulfide are generally yellow to brick red owing to the presence of polysulfides. It is commonly supplied as a crystalline mass, in flake form, or as a fused solid. They are water-soluble, giving strongly alkaline solutions. When exposed to moisture, Na2S is immediately hydrolyzed to give sodium hydrosulfide. Sodium sulfide has an unpleasant rotten egg smell due to the formation of hydrogen sulfide by hydrolysis in moist air.
Some commercial samples are described 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. These grades of sodium sulfide are often marketed as “sodium sulfide flakes”. These samples consist of NaSH, NaOH, and water.
The structures of sodium sulfides have been determined by X-ray crystallography. The nonahydrate features S2- hydrogen-bonded to 12 water molecules.[2] The pentahydrate consists of S2- centers bound to Na+ and encased by an array of hydrogen bonds.[3] Anhydrous Na2S, which is rarely encountered, adopts the antifluorite structure,[4][5] which means that the Na+ centers occupy sites of the fluoride in the CaF2 framework, and the larger S2− occupy the sites for Ca2+.
Industrially Na2S is produced by carbothermic reduction of sodium sulfate often using coal:[6]
Na2SO4 + 2 C → Na2S + 2 CO2
In the laboratory, the salt can be prepared by reduction of sulfur with sodium in anhydrous ammonia, or by sodium in dry THF with a catalytic amount of naphthalene (forming sodium naphthalenide):[7]
2 Na + S → Na2S
The sulfide ion in sulfide salts such as sodium sulfide can incorporate a proton into the salt by protonation:
S2− + H+ → SH−
Because of this capture of the proton ( H+), sodium sulfide has basic character. Sodium sulfide is strongly basic, able to absorb two protons. Its conjugate acid is sodium hydrosulfide (SH−). An aqueous solution contains a significant portion of sulfide ions that are singly protonated.
S2− + H2O ↽ − ⇀ {displaystyle {ce {<=>>}}} SH− + OH− SH− + H2O ↽ − ⇀ {displaystyle {ce {<<=>}}} H2S + OH−
Sodium sulfide is unstable in the presence of water due to the gradual loss of hydrogen sulfide into the atmosphere.
When heated with oxygen and carbon dioxide, sodium sulfide can oxidize to sodium carbonate and sulfur dioxide:
2 Na2S + 3 O2 + 2 CO2 → 2 Na2CO3 + 2 SO2
Oxidation with hydrogen peroxide gives sodium sulfate:[8]
Na2S + 4 H2O2 → 4 H2O + Na2SO4
Upon treatment with sulfur, sodium polysulfides are formed:
2 Na2S + S8 → 2 Na2S5
In terms of its dominant use, sodium sulfide is primarily used in the kraft process in the pulp and paper industry. It aids in the delignification process, affording cellulose, which is the main component of paper.
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 the textile industry as a bleaching agent, for desulfurising and as a dechlorinating agent; and in the 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 during leather processing, as an unhairing agent in the liming operation.
Alkylation of sodium sulfide give thioethers:
Na2S + 2 RX → R2S + 2 NaX
Even aryl halides participate in this reaction.[9] By a broadly similar process sodium sulfide can react with alkenes in the thiol-ene reaction to give thioethers. Sodium sulfide can be used as nucleophile in Sandmeyer type reactions.[10]
Aqueous solution of sodium sulfide will reduce nitro groups to amine. This conversion is applied to production of some azo dyes since other reducible groups, e.g. azo group, remain intact.[11] The reduction of nitro aromatic compounds to amines using sodium sulfide is known as the Zinin reaction in honor of its discoverer.[12] Hydrated sodium sulfide reduces 1,3-dinitrobenzene derivatives to the 3-nitroanilines.[13]
Sulfide has also been employed in photocatalytic applications.[14]
Consisting of the equivalent of sodium hydroxide, sodium sulfide is strongly alkaline and can cause chemical burns. It reacts rapidly with acids to produce hydrogen sulfide, a gas which is both highly toxic and potentially explosive. Sodium sulfide hydrolyses in water to form smaller amounts of hydrogen sulfide which also makes it very toxic to aquatic life.
