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| Names | |
|---|---|
| Other names Disodium sulfide | |
| Identifiers | |
3D model (JSmol) | |
| ChEBI | |
| ChemSpider |
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| ECHA InfoCard | 100.013.829 |
| EC Number |
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| RTECS number |
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| UNII |
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| UN number | 1385 (anhydrous) 1849 (hydrate) |
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| 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 inatmosphericmoisture tohydrogen 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 inether slightly soluble inalcohol[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 | |
Otheranions | Sodium oxide Sodium selenide Sodium telluride Sodium polonide |
Othercations | Lithium sulfide Potassium sulfide Rubidium sulfide Caesium sulfide |
Related compounds | Sodium hydrosulfide |
Except where otherwise noted, data are given for materials in theirstandard state (at 25 °C [77 °F], 100 kPa). | |
Sodium sulfide is achemical compound with theformulaNa2S, or more commonly itshydrate 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 ofpolysulfides. It is commonly supplied as a crystalline mass, in flake form, or as a fused solid. They are water-soluble, giving stronglyalkaline solutions. When exposed to moisture, Na2S immediately hydrates to givesodium hydrosulfide. Sodium sulfide has an unpleasant rotten egg smell due to the hydrolysis tohydrogen sulfide 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 thatx 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 byX-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 theantifluorite 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 bycarbothermic reduction ofsodium sulfate often using coal:[6]
In the laboratory, the salt can be prepared by reduction ofsulfur withsodium in anhydrousammonia, or by sodium in dryTHF with a catalytic amount ofnaphthalene (formingsodium naphthalenide):[7]
The sulfide ion in sulfide salts such as sodium sulfide can incorporate a proton into the salt by protonation:
Because of this capture of theproton ( H+), sodium sulfide has basic character. Sodium sulfide is strongly basic, able to absorb two protons. Itsconjugate acid issodium hydrosulfide (SH−
). An aqueous solution contains a significant portion of sulfide ions that are singly protonated.
Sodium sulfide is unstable in the presence of water due to the gradual loss ofhydrogen sulfide into the atmosphere.
When heated withoxygen andcarbon dioxide, sodium sulfide can oxidize tosodium carbonate andsulfur dioxide:
Oxidation withhydrogen peroxide givessodium sulfate:[8]
Upon treatment withsulfur,sodium polysulfides are formed:
In terms of its dominant use, "sodium sulfide" is primarily used in thekraft process in thepulp 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 givethioethers:
Evenaryl halides participate in this reaction.[9] By a broadly similar process sodium sulfide can react with alkenes in thethiol-ene reaction to give thioethers.Sodium sulfide can be used asnucleophile inSandmeyer type reactions.[10]
Aqueous solution of sodium sulfide will reducenitro groups toamine. This conversion is applied to production of someazo 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 theZinin reaction in honor of its discoverer.[12] Hydrated sodium sulfide reduces1,3-dinitrobenzene derivatives to the3-nitroanilines.[13]
Sulfide has also been employed inphotocatalytic applications.[14]
Consisting of the equivalent ofsodium hydroxide, sodium sulfide is strongly alkaline and can causechemical burns. It reacts rapidly with acids to producehydrogen 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.
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