Sodium sulfide
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Other names
Disodium sulfide
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Identifiers | |
3D model (
JSmol ) |
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ChEBI | |
ChemSpider | |
ECHA InfoCard
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100.013.829 |
EC Number |
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PubChem CID
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RTECS number
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UNII |
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UN number | 1385 (anhydrous) 1849 (hydrate) |
CompTox Dashboard (EPA)
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Properties | |
Na2S | |
Molar mass | 78.0452 g/mol (anhydrous) 240.18 g/mol (nonahydrate) |
Appearance | colorless, hygroscopic solid |
Odor | none |
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
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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 | |
Other anions
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Sodium oxide Sodium selenide Sodium telluride Sodium polonide |
Other cations
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Lithium sulfide Potassium sulfide Rubidium sulfide Caesium sulfide |
Related compounds
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Sodium hydrosulfide |
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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Sodium sulfide is a
Some commercial samples are 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. These grades of sodium sulfide are often marketed as 'sodium sulfide flakes'.
Structure
Na2S adopts the
which means that the Na+ centers occupy sites of the fluoride in the CaF2 framework, and the larger S2− occupy the sites for Ca2+.Production
Industrially Na2S is produced by
- Na2SO4 + 2 C → Na2S + 2 CO2
In the laboratory, the salt can be prepared by reduction of
- 2 Na + S → Na2S
Reactions with inorganic reagents
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
). 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 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:[6]
- Na2S + 4 H2O2 → 4 H2O + Na2SO4
Upon treatment with sulfur, polysulfides are formed:
- 2 Na2S + S8 → 2 Na2S5
Uses
Sodium sulfide is primarily used in the kraft process in the pulp and paper industry.
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,
Reagent in organic chemistry
Alkylation of sodium sulfide give
- Na2S + 2 RX → R2S + 2 NaX
Even
Sodium sulfide is the active ingredient in
Safety
Like
References
- .
- ^ 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.
- ISBN 0-19-855370-6.
- ISBN 0-12-352651-5.
- )
- S2CID 196805424.
- ^ Hartman, W. W.; Silloway, H. L. (1955). "2-Amino-4-nitrophenol". Organic Syntheses
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: CS1 maint: multiple names: authors list (link); Collected Volumes, vol. 3, p. 82. - .
- .