Sodium sulfide


















































































































































Sodium sulfide

Natriumsulfid.jpg

Fluorite-unit-cell-3D-ionic.png
Names
Other names
Disodium sulfide

Identifiers

CAS Number




  • 1313-82-2 ☑Y


  • 1313-84-4 (pentahydrate) ☒N


  • 1313-84-4 (nonahydrate) ☒N



3D model (JSmol)


  • Interactive image


ChEBI


  • CHEBI:76208 ☒N


ChemSpider


  • 14120 ☒N


ECHA InfoCard

100.013.829

EC Number
215-211-5


PubChem CID


  • 237873


RTECS number
WE1905000

UN number
1385 (anhydrous)
1849 (hydrate)




Properties

Chemical formula

Na2S

Molar mass
78.0452 g/mol (anhydrous)
240.18 g/mol (nonahydrate)
Appearance
colorless, hygroscopic solid

Odor
rotten eggs

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)

Solubility in water

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]


Magnetic susceptibility (χ)

−39.0·10−6 cm3/mol
Structure

Crystal structure


Antifluorite (cubic), cF12

Space group

Fm3m, No. 225

Coordination geometry

Tetrahedral (Na+); cubic (S2−)
Hazards

Safety data sheet

ICSC 1047


EU classification (DSD) (outdated)

Corrosive (C)
Dangerous for the environment (N)

R-phrases (outdated)

R31, R34, R50

S-phrases (outdated)

(S1/2), S26, S45, S61

NFPA 704



Flammability code 1: Must be pre-heated before ignition can occur. Flash point over 93 °C (200 °F). E.g., canola oil
Health code 3: Short exposure could cause serious temporary or residual injury. E.g., chlorine gas
Reactivity code 1: Normally stable, but can become unstable at elevated temperatures and pressures. E.g., calcium
Special hazards (white): no code
NFPA 704 four-colored diamond


1


3


1



Autoignition
temperature

> 480 °C (896 °F; 753 K)
Related compounds

Other anions


Sodium oxide
Sodium selenide
Sodium telluride

Other cations


Lithium sulfide
Potassium sulfide

Related compounds


Sodium hydrosulfide

Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).


☒N verify (what is ☑Y☒N ?)

Infobox references



Sodium sulfide is the chemical compound with the formula Na2S, or more commonly its hydrate Na2S·9H2O. Both are colorless water-soluble salts that give strongly alkaline solutions. When exposed to moist air, Na2S and its hydrates emit hydrogen sulfide, which smells like rotten eggs. Some commercial samples are specified as Na2xH2O, where a weight percentage of Na2S is specified. Commonly available grades have around 60% Na2S by weight, which means that x is around 3. Such technical grades of sodium sulfide have a yellow appearance owing to the presence of polysulfides. These grades of sodium sulfide are marketed as 'sodium sulfide flakes'. Although the solid is yellow, solutions of it are colorless.




Contents






  • 1 Structure


  • 2 Production


  • 3 Reactions with inorganic reagents


  • 4 Uses


    • 4.1 Reagent in organic chemistry




  • 5 Safety


  • 6 References





Structure


Na2S adopts the antifluorite structure,[2][3] 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 carbothermic reduction of sodium sulfate often using coal:[4]


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):[5]


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 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 {<=>>}}}{displaystyle {ce {<=>>}}} SH
+ OH












 



 



 



 





(1)










SH
+ H2O
{displaystyle {ce {<<=>}}}{displaystyle {ce {<<=>}}} H2S + OH












 



 



 



 





(2)




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, oil recovery, making dyes, and detergent. It is also used during leather processing, as an unhairing agent in the liming operation.



Reagent in organic chemistry


Alkylation of sodium sulfide give thioethers:


Na2S + 2 RX → R2S + 2 NaX

Even aryl halides participate in this reaction.[7]
Sodium sulfide can be used as nucleophile in Sandmeyer type reactions.[8] Sodium sulfide reduces1,3-dinitrobenzene derivatives to the 3-nitroanilines.[9] Aqueous solution of sodium sulfide can be refluxed with nitro carrying azo dyes dissolved in dioxane and ethanol to selectively reduce the nitro groups to amine; while other reducible groups, e.g. azo group, remain intact.[10] Sulfide has also been employed in photocatalytic applications.[11]



Safety


Like sodium hydroxide, sodium sulfide is strongly alkaline and can cause skin burns. Acids react with it to rapidly produce hydrogen sulfide, which is highly toxic.



References





  1. ^ Kurzin, Alexander V.; Evdokimov, Andrey N.; Golikova, Valerija S.; Pavlova, Olesja S. (June 9, 2010). "Solubility of Sodium Sulfide in Alcohols". J. Chem. Eng. Data. 55 (9): 4080–4081. doi:10.1021/je100276c..mw-parser-output cite.citation{font-style:inherit}.mw-parser-output q{quotes:"""""""'""'"}.mw-parser-output code.cs1-code{color:inherit;background:inherit;border:inherit;padding:inherit}.mw-parser-output .cs1-lock-free a{background:url("//upload.wikimedia.org/wikipedia/commons/thumb/6/65/Lock-green.svg/9px-Lock-green.svg.png")no-repeat;background-position:right .1em center}.mw-parser-output .cs1-lock-limited a,.mw-parser-output .cs1-lock-registration a{background:url("//upload.wikimedia.org/wikipedia/commons/thumb/d/d6/Lock-gray-alt-2.svg/9px-Lock-gray-alt-2.svg.png")no-repeat;background-position:right .1em center}.mw-parser-output .cs1-lock-subscription a{background:url("//upload.wikimedia.org/wikipedia/commons/thumb/a/aa/Lock-red-alt-2.svg/9px-Lock-red-alt-2.svg.png")no-repeat;background-position:right .1em center}.mw-parser-output .cs1-subscription,.mw-parser-output .cs1-registration{color:#555}.mw-parser-output .cs1-subscription span,.mw-parser-output .cs1-registration span{border-bottom:1px dotted;cursor:help}.mw-parser-output .cs1-hidden-error{display:none;font-size:100%}.mw-parser-output .cs1-visible-error{font-size:100%}.mw-parser-output .cs1-subscription,.mw-parser-output .cs1-registration,.mw-parser-output .cs1-format{font-size:95%}.mw-parser-output .cs1-kern-left,.mw-parser-output .cs1-kern-wl-left{padding-left:0.2em}.mw-parser-output .cs1-kern-right,.mw-parser-output .cs1-kern-wl-right{padding-right:0.2em}


  2. ^ 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.


  3. ^ Wells, A.F. (1984) Structural Inorganic Chemistry, Oxford: Clarendon Press.
    ISBN 0-19-855370-6.



  4. ^ Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 2001.
    ISBN 0-12-352651-5.



  5. ^ So, J.-H; Boudjouk, P; Hong, Harry H.; Weber, William P. (1992). Hexamethyldisilathiane. Inorg. Synth. Inorganic Syntheses. 29. p. 30. doi:10.1002/9780470132609.ch11. ISBN 978-0-470-13260-9.


  6. ^ L. Lange, W. Triebel, "Sulfides, Polysulfides, and Sulfanes" in Ullmann's Encyclopedia of Industrial Chemistry 2000, Wiley-VCH, Weinheim. doi:10.1002/14356007.a25_443


  7. ^ Charles C. Price, Gardner W. Stacy "p-Aminophenyldisulfide" Org. Synth. 1948, vol. 28, 14. doi:10.15227/orgsyn.028.0014


  8. ^ Khazaei; et al. (2012). "synthesis of thiophenols". Synthesis Letters - Thieme Chemistry. 23 (13): 1893–1896. doi:10.1055/s-0032-1316557.CS1 maint: Explicit use of et al. (link)


  9. ^ Hartman, W. W.; Silloway, H. L. (1955). "2-Amino-4-nitrophenol". Organic Syntheses.CS1 maint: Multiple names: authors list (link) ; Collective Volume, 3, p. 82


  10. ^ Yu; et al. (2006). "Syntheses of functionalized azobenzenes". Tetrahedron. 62 (44): 10303–10310. doi:10.1016/j.tet.2006.08.069.CS1 maint: Explicit use of et al. (link)


  11. ^ Savateev, A.; Dontsova, D.; Kurpil, B.; Antonietti, M. (June 2017). "Highly crystalline poly(heptazine imides) by mechanochemical synthesis for photooxidation of various organic substrates using an intriguing electron acceptor – Elemental sulfur". Journal of Catalysis. 350: Pages 203–211. doi:10.1016/j.jcat.2017.02.029.CS1 maint: Multiple names: authors list (link)











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