Sodium sulfide

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Sodium sulfide
Names
Other names
Disodium sulfide
Identifiers
3D model (
JSmol
)
ChEBI
ChemSpider
ECHA InfoCard
100.013.829 Edit this at Wikidata
EC Number
  • 215-211-5
RTECS number
  • WE1905000
UNII
UN number 1385 (anhydrous)
1849 (hydrate)
  • InChI=1S/2Na.S/q2*+1;-2 ☒N
    Key: GRVFOGOEDUUMBP-UHFFFAOYSA-N ☒N
  • InChI=1/2Na.S/q2*+1;-2
    Key: GRVFOGOEDUUMBP-UHFFFAOYAP
  • [Na+].[Na+].[S-2]
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
Fm3m, No. 225
Tetrahedral (Na+); cubic (S2−)
Hazards
GHS labelling:
GHS05: Corrosive GHS06: Toxic GHS07: Exclamation mark GHS09: Environmental hazard
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)
NFPA 704 four-colored diamondHealth 3: Short exposure could cause serious temporary or residual injury. E.g. chlorine gasFlammability 1: Must be pre-heated before ignition can occur. Flash point over 93 °C (200 °F). E.g. canola oilInstability 1: Normally stable, but can become unstable at elevated temperatures and pressures. E.g. calciumSpecial hazards (white): no code
3
1
1
> 480 °C (896 °F; 753 K)
Safety data sheet (SDS) ICSC 1047
Related compounds
Other anions
Sodium oxide
Sodium selenide
Sodium telluride
Sodium polonide
Other cations
Lithium sulfide
Potassium sulfide
Rubidium sulfide
Caesium 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).
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Sodium sulfide is a

Na2S, or more commonly its hydrate Na2S·9H2O. Both the anhydrous and the hydrated salts in pure crystalline form are colorless solids, although technical grades of sodium sulfide are generally yellow to brick red owing to the presence of polysulfides and 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 moist air, Na2S and its hydrates emit hydrogen sulfide
, an extremely toxic, flammable and corrosive gas 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. 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

conjugate acid is sodium hydrosulfide
(SH
). An aqueous solution contains a significant portion of sulfide ions that are singly protonated.

S2−
+
H2O SH
+ OH
SH
+ H2O 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:[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] 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
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]

Sodium sulfide is the active ingredient in

ingrown toenail
pain treatment.

Safety

Like

chemical burns. Acids react with it to rapidly produce hydrogen sulfide
, which is highly toxic.

References

  1. .
  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. .
  4. .
  5. ISBN 978-0-470-13260-9. {{cite book}}: |journal= ignored (help
    )
  6. .
  7. ^ Hartman, W. W.; Silloway, H. L. (1955). "2-Amino-4-nitrophenol". Organic Syntheses{{cite journal}}: CS1 maint: multiple names: authors list (link); Collected Volumes, vol. 3, p. 82.
  8. .
  9. .