Sodium Sulfide 250g

Sodium Sulfide 250g

Sodium Sulfide 250g hydrate

Sodium Sulfide is the chemical compound with the formula Na₂S, or more commonly its hydrate Na₂S·9H₂O. Both the anhydrous and the hydrated salts are colorless solids. 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 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. 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’.

Structure

Na₂S adopts the antifluorite structure, 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: 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):  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 {<=>>}}} 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: 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. 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. Sodium sulfide reduces 1,3-dinitrobenzene derivatives to the 3-nitroanilines. 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. Sulfide has also been employed in photocatalytic applications.

Safety

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