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Sodium sulphite

anhydrous hydrate
Names
IUPAC name Sodium sulfite
Other names Hypo clear (photography) E221
Identifiers
CAS Number	7757-83-7 Y
3D model (JSmol)	Interactive image
ChEBI	CHEBI:86477 N
ChemSpider	22845 Y
ECHA InfoCard	100.028.929
EC Number	231-821-4
E number	E221 (preservatives)
PubChem CID	24437
RTECS number	WE2150000
UNII	VTK01UQK3G Y
InChI InChI=1S/2Na.H2O3S/c;;1-4(2)3/h;;(H2,1,2,3)/q2*+1;/p-2 Y Key: GEHJYWRUCIMESM-UHFFFAOYSA-L Y InChI=1/2Na.H2O3S/c;;1-4(2)3/h;;(H2,1,2,3)/q2*+1;/p-2 Key: GEHJYWRUCIMESM-NUQVWONBAK
SMILES [O-]S(=O)[O-].[Na+].[Na+]
Properties
Chemical formula	Na2SO3
Molar mass	126.043 g/mol
Appearance	White solid
Odor	Odorless
Density	2.633 g/cm3 (anhydrous) 1.561 g/cm3 (heptahydrate)
Melting point	33.4 °C (92.1 °F; 306.5 K) (dehydration of heptahydrate) 500 °C (anhydrous)
Boiling point	Decomposes
Solubility in water	27.0 g/100 mL water (20 °C)
Solubility	Soluble in glycerol Insoluble in ammonia, chlorine
log P	−4
Acidity (pKa)	~9 (heptahydrate)
Refractive index (nD)	1.565
Structure
Crystal structure	Hexagonal (anhydrous) Monoclinic (heptahydrate)
Hazards
Safety data sheet	ICSC 1200
NFPA 704	0 2 0
Flash point	Non-flammable
Related compounds
Other anions	Sodium selenite
Other cations	Potassium sulfite
Related compounds	Sodium bisulfite Sodium metabisulfite Sodium sulfate
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
N verify (what is YN ?)
Infobox references

Sodium sulfite (sodium sulphite) is a soluble sodium salt of sulfurous acid (sulfite) with the chemical formula Na₂SO₃. It is also used as a preservative to prevent dried fruit from discoloring, and for preserving meats, and is used in the same way as sodium thiosulfate to convert elemental halogens to their respective hydrohalic acids, in photography and for reducing chlorine levels in pools. In boiler systems, sulfite and bisulfite are the most commonly employed oxygen scavengers used to prevent pitting corrosion. Sodium sulfite is also a byproduct of sulfur dioxide scrubbing, a part of the flue-gas desulfurization process.

Preparation

Sodium sulfite can be prepared in lab by reacting sodium hydroxide solution with sulfur dioxide gas:

2 NaOH + SO₂ → Na₂SO₃ + H₂O

Evolution of SO₂ by adding few drops of concentrated hydrochloric acid will indicate if sodium hydroxide is nearly gone, turned to aqueous sodium sulfite:

Na₂SO₃ + 2 HCl → 2 NaCl + SO₂ + H₂O

Sodium sulfite is made industrially by reacting sulfur dioxide with a solution of sodium carbonate. The initial combination generates sodium bisulfite (NaHSO₃), which is converted to the sulfite by reaction with sodium hydroxide or sodium carbonate.^[1]

The overall reaction is:

SO₂ + Na₂CO₃ → Na₂SO₃ + CO₂

Applications

Sodium sulfite is primarily used in the pulp and paper industry. It is used in water treatment as an oxygen scavenger agent, to treat water being fed to steam boilers to avoid corrosion problems,^[2] in the photographic industry to protect developer solutions from oxidation and (as hypo clear solution) to wash fixer (sodium thiosulfate) from film and photo-paper emulsions, in the textile industry as a bleaching, desulfurizing and dechlorinating agent and in the leather trade for the sulphonation of tanning extracts. It is used in the purification of TNT for military use. It is used in chemical manufacturing as a sulfonation and sulfomethylation agent. It is used in the production of sodium thiosulfate. It is used in other applications, including froth flotation of ores, oil recovery, food preservatives, and making dyes.

Reactions

Sodium sulfite forms a bisulfite adduct with aldehydes, and with ketones forms a sulfonic acid. It is used to purify or isolate aldehydes and ketones.

Descriptive chemistry

Sodium sulfite is decomposed by even weak acids, giving up sulfur dioxide gas.

Na₂SO₃ + 2 H⁺ → 2 Na⁺ + H₂O + SO₂

A saturated aqueous solution has pH of ~9. Solutions exposed to air are eventually oxidized to sodium sulfate. If sodium sulfite is allowed to crystallize from aqueous solution at room temperature or below, it does so as a heptahydrate. The heptahydrate crystals effloresce in warm dry air. Heptahydrate crystals also oxidize in air to form the sulfate. The anhydrous form is much more stable against oxidation by air.^[3]

References