Sodium sulfite

Sodium sulphite
anhydrous
hydrate
Names
IUPAC name
Sodium sulfite
Other names
Hypo clear(photography)
E221
Identifiers
7757-83-7 YesY
ChEBI CHEBI:86477 N
ChemSpider 22845 YesY
EC Number 231-821-4
Jmol 3D model Interactive image
PubChem 24437
RTECS number WE2150000
UNII VTK01UQK3G YesY
Properties
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(separate (substances) into constituent elements)
27.0 g/100 mL water (20 °C)
Solubility soluble in glycerol
insoluble in ammonia, chlorine
log P -4
Acidity (pKa) ~9 (heptahydrate)
1.565
Structure
hexagonal (anhydrous)
monoclinic (heptahydrate)
Hazards
Safety data sheet ICSC 1200
NFPA 704
Flammability code 0: Will not burn. E.g., water Health code 2: Intense or continued but not chronic exposure could cause temporary incapacitation or possible residual injury. E.g., chloroform Reactivity code 0: Normally stable, even under fire exposure conditions, and is not reactive with water. E.g., liquid nitrogen Special hazards (white): no codeNFPA 704 four-colored diamond
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 YesYN ?)
Infobox references

Sodium sulfite (sodium sulphite) is a soluble sodium salt of sulfurous acid (sulfite) with the chemical formula Na2SO3. It is a product of sulfur dioxide scrubbing, a part of the flue-gas desulfurization process. 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.

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 sulfitization 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

  1. Weil, Edward D.; Sandler, Stanley R. (1999). "Sulfur Compounds". In Kroschwitz, Jacqueline I. Kirk-Othmer Concise Encylclopedia of Chemical Technology (4th ed.). New York: John Wiley & Sons, Inc. p. 1937. ISBN 978-0471419617.
  2. {http://www.gewater.com/handbook/boiler_water_systems/ch_11_preboiler.jsp}
  3. Merck Index of Chemicals and Drugs, 9th ed. monograph 8451
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