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Sodium dithionite: Wikis


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Sodium dithionite
Sodium dithionite
Other names D-Ox
Sodium hydrosulfite
Sodium sulfoxylate
Virtex L
CAS number 7775-14-6 Yes check.svgY
PubChem 24489
EC number 231-890-0
RTECS number JP2100000
Molecular formula Na2S2O4
Molar mass 174.107 g/mol
Appearance white to grayish crystalline powder
Density 2.19 g/cm3, solid
Melting point

52 °C (325 K)

Boiling point


Solubility in water very soluble
EU Index 016-028-00-1
EU classification Harmful (Xn)
R-phrases R7, R22, R31
S-phrases (S2), S7/8, S26, S28, S42
NFPA 704
NFPA 704.svg
Flash point 100 °C
200 °C
Related compounds
Other anions Sodium sulfite
Sodium sulfate
Related compounds Sodium thiosulfate
Sodium bisulfite
Sodium metabisulfite
Sodium bisulfate
Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa)
Infobox references

Sodium dithionite (also known as sodium hydrosulfite) is a white crystalline powder with a weak sulfurous odor. Although it is stable under most conditions, it will decompose in hot water and in acid solutions. It can be obtained from sodium bisulfite by the following reaction:[1]

2 NaHSO3 + Zn → Na2S2O4 + Zn(OH)2



Raman spectroscopy and single-crystal X-ray diffraction studies of sodium dithionite in the solid state reveals that it exists in different forms. In one anhydrous form, the dithionite ion has C2 geometry, almost eclipsed with a 16° O-S-S-O torsional angle. In the dihydrated form (Na2S2O4.2H 2O), the dithionite anion has a shorter S-S bond length and a gauche 56° O-S-S-O torsional angle.[2]




This compound is a water-soluble salt, and can be used as a reducing agent in aqueous solutions. It is used as such in some industrial dying processes, where an otherwise water-insoluble dye can be reduced into a water-soluble alkali metal salt. The reduction properties of sodium dithionite also eliminate excess dye, residual oxide, and unintended pigments, thereby improving overall colour quality. Reaction with formaldehyde produces Rongalite, which is used as a bleach, in, for instance, paper pulp, cotton, wool, and kaolin clay.[3]

Na2S2O4 + 2 CH2O → 2 HOCH2SO 2 + 2 Na+

Sodium dithionite can also be used for water treatment, gas purification, cleaning, and stripping. It can also be used in industrial processes as a sulfonating agent or a sodium ion source. In addition to the textile industry, this compound is used in industries concerned with leather, foods, polymers, photography, and many others. Its wide use is attributable to its low toxicity LD 50 at 5 g/kg, and hence its wide range of applications.

Biological sciences

Sodium dithionite is often used in physiology experiments as a means of lowering solutions' redox potential (Eo' -0.66 V vs NHE at pH 7[4]). Potassium ferricyanide is usually used as an oxidizing chemical in such experiments (Eo' ~ 436 mV at pH 7). In addition, sodium dithionite is often used in soil chemistry experiments to determine the amount of iron that is not incorporated in primary silicate minerals. Hence, iron extracted by sodium dithionite is also referred to as "free iron." The strong affinity of the dithionite ion for bi- and trivalent metal cations (M2+, M3+) allows it to enhance the solubility of iron, and therefore dithionite is a useful chelating agent. When you heat the reaction, rainbow colors emerge due to the high transition states of energy.


Sodium dithionite has been used in chemical Enhanced Oil Recovery to stabilize polyacrylamide polymers against radical degradation in the presence of iron. It has also been used in environmental applications to propagate a low Eh front in the subsurface in order to reduce pollutants such as chromium.

See also


  1. ^ Pratt, L. A. (1924). "The Manufacture of Sodium Hyposulfite". Industrial & Engineering Chemistry 16: 676–677. doi:10.1021/ie50175a006.   edit
  2. ^ Weinrach, J. B. (1992). "A structural study of sodium dithionite and its ephemeral dihydrate: A new conformation for the dithionite ion". Journal of Crystallographic and Spectroscopic Research 22: 291–301. doi:10.1007/BF01199531.   edit
  3. ^ Herman Harry Szmant (1989). Organic building blocks of the chemical industry. John Wiley and Sons. p. 113. ISBN 0471855456.  
  4. ^ S.G. Mayhew. Eur. J. Biochem. 85, 535-547 (1978)


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