|Jmol-3D images||Image 1|
|Molar mass||174.107 g/mol (anhydrous)
210.146 g/mol (dihydrate)
|Appearance||white to grayish crystalline powder
light-lemon colored flakes
|Odor||faint sulfur odor|
|Density||2.38 g/cm3 (anhydrous)
1.58 g/cm3 (dihydrate)
52 °C, 325 K, 126 °F
|Solubility in water||18.2 g/100 mL (anhydrous, 20 °C)
21.9 g/100 mL (Dihydrate, 20 °C)
|Solubility||slightly soluble in alcohol|
|EU classification||Harmful (Xn)|
|R-phrases||R7, R22, R31|
|S-phrases||(S2), S7/8, S26, S28, S43|
|Flash point||100 °C|
|Other anions||Sodium sulfite
|Related compounds||Sodium thiosulfate
| (what is: / ?)
Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa)
Sodium dithionite (also known as sodium hydrosulfite) is a white crystalline powder with a weak sulfurous odor. It is a sodium salt of dithionous acid. 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:
- 2 NaHSO3 + Zn → Na2S2O4 + Zn(OH)2
Properties and reactions 
Sodium dithionite is stable when dry, but is slowly oxidized by air when in solution. Even with the absence of air, solutions of sodium dithionite deteriorate due to the following reaction:
- 2 S2O42- + H2O → S2O32- + 2 HSO3-
Thus solutions of sodium dithionite cannot be stored for a long period of time.
Anhydrous is a monoclinic crystal with slightly sulfuric odor. It is soluble in water and slightly soluble in ethanol. Dihydrate is a columnar crystal, and it is so unstable that it easily dehydrates to anhydrous and is easily oxidized by oxygen in the air.
Anhydrous gradually decomposes to sodium sulfate and sulfur dioxide above 90 °C in the air. In absence of air, it decomposes quickly above 150 °C to sodium sulfite, sodium thiosulfate, sulfur dioxide and trace amount of sulfur.
Powdered anhydrous sodium dithionite with a small amount of water may ignite in air by the heat of decomposition. In absence of air (oxygen), it only decomposes slowly.
An aqueous solution of sodium dithionite is acidic and decomposes to sodium thiosulfate and sodium bisulfite. The reaction rate increases with increasing temperature. In addition, the rate is higher under stronger acidity.
- 2 Na2S2O4 + H2O → Na2S2O3 + 2 NaHSO3
- Na2S2O4 + O2 + H2O → NaHSO4 + NaHSO3
Sodium bisulfate and sodium bisulfite decrease the pH and therefore accelerate the reaction. Sulfur dioxide is formed under strongly acidic conditions.
- 2 H2S2O4 → 3 SO2 + S + 2 H2O
- 3 H2S2O4 → 5 SO2 + H2S + 2 H2O
- 3 Na2S2O4 + 6 NaOH → 5 Na2SO3 + Na2S + 3 H2O
Raman spectroscopy and single-crystal X-ray diffraction studies of sodium dithionite in the solid state reveal that sodium dithionite exists in different forms, such as tierld. In one anhydrous form, the dithionite ion has C
2 geometry, almost eclipsed with a 16° O-S-S-O torsional angle. In the dihydrated form (Na2S2O4.2H2O), the dithionite anion has a shorter S-S bond length and a gauche 56° O-S-S-O torsional angle.
The weak S-S bond causes the dithionite anion to dissociate into the [SO2]- radical anion in aqueous solution, which has been confirmed by ESR spectroscopy. It is also observed that 35S undergoes rapid exchange between S2O42- and SO2 in neutral or acidic solution, consistent with the weak S-S bond in the anion. 
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 dyeing 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, Leather, Chrome Tanning agent and kaolin clay.
- 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 LD50 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). 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.
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.
It can be used as a developer, but it is a very uncommon choice. It is prone to reduce film speed and, if improperly used, quickly fogs the image.
See also 
- Pratt, L. A. (1924). "The Manufacture of Sodium Hyposulfite". Industrial & Engineering Chemistry 16 (7): 676–677. doi:10.1021/ie50175a006.
- Catherine E. Housecroft; Alan G. Sharpe (2008). "Chapter 16: The group 16 elements". Inorganic Chemistry, 3rd Edition. Pearson. p. 520. ISBN 978-0-13-175553-6.
- Weinrach, J. B.; Meyer, D. R.; Guy, J. T.; Michalski, P. E.; Carter, K. L.; Grubisha, D. S.; Bennett, D. W. (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 (3): 291–301. doi:10.1007/BF01199531.
- Herman Harry Szmant (1989). Organic building blocks of the chemical industry. John Wiley and Sons. p. 113. ISBN 0-471-85545-6.
- S.G. Mayhew. Eur. J. Biochem. 85, 535-547 (1978)