|Two sodium cations and one dithionate anion|
|Ball-and-stick model of the component ions|
|Jmol 3D model||Interactive image|
|Molar mass||206.106 g/mol|
|Appearance||White crystalline powder|
|Melting point||190 °C (374 °F; 463 K) (decomposes)
52 °C (dihydrate)
|Boiling point||267 °C (513 °F; 540 K) decomposes|
|6.27 g/100 mL (0 °C)
15.12 g/100 mL (20 °C)
64.74 g/100 mL (100 °C)
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
|what is ?)(|
Sodium dithionate Na2S2O6 is an important compound for inorganic chemistry. It is also known under names disodium dithionate, sodium hyposulfate, and sodium metabisulfate. The sulfur can be considered to be in its +5 oxidation state.
It should not be confused with sodium dithionite, Na2S2O4, which is a very different compound, and is a powerful reducing agent with many uses in chemistry and biochemistry. Confusion between dithionate and dithionite is commonly encountered, even in manufacturers' catalogues.
- 2 NaHSO3 + MnO2 → Na2S2O6 + MnO + H2O
3 + 2 Ag+
3 → Na
6 + 2 Ag
- 3 Cl2 + Na2S2O3·5H2O + 6 NaOH → Na2S2O6 + 6 NaCl + 8 H2O
And another method to produce sodium dithionate is treating sodium thiosulfate with sodium hypochlorite solution.
The dithionate ion represents sulfur that is oxidized relative to elemental sulfur, but not totally oxidized. Sulfur can be reduced to sulfide or totally oxidized to sulfate, with numerous intermediate oxidation states in inorganic moieties, as well as organosulfur compounds. Example inorganic ions include sulfite and thiosulfate.
Large beautiful single crystals of (Na
2O) has been grown and studied for pulsed lasing purposes (pico second spectroscopy) with great success by E. Haussühl and cols.
Sodium dithionate is a very stable compound which is not oxidized by permanganate, dichromate or bromine. It can be oxidized to sulfate under strongly oxidizing conditions: these include boiling for one hour with 5 M sulfuric acid with an excess of potassium dichromate, or treating with an excess of hydrogen peroxide then boiling with concentrated hydrochloric acid. The Gibbs free energy change for the oxidation is about −300 kJ/mol. In addition, the (S
6) anion is not a good reducing group. Therefore, it has been used to form single crystals of large cation complexes in high oxidation states without reduction of the metallic complex.
- W. G. Palmer (1954). Experimental Inorganic Chemistry. CUP Archive. pp. 361–365. ISBN 0-521-05902-X.<templatestyles src="Module:Citation/CS1/styles.css"></templatestyles>
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- Haussühl, E.; A. A. Kaminskii (2010). Laser Physics. 15 (5): 714–727. Missing or empty