Sodium dithionite

Sodium dithionite (also known as sodium hydrosulfite) is a white crystalline powder with a weak sulfurous odor. Although it is stable in the absence of air, it decomposes in hot water and in acid solutions.

"Sodium hydrosulfite" redirects here. For sodium hydrogen sulfite (NaHSO3), sometimes erroneously referred to as sodium hydrosulfite, see Sodium bisulfite.
Not to be confused with Sodium dithionate.
Sodium dithionite
Names
Other names
D-Ox, Hydrolin, Reductone
sodium hydrosulfite, sodium sulfoxylate, Sulfoxylate
Vatrolite, Virtex L
Hydrosulfit, Prayon
Blankit, Albite A, Konite
Zepar, Burmol, Arostit
Identifiers
3D model (JSmol)
ChEBI
ChemSpider
ECHA InfoCard 100.028.991
EC Number
  • 231-890-0
RTECS number
  • JP2100000
UNII
UN number 1384
Properties
Na2S2O4
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)
Melting point 52 °C (126 °F; 325 K)
Boiling point Decomposes
18.2 g/100 mL (anhydrous, 20 °C)
21.9 g/100 mL (Dihydrate, 20 °C)
Solubility slightly soluble in alcohol
Hazards
Harmful (Xn)
R-phrases (outdated) R7, R22, R31
S-phrases (outdated) (S2), S7/8, S26, S28, S43
NFPA 704 (fire diamond)
Flammability code 3: Liquids and solids that can be ignited under almost all ambient temperature conditions. Flash point between 23 and 38 °C (73 and 100 °F). E.g. gasolineHealth code 2: Intense or continued but not chronic exposure could cause temporary incapacitation or possible residual injury. E.g. chloroformReactivity code 1: Normally stable, but can become unstable at elevated temperatures and pressures. E.g. calciumSpecial hazards (white): no code
3
2
1
Flash point 100 °C (212 °F; 373 K)
200 °C (392 °F; 473 K)
Related compounds
Other anions
 Sodium sulfite
Sodium sulfate
Related compounds
 Sodium thiosulfate
Sodium bisulfite
Sodium metabisulfite
Sodium bisulfate
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

Structure

The structure has been examined by Raman spectroscopy and single-crystal X-ray diffraction. The dithionite dianion has C
2 symmetry, with almost eclipsed with a 16° O-S-S-O torsional angle. In the dihydrated form (Na
2S
2O
4·2H
2O), the dithionite anion has gauche 56° O-S-S-O torsional angle.[1]

A weak S-S bond is indicated by the S-S distance of 239 pm, which is elongated by ca. 30 pm relative to a typical S-S bond.[2] Because this bond is fragile, the dithionite anion dissociates in solution into the [SO2]− radicals, as has been confirmed by EPR 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.[3]

Preparation

Sodium dithionite is produced industrially by reduction of sulfur dioxide. Approximately 300,000 tons were produced in 1990.[4] The route using zinc powder is a two-step process:

2 SO2 + Zn → ZnS2O4
ZnS2O4 + 2 NaOH → Na2S2O4 + Zn(OH)2

The sodium borohydride method obeys the following stoichiometry:

NaBH4 + 8 NaOH + 8 SO2 → 4 Na2S2O4 + NaBO2 + 6 H2O

Each equivalent of H- reduces two equivalents of sulfur dioxide. Formate has also been used as the reductant.

Properties and reactions

Hydrolysis

Sodium dithionite is stable when dry, but aqueous solutions deteriorate due to the following reaction:

2 S2O42− + H2O → S2O32− + 2 HSO3−

This behavior is consistent with the instability of dithionous acid. Thus, solutions of sodium dithionite cannot be stored for a long period of time.[3]

Anhydrous sodium dithionite 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.

Redox reactions

Sodium dithionite is a reducing agent. At pH=7, the potential is -0.66 V vs NHE. Redox occurs with formation of sulfite:[5]

S2O42- + 2 H2O → 2 HSO3− + 2 e− + 2 H+

Sodium dithionite reacts with oxygen:

Na2S2O4 + O2 + H2O → NaHSO4 + NaHSO3

These reactions exhibit complex pH-dependent equilibria involving bisulfite, thiosulfate, and sulfur dioxide.

With organic carbonyls

In the presence of aldehydes, sodium dithionite reacts either to form α-hydroxy-sulfinates at room temperature or to reduce the aldehyde to the corresponding alcohol above a temperature of 85 °C.[6][7] Some ketones are also reduced under similar conditions.

Applications

Industry

Being water-soluble, sodium dithionite is used as a reducing agent in some industrial dyeing processes. In the case of sulfur dyes and vat dyes, an otherwise water-insoluble dye can be reduced into a water-soluble alkali metal salt (e.g. indigo dye).[8]

Sodium dithionite can also be used for water treatment, gas purification, cleaning, and stripping. It has also been applied as a sulfonating agent. In addition to the textile industry, this compound is used in industries concerned with leather, foods, polymers, photography, and many others, often as a decolourising agent. It is even used domestically as a decoloring agent for white laundry, when it has been accidentally stained by way of a dyed item slipping into the high temperature washing cycle. It is usually available in 5 gram sachets termed hydrosulfite after the antiquated name of the salt.

Laboratory

Sodium dithionite is often used in physiology experiments as a means of lowering solutions' redox potential (Eo' -0.66 V vs SHE at pH 7).[9] Potassium ferricyanide is usually used as an oxidizing chemical in such experiments (Eo' ~ .436 V 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.

Aqueous solutions of sodium dithionite were once used to produce "Fieser's solution' for the removal of oxygen from a gas stream.[10] Pyrithione can be prepared in a two-step synthesis from 2-bromopyridine by oxidation to the N-oxide with a suitable peracid followed by substitution using sodium dithionite to introduce the thiol functional group.[11]

Photography

It is used in Kodak fogging developer, FD-70. This is used in the second step in processing black and white positive images, for making slides. It is part of the Kodak Direct Positive Film Developing Outfit.[12]

Safety

The wide use of sodium dithionite is attributable in part to its low toxicity LD50 at 2.5 g/kg (rats, oral).[4]

See also

References
  1. 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. S2CID 97124638.
  2. Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 978-0-08-037941-8.
  3. 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.
  4. José Jiménez Barberá; Adolf Metzger; Manfred Wolf (15 June 2000). "Sulfites, Thiosulfates, and Dithionites". Ullmann's Encyclopedia of Industrial Chemistry. Wiley Online Library. doi:10.1002/14356007.a25_477. ISBN 978-3527306732.
  5. Mayhew, S. G. (2008). "The Redox Potential of Dithionite and SO−2 from Equilibrium Reactions with Flavodoxins, Methyl Viologen and Hydrogen plus Hydrogenase". European Journal of Biochemistry. 85 (2): 535–547. doi:10.1111/j.1432-1033.1978.tb12269.x. PMID 648533.
  6. J. Org. Chem., 1980, 45 (21), pp 4126–4129, http://pubs.acs.org/doi/abs/10.1021/jo01309a011
  7. "Aldehyde sulfoxylate systemic fungicides". google.com. Archived from the original on 27 April 2018. Retrieved 27 April 2018.
  8. Božič, Mojca; Kokol, Vanja (2008). "Ecological alternatives to the reduction and oxidation processes in dyeing with vat and sulphur dyes". Dyes and Pigments. 76 (2): 299–309. doi:10.1016/j.dyepig.2006.05.041.
  9. MAYHEW, Stephen G. (1978). "The Redox Potential of Dithionite and SO-2 from Equilibrium Reactions with Flavodoxins, Methyl Viologen and Hydrogen plus Hydrogenase". European Journal of Biochemistry. 85 (2): 535–547. doi:10.1111/j.1432-1033.1978.tb12269.x. ISSN 0014-2956. PMID 648533.
  10. Kenneth L. Williamson "Reduction of Indigo: Sodium Hydrosulfite as a Reducing Agent" J. Chem. Educ., 1989, volume 66, p 359. doi:10.1021/ed066p359.2
  11. Knight, David W.; Hartung, Jens (15 September 2006). "1-Hydroxypyridine-2(1H)-thione". 1-Hydroxypyridine-2(1H)-thione. Encyclopedia of Reagents for Organic Synthesis. John Wiley & Sons. doi:10.1002/047084289X.rh067.pub2. ISBN 978-0471936237.
  12. "Kodak Direct Positive Film 5246" (PDF). 125px.com. Kodak. Retrieved 6 November 2019.
  • Sodium dithionite - ipcs inchem[1]
  1. "Sodium dithionite - ipcs inchem" (PDF). www.inchem.org. Berliln, Germany. 2004. Retrieved 15 June 2018.
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