The dithionite is the oxyanion with the formula [S2O4]2−.[1] It is commonly encountered as the salt sodium dithionite. For historical reasons, it is sometimes called hydrosulfite, but it contains no hydrogen and is not a sulfite.[2] The dianion has a steric number of 4 and trigonal pyramidal geometry.
Production and reactions
editIn its main applications, dithionite is generally prepared in situ by reduction of sulfur dioxide by sodium borohydride, described by the following idealized equation:[3]
- NaBH4 + 8 SO2 + 8 NaOH → 4 Na2S2O4 + NaBO2 + 6 H2O
Dithionite is a reducing agent. At pH 7, its reduction potential is −0.66 V vs SHE. Its oxidation occurs with formation of sulfite:[4]
- S
2O2−
4 + 2 H2O → 2 HSO−
3 + 2 e− + 2 H+
Dithionite undergoes acid hydrolytic disproportionation to thiosulfate and bisulfite:[2]
- 2 S
2O2−
4 + H2O → S
2O2−
3 + 2 HSO−
3
It also undergoes alkaline hydrolytic disproportionation to sulfite and sulfide:[2]
- 3 Na2S2O4 + 6 NaOH → 5 Na2SO3 + Na2S + 3 H2O
It is formally derived from dithionous acid (H2S2O4), but this acid does not exist in any practical sense.
Use and occurrence
editSodium dithionite finds widespread use in industry as a reducing agent. It is for example used in bleaching of wood pulp and some dyes.[3]
Chemical analyses
editDithionite is used in conjunction with complexing agents (for example, citric acid) to reduce iron(III) oxy-hydroxide into soluble iron(II) compounds and to remove amorphous iron(III)-bearing mineral phases in soil analyses (selective extraction).
Harmful properties
editThe decomposition of dithionite produces reduced species of sulfur that can be very aggressive for the corrosion of steel and stainless steel. Thiosulfate (S
2O2−
3) is known to induce pitting corrosion, whereas sulfide (S2−, HS−) is responsible for stress corrosion cracking (SCC).
References
edit- ^ International Union of Pure and Applied Chemistry (2005). Nomenclature of Inorganic Chemistry (IUPAC Recommendations 2005). Cambridge (UK): RSC–IUPAC. ISBN 0-85404-438-8. p. 130. Electronic version.
- ^ a b c José Jiménez Barberá; Adolf Metzger; Manfred Wolf (2000). "Sulfites, Thiosulfates, and Dithionites". Ullmann's Encyclopedia of Industrial Chemistry. Weinheim: Wiley-VCH. doi:10.1002/14356007.a25_477. ISBN 978-3527306732.
- ^ a b Wietelmann, Ulrich; Felderhoff, Michael; Rittmeyer, Peter (2016-09-29) [2002]. "Hydrides". Ullmann's Encyclopedia of Industrial Chemistry. Weinheim, Germany: Wiley-VCH Verlag GmbH & Co. KGaA. doi:10.1002/14356007.a13_199.pub2. ISBN 978-3-527-30673-2. OCLC 751968805.
- ^ Mayhew, S. G. (2008). "The Redox Potential of Dithionite and SO2− 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.
Further reading
edit- Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 978-0-08-037941-8.
