Dithionite
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
[edit]In 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
[edit]Sodium 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
[edit]Dithionite 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
[edit]The 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.
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