View source for Sodium dithionite

{{redirect|Sodium hydrosulfite|sodium hydrogen sulfite (NaHSO<sub>3</sub>), sometimes erroneously referred to as sodium hydrosulfite|Sodium bisulfite}}
{{not to be confused with|Sodium hydrosulfide|Sodium dithionate}}
{{redirect|D-Ox|other uses|dox (disambiguation)}}
{{chembox
|Verifiedfields = changed
|Watchedfields = changed
|verifiedrevid = 448784510
|Name = Sodium dithionite
|ImageFile = Sodium-dithionite-2D.png
|ImageName = Sodium dithionite
|ImageFile2 = Sample of sodium dithionite.jpg
|OtherNames =D-Ox, Hydrolin, Reductone<br/>sodium hydrosulfite, sodium sulfoxylate, Sulfoxylate<br/>Vatrolite, Virtex L<br/>Hydrosulfit, Prayon<br/>Blankit, Albite A, Konite<br/>Zepar, Burmol, Arostit 
|Section1={{Chembox Identifiers
|CASNo = 7775-14-6
|CASNo_Ref = {{cascite|correct|CAS}}
|UNII_Ref = {{fdacite|correct|FDA}}
|UNII = 2K5B8F6ES1
|PubChem = 24489
|ChEBI_Ref = {{ebicite|changed|EBI}}
|ChEBI = 66870
|RTECS = JP2100000
|ChemSpiderID = 22897
|EINECS = 231-890-0
|StdInChI = 1S/2Na.H2O4S2/c;;1-5(2)6(3)4/h;;(H,1,2)(H,3,4)/q2*+1;/p-2
|StdInChIKey = JVBXVOWTABLYPX-UHFFFAOYSA-L
|SMILES = [O-]S(=O)S(=O)[O-].[Na+].[Na+]
|UNNumber = 1384
}}
|Section2={{Chembox Properties
|Formula = Na<sub>2</sub>S<sub>2</sub>O<sub>4</sub>
|MolarMass = 174.107 g/mol (anhydrous) <br/> 210.146 g/mol (dihydrate)
|Appearance = white to grayish crystalline powder <br/> light-lemon colored flakes
|Odor = faint sulfur odor
|Density = 2.38 g/cm<sup>3</sup> (anhydrous) <br/> 1.58 g/cm<sup>3</sup> (dihydrate)
|Solubility = 18.2 g/100 mL (anhydrous, 20 °C) <br/> 21.9 g/100 mL (Dihydrate, 20 °C)
|SolubleOther = slightly soluble in [[ethanol|alcohol]]
|MeltingPtC = 52
|BoilingPt = Decomposes
}}
|Section3={{Chembox Hazards
|ExternalSDS = 
|GHSPictograms = {{GHS02}}{{GHS07}}
|GHSSignalWord = Danger
|HPhrases = {{H-phrases|251|302}}
|PPhrases = {{P-phrases|235+410|264|270|280|301+312|330|407|413|420|501}}
|NFPA-H = 2
|NFPA-F = 3
|NFPA-R = 1
|FlashPtC = 100
|AutoignitionPtC = 200
}}
|Section4={{Chembox Related
|OtherAnions = [[Sodium sulfite]]<br/>[[Sodium sulfate]]
|OtherCompounds = [[Sodium thiosulfate]]<br/>[[Sodium bisulfite]]<br/>[[Sodium metabisulfite]]<br/>[[Sodium bisulfate]]
}}
}}

'''Sodium dithionite''' (also known as '''sodium hydrosulfite''') is a white [[crystal]]line powder with a sulfurous odor. Although it is stable in dry air, it [[Chemical decomposition|decomposes]] in hot [[water]] and in [[acid]] [[Solution (chemistry)|solution]]s.

==Structure==
[[image:Sodium-dithionite-xtal-1992-3D-balls.png|left|220px|thumb|Packing of sodium dithionite ions in a crystal, showing the saw-horse geometry of the dianion.  Color code: red = O, yellow = S.]]
The structure has been examined by [[Raman spectroscopy]] and [[X-ray crystallography]]. The dithionite dianion has C{{su|b=2}} [[symmetry group|symmetry]], with almost eclipsed with a 16° O-S-S-O torsional angle. In the [[Hydrate|dihydrated]] form ({{chem|Na|2|S|2|O|4|·2H|2|O}}), the dithionite anion has [[Alkane stereochemistry|gauche]] 56° O-S-S-O torsional angle.<ref>{{cite journal| first1 = J. B.| first2 = D. R.| first3 = J. T.| first4 = P. E.| first5 = K. L.| last1 = Weinrach| first6 = D. S.| first7 = D. W.| title = A structural study of sodium dithionite and its ephemeral dihydrate: A new conformation for the dithionite ion| journal = Journal of Crystallographic and Spectroscopic Research| volume = 22| issue = 3| pages = 291–301| year = 1992| doi = 10.1007/BF01199531| last2 = Meyer| last3 = Guy| last4 = Michalski| last5 = Carter| last6 = Grubisha| last7 = Bennett| s2cid = 97124638}}</ref>

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.<ref>{{Greenwood&Earnshaw2nd}}</ref>  Because this bond is fragile, the dithionite anion dissociates in solution into the [SO<sub>2</sub>]<sup>−</sup> radicals, as has been confirmed by [[EPR spectroscopy]]. It is also observed that <sup>35</sup>S undergoes rapid exchange between S<sub>2</sub>O<sub>4</sub><sup>2−</sup> and SO<sub>2</sub> in neutral or acidic solution, consistent with the weak S-S bond in the anion.<ref name="InorgChem" />

==Preparation==
Sodium dithionite is produced industrially by reduction of [[sulfur dioxide]]. Approximately 300,000 tons were produced in 1990.<ref name=Ullmann>{{cite book |author1=José Jiménez Barberá |author2=Adolf Metzger |author3=Manfred Wolf |chapter=Sulfites, Thiosulfates, and Dithionites |title=Ullmann's Encyclopedia of Industrial Chemistry |date=15 June 2000 |publisher=Wiley Online Library |doi=10.1002/14356007.a25_477|isbn=978-3527306732 }}</ref> The route using zinc powder is a two-step process:
:2{{nbsp}}SO<sub>2</sub>  +  Zn  →  ZnS<sub>2</sub>O<sub>4</sub>
:ZnS<sub>2</sub>O<sub>4</sub>  +  2{{nbsp}}NaOH  →  Na<sub>2</sub>S<sub>2</sub>O<sub>4</sub>  +  Zn(OH)<sub>2</sub>

The [[sodium borohydride]] method obeys the following stoichiometry:
:NaBH<sub>4</sub>  +  8{{nbsp}}NaOH  +  8{{nbsp}}SO<sub>2</sub>  →  4{{nbsp}}Na<sub>2</sub>S<sub>2</sub>O<sub>4</sub>  +  NaBO<sub>2</sub>   +  6{{nbsp}}H<sub>2</sub>O
Each equivalent of H<sup>−</sup> 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 S<sub>2</sub>O<sub>4</sub><sup>2−</sup> + H<sub>2</sub>O → S<sub>2</sub>O<sub>3</sub><sup>2−</sup> + 2 HSO<sub>3</sub><sup>−</sup>
This behavior is consistent with the instability of [[dithionous acid]].  Thus, solutions of sodium dithionite cannot be stored for a long period of time.<ref name="InorgChem">{{cite book
| title = Inorganic Chemistry, 3rd Edition
| chapter = Chapter 16: The group 16 elements
| author1 = Catherine E. Housecroft
| author2 = Alan G. Sharpe
| publisher = Pearson
| year = 2008
| isbn = 978-0-13-175553-6
| page = 520
}}</ref>

Anhydrous sodium dithionite decomposes to [[sodium sulfate]] and [[sulfur dioxide]] above 90&nbsp;°C in the air. In absence of air, it decomposes quickly above 150&nbsp;°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 compared to the [[standard hydrogen electrode|normal hydrogen electrode]].  Redox occurs with formation of bisulfite:<ref>{{cite journal|author=Mayhew, S. G.|title=The Redox Potential of Dithionite and SO−2 from Equilibrium Reactions with Flavodoxins, Methyl Viologen and Hydrogen plus Hydrogenase|journal=European Journal of Biochemistry|year=2008|volume=85|issue=2|pages=535–547|doi=10.1111/j.1432-1033.1978.tb12269.x|pmid=648533|doi-access=free}}</ref>
:S<sub>2</sub>O<sub>4</sub><sup>2-</sup>  +  2 H<sub>2</sub>O   →   2 HSO<sub>3</sub><sup>−</sup>  +  2 e<sup>−</sup> +  2 H<sup>+</sup>

Sodium dithionite reacts with oxygen:
:Na<sub>2</sub>S<sub>2</sub>O<sub>4</sub> + O<sub>2</sub> + H<sub>2</sub>O → NaHSO<sub>4</sub> + NaHSO<sub>3</sub>
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&nbsp;°C.<ref>J. Org. Chem., 1980, 45 (21), pp 4126–4129, http://pubs.acs.org/doi/abs/10.1021/jo01309a011</ref><ref>{{cite web|url=https://patents.google.com/patent/US5270058A/en|title=Aldehyde sulfoxylate systemic fungicides|website=google.com|access-date=27 April 2018|url-status=live|archive-url=https://web.archive.org/web/20180427175353/https://patents.google.com/patent/US5270058A/en|archive-date=27 April 2018}}</ref>  Some ketones are also reduced under similar conditions.

==Uses==

===Industry===
Sodium dithionite is used as a water-soluble [[reducing agent]] in some industrial dyeing processes.  In the case of [[sulfur dye]]s and [[vat dye]]s, an otherwise water-insoluble dye can be reduced into its water-soluble [[alkali metal]] [[leuco salt]].  [[Indigo dye]] is sometimes processed in this way.<ref>{{cite journal|last1=Božič|first1=Mojca|last2=Kokol|first2=Vanja|title=Ecological alternatives to the reduction and oxidation processes in dyeing with vat and sulphur dyes|journal=Dyes and Pigments|date=2008|volume=76|issue=2|pages=299–309|doi=10.1016/j.dyepig.2006.05.041}}</ref>

===Domestic and hobby uses===
Sodium dithionite can also be used for water treatment, aquarium water conditioners, gas purification, cleaning, and stripping. 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.

It is the active ingredient in "Iron Out Rust Stain Remover", a commercial rust product.<ref>{{cite web | url=https://www.popularmechanics.com/home/tools/g2312/best-liquid-rust-removers-test/#product-13a4a293-8e39-4bbf-93b1-74a14dc07cfa-anchor | title=The Best Rust Removers for Restoring Every Surface | date=23 March 2023 }}</ref>

===Laboratory===
Sodium dithionite is often used in [[physiology]] experiments as a means of lowering solutions' [[redox potential]] (E<sup>o'</sup>  -0.66 V vs [[Standard hydrogen electrode|SHE]] at pH 7).<ref>{{cite journal|last1=MAYHEW|first1=Stephen G.|title=The Redox Potential of Dithionite and SO-2 from Equilibrium Reactions with Flavodoxins, Methyl Viologen and Hydrogen plus Hydrogenase|journal=European Journal of Biochemistry|volume=85|issue=2|year=1978|pages=535–547|issn=0014-2956|doi=10.1111/j.1432-1033.1978.tb12269.x|pmid=648533|doi-access=free}}</ref>  [[Potassium ferricyanide]] is usually used as an oxidizing chemical in such experiments (E<sup>o'</sup> ~ .436 V at pH 7). In addition, sodium dithionite is often used in [[soil science|soil chemistry]] experiments to determine the amount of [[iron]] that is not incorporated in primary [[silicate mineral]]s. Hence, iron extracted by sodium dithionite is also referred to as "free iron." 

Aqueous solutions of sodium dithionite were once used to produce '[[Louis Fieser|Fieser's]] solution' for the removal of oxygen from a gas stream.<ref>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}}</ref>  [[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 reaction|substitution]] using sodium dithionite to introduce the [[thiol]] [[functional group]].<ref>{{cite encyclopedia|title = 1-Hydroxypyridine-2(1''H'')-thione|first1 = David W.|last1 = Knight|first2 = Jens|last2 = Hartung|date = 15 September 2006|doi = 10.1002/047084289X.rh067.pub2|publisher = [[John Wiley & Sons]]|encyclopedia = [[Encyclopedia of Reagents for Organic Synthesis]]|chapter = 1-Hydroxypyridine-2(1H)-thione|isbn = 978-0471936237}}</ref>

===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''.<ref>{{cite web |title=Kodak Direct Positive Film 5246 |url=https://125px.com/docs/techpubs/kodak/j6.pdf |website=125px.com |publisher=Kodak |access-date=6 November 2019}}</ref>

==Safety==
The wide use of sodium dithionite is attributable in part to its low toxicity {{LD50}} at 2.5 g/kg (rats, oral).<ref name=Ullmann/>

==See also==
*[[Dithionite]]

==References==
{{Reflist}}

==External links==
*Sodium dithionite - ipcs inchem<ref name="inchem">{{cite web |title=Sodium dithionite - ipcs inchem |url=http://www.inchem.org/documents/sids/sids/7775146.pdf |website=www.inchem.org |access-date=15 June 2018 |location=Berliln, Germany |date=2004 |archive-date=17 April 2018 |archive-url=https://web.archive.org/web/20180417180337/http://www.inchem.org/documents/sids/sids/7775146.pdf |url-status=dead }}</ref>

{{Sodium compounds}}

[[Category:Bleaches]]
[[Category:Dithionites]]
[[Category:Sodium compounds]]
[[Category:Reducing agents]]