clearly_not_atara
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thoughts on practical reduction with dithionite
The mechanism of reductions by sodium dithionite turns out to be known; the first step is thermal fragmentation of the molecule to release
sulfoxylate:
S2O4(2-) + heat >> SO2 + SO2(2-)
Sulfoxylate dianion is a strong reducing agent but also a strong base. The pKa2 of the reaction:
SO2(2-) + H+ >> SO2H-
is about 13.5 at 20 C. Correspondingly, reductions with dithionite become much faster above pH 13. So in order to generate the active reducing agent
you need:
* heat
* pH >= 13
This is not achievable with carbonate bases because pKa2 = 10.6 and a solution of sodium carbonate in water does not reach pH 13 at practical
concentrations. Additionally the byproducts of dithionite reduction are acidic (sulfur dioxide) and so sufficient alkalinity must be maintained so
that e.g. dilute NaOH is out of the question.
However heating concentrated NaOH solutions is damaging to glassware. It is noteworthy that pKa3 of trisodium cyanurate is about 13.5:
https://en.wikipedia.org/wiki/Cyanuric_acid#Properties
So a solution of trisodium cyanurate can be used to stabilize the sulfoxylate ion without using large amounts of sodium hydroxide. The pKa of water is
15.7 so generally [C3N3O3(3-)] will be 100 times larger than [OH-] which means that cyanurate solutions should damage glassware 100 times more slowly
than hydroxide solutions.
Somewhat inconveniently, for every two electrons, dithionite releases four protons at pH 13:
S2O4(2-) + 4 H2O >> 2 e- + 2 SO3(2-) + 4 H+
S2O4(2-) + 4 C3N3O3(3-) + 4 H2O >> 2 e- + 2SO3(2-) + 4HC3N3O3(2-)
So the molar quantity of trisodium cyanurate or other base must be four times the molar quantity of dithionite. This is annoying but not unachievable
because cyanurate is very cheap.
As far as I'm aware there are no literature references for reduction with sodium dithionite using trisodium cyanurate as the base, but the basic
science says it should work. Cyanuric acid is OTC as TCCA but I don't know if anyone has really tried to convert this to the trisodium salt.
Attached is the ref for pKa of sulfoxylate.
Attachment: makarov2006.pdf (157kB)
This file has been downloaded 318 times
EDIT: if TCCA is unavailable, cyanuric acid may also be produced by the thermal decomposition of urea.
[Edited on 4-5-2017 by clearly_not_atara]
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Melgar
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If you're worried about glassware corrosion, it may interest you to know that, believe it or not, potassium corrodes glass less than sodium does,
despite being a stronger base.
http://www.springerlink.com/content/n5523871x634l523/
"Sodium hydroxide corrodes silicate glasses of all compositions to a greater extent than all the other hydroxides. It is proposed to call this
phenomenon the sodium anomaly. The sodium anomaly probably arises from peculiarities in the peptization of silicon acids by hydroxyl ions in presence
of sodium ions."
Calcium hydroxide is the best out of all of them though. Its corrosive effect on glass is less than that of distilled water.
I don't think that using cyanurate as a buffer like that is very likely to work. There would just be an equilibrium with -OH ions and cyanurate ions,
as well as silicate ions from the glass. Also, the fact that it's an organic molecule that can participate in the reaction isn't good. What about a
different acid that's very weak and frequently used as a buffer? Like say, boric acid? Combine that with some other hydroxide, and you should be
set.
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clearly_not_atara
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| Quote: | | it may interest you to know that, believe it or not, potassium corrodes glass less than sodium does, despite being a stronger base.
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That's very cool, I didn't know that.
| Quote: | | Calcium hydroxide is the best out of all of them though. Its corrosive effect on glass is less than that of distilled water. |
Come to think of it, this would probably be the best solution, since Ca will also trap any formed SO2 as insoluble CaSO3. That's a big advantage.
EDIT: No, it sounds nice, but calcium dithionite will precipitate and stop the reaction.
| Quote: | | There would just be an equilibrium with -OH ions |
It's true that the equilibrium concentration of [OH-] is somewhat higher than I anticipated because the pKb is 14-pKa and not 15.7-pKa as I
misremembered. However, relative to 3M KOH, a 3M solution of K3C3N3O3 contains only 0.84M [OH-]. And the [OH-] concentration decreases
faster than the [C3N3O3(3-)] concentration as the rxn progresses. After 0.1 moles of protons are released, the [OH-] concentration drops to
0.75M while [C3N3O3(3-)] stays at 2.15.
You can actually make the [OH-] concentration lower at a cost of a slightly lower rxn rate by using some K2HN3C3O3. If you use 1 M K2HN3C3O3 and 2 M
K3N3C3O3 then [OH-] drops to 0.374.
| Quote: | | Also, the fact that it's an organic molecule that can participate in the reaction isn't good. What about a different acid that's very weak and
frequently used as a buffer? Like say, boric acid? |
Cyanurate is very symmetric, which gives it a low nucleophilicity, which is important because the reactant we're concerned about (sulfoxylate)
generally reacts by nucleophilic addition. Hydroxide is very nucleophilic, but will be present at some concentration in all cases. Boric acid by
contrast reacts spontaneously with many oxygen-containing compounds (and if it is used as a buffer there must be free B(OH)3 present).
Tetrahydroxyborate (the only true counterion of boric acid) in any case is not nearly basic enough to stabilize the sulfoxylate anion. Triphosphate is
okay with pKb = 1.7.
Trialkalium orthovanadates are basic enough, but in theory it should be energetically favorable for them to react with dithionite which may
cause both a drop in pH and also an explosion neither of which is desirable. Guanidines and DBU are basic enough but may hydrolyze under the rxn
conditions. Pyrrolidine is the most basic simple amine with pKb = 2.7.
Overall I have not found anything which replicates the desirable properties of cyanurate. Benzhydrol is the only compound I can think of with the
right pKa and it is much more nucleophilic. Phosphate is closest, followed by the anion of oxazolidinone (pKb = 2).
[Edited on 10-5-2017 by clearly_not_atara]
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Melgar
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Ok, cool, I'm glad I could help. However, due to Ca(OH)2's low solubility, I think it's maximum pH is 12.3. However, if less-soluble calcium salts
precipitate out, they'll be replenished with new Ca(OH)2 entering solution, thus stabilizing the pH. Thinking about it, I vaguely remember a
chemistry lecture where the professor explained how Nair (hair-removing cream) works. Apparently, there's a lot of calcium hydroxide in it and a
small amount of sodium hydroxide. The sodium hydroxide raises the pH, and reacts with a divalent anion, which is then pulled out of solution by the
calcium hydroxide, replenishing the sodium hydroxide, and allowing more calcium hydroxide to enter solution. Because pH depends primarily on calcium
salt solubility, the reaction is slow and controlled, and it only eats through your skin at a very slow rate.
So I guess if you have divalent ions that form insoluble calcium salts, you could raise the pH with just a small amount of KOH to react first, then
let the Ca(OH)2 pull it out of solution as soon as it's there. The Ca(OH)2 should also prevent any alkali hydroxides from attacking the glass, and
regenerate the KOH via a solubility-driven metathesis.
edit: I didn't mean that three-proton boric acid should be used, just that a borate salt of whatever base you use should be used. So tripotassium
borate then, with a pKa of 13.3. Seems ideal. Dipotassium borate would be 12.4, if I'm not mistaken.
[Edited on 5/10/17 by Melgar]
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