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Author: Subject: Why is sodium carbonate not used as starting point for hypochlorite/chlorate production?
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[*] posted on 21-1-2020 at 14:47
Why is sodium carbonate not used as starting point for hypochlorite/chlorate production?


Someone suggested this equation:

3 Na2CO3 + 3 Cl2 → 5 NaCl + NaClO3 + 3 CO2

Looks reasonable on paper, and would give the same yield as NaOH.
Is there any particular reason why this reaction would be disfavoured?
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Tsjerk
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[*] posted on 21-1-2020 at 15:32


Try to write down all reaction steps and see where the different equilibriums are and what the K values are. I have no clue but I guess at least one is pretty far from your proposed product.
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[*] posted on 21-1-2020 at 15:38


I think the process will be less efficient. Consider that you are bubbling Cl2 gas through a solution. At the reaction interface you will also have CO2 diffusing back into the bubble. This is going to dilute the chlorine in the bubbles and the amount of unreacted chlorine will be greater.

It will also be more difficult to monitor the reaction. With NaOH you can observe visually how the bubble size diminishes and get an idea of how much chlorine is reacting. If you have CO2 being produced you lose some of that visual feedback.
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[*] posted on 22-1-2020 at 08:09


you can't. in fact adding soda water to bleach will make toxic Cl2 gas! more so with Ca(OCl)2 based ones because of insoluble CaCO3.



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clearly_not_atara
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[*] posted on 22-1-2020 at 08:32


Chlorine only dissolves well in water when hypochlorite is deprotonated, pKa 7.53, so you need pH > 9. A buffered sol'n of sodium sesquicarbonate has a pH of 10.5 but a buffered sol'n of bicarbonate has a pKa of about 8. The latter is not basic enough to dissolve Cl2, so the conjugate acid for this reaction is sesquicarbonate.

So your equation should really be:

4 Na2CO3 + Cl2 + H2O >> NaCl + NaOCl + 2 Na3H(CO3)2

Here it is evident that you need eight moles of sodium to get one mole of hypochlorite! What a waste!




[Edited on 04-20-1969 by clearly_not_atara]
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[*] posted on 22-1-2020 at 11:58


Wait, what? Where does the Cl2 come from? This would assume you had an excess of Cl2 from another process, or that NaOH is far more valuable than Na2CO3?




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[*] posted on 22-1-2020 at 23:39


Doesn't matter, the question is about making bleach from Cl2 and an alkaline solution. the Cl2 indeed comes from an external process/source.

The last answer of clearly_not_atara makes sense. Carbonate ion is just a weak base and a little added acidity (indirectly from the disproportionation reaction of Cl2) quickly neutralizes it.

[Edited on 23-1-20 by woelen]




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[*] posted on 26-1-2020 at 13:32


Quote: Originally posted by clearly_not_atara  
Chlorine only dissolves well in water when hypochlorite is deprotonated, pKa 7.53, so you need pH > 9. A buffered sol'n of sodium sesquicarbonate has a pH of 10.5 but a buffered sol'n of bicarbonate has a pKa of about 8. The latter is not basic enough to dissolve Cl2, so the conjugate acid for this reaction is sesquicarbonate.

So your equation should really be:

4 Na2CO3 + Cl2 + H2O >> NaCl + NaOCl + 2 Na3H(CO3)2

Here it is evident that you need eight moles of sodium to get one mole of hypochlorite! What a waste!


Yes, because HClO is a weak acid.
But HClO3 is a strong acid, like HCl.

At which pH does the reaction
3Cl2+3H2O<->5HCl+HClO3
have its equilibrium on the left?
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clearly_not_atara
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[*] posted on 26-1-2020 at 19:11


Chloric acid is not stable with respect to decomposition:

2 HClO3 >> H2O + Cl2 + 2 O2 + ~275 kJ/mol

It is itself produced only by the disproportionation of hypochlorite. It cannot be produced at any pH where hypochlorite does not form.




[Edited on 04-20-1969 by clearly_not_atara]
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[*] posted on 24-2-2020 at 13:57


Per a source (see https://books.google.com/books?id=dXn3aB1DKk4C&pg=PA454&...), Cl2 first reacts to neutralize the alkaline aqueous Na2CO3 per the cited reaction:

Cl2 + H2O + 2 Na2CO3 -> NaOCl + NaCl+ 2 NaHCO3

And then, neutralizes the more alkaline sodium hypochlorite:

Cl2 + H2O + NaOCl -> 2 HOCl + NaCl

where the two reactions above are stated by the reference. However, I would further argue even without more added chlorine, the bicarbonate can act on hypochlorite also liberating hypochlorous acid in an equilibrium reaction (and, with an insoluble carbonate, like CaCO3, the reaction moves completely to the right):

HCO3- + OCl- <--> HOCl + CO3(2-)

So on net, the action of chlorine on Na2CO3 is NOT a good path to NaOCl, but more powerful, at least in terms of disinfecting, or especially, bleaching in the presence of ferrous ions (as are found in hot tap water) with the formation hypochlorous acid, and associated radicals created therefrom. With respect to radicals, a Fenton-type reaction based on HOCl (in place of H2O2 with the classic Fenton reaction) proceeds as follows leading to active radicals:

Fe(II) + HOCl -> Fe(lll) + •ClOH-

where at pH > 5:

•ClOH- -> Cl- + •OH (See cited sources at https://www.sciencemadness.org/whisper/viewthread.php?tid=15... )

where the hydroxyl radical can breakdown organics (like stains) or the DNA of bacteria,...

Also, in the presence of chloride:

•ClOH- + Cl- -> •Cl2- + OH-

Also importantly, created HOCl here is not stable so one cannot place it on a shelf at the supermarket (so, the action of chlorine on sodium carbonate does NOT create even a stable bleach alternative).

[Edited on 24-2-2020 by AJKOER]
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