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Author: Subject: Why is sodium carbonate not used as starting point for hypochlorite/chlorate production?
Merryp
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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.



far as i know everything is trying to be an equilibrium.
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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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chornedsnorkack
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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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