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Author: Subject: Iodate to chlorate?
tahallium
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[*] posted on 15-4-2020 at 13:00
Iodate to chlorate?


Okay I know this sounds stupid but can chlorine gas be oxidized to chlorite/chlorate? Without using nitrate cuz I know sodium nitrate can oxidize chlorine to sodium chlorite
Can iodate/permanganate/chromate oxidize chlorine?
How can I test for chlorate?
Can chlorine react with oxygen at high temperature?
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clearly_not_atara
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[*] posted on 15-4-2020 at 13:14


What does this have to do with iodate?

But anyway, this is pretty much impossible, and wouldn't be easier even if it worked. The reason is that Cl2 gas has a relatively high entropy which forbids any endothermic reactions from occurring. Instead oxygen is produced, even with very strong oxidizers:

Cl2 + O3 >> Cl2O + O2

Cl2O + O3 >> Cl2 + 2 O2

etc...

Alkaline conditions are necessary for the formation of chlorine oxyanions.




[Edited on 04-20-1969 by clearly_not_atara]
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tahallium
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[*] posted on 15-4-2020 at 15:18


Thanks but can bubbling chlorine gas through hot water produce chloric acid?
https://en.wikipedia.org/wiki/Chloric_acid?wprov=sfla1
Another method is the heating of hypochlorous acid, producing chloric acid and hydrogen chloride:

3 HClO → HClO3 + 2 HCl
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Texium (zts16)
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[*] posted on 15-4-2020 at 15:58


Quote: Originally posted by tahallium  
Thanks but can bubbling chlorine gas through hot water produce chloric acid?
No! Why would you even think that could work?



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Texium (zts16)
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[*] posted on 16-4-2020 at 01:22


You can convert chlorine to chlorate, by bubbling it through a concentrated solution of KOH or K2CO3, until no more chlorine is absorbed. After that, you have to heat the solution to 70 C or so for a while. In that reaction you get chlorate. Keep in mind though that for each single chlorate ion you get 5 chloride ions. After this, you have to boil down the solution and allow it to cool down slowly to 0 C. You will get some crystalline KClO3.

Better is using electrolysis of KCl, using a graphite anode or an MMO anode. The solution also shouild contain a tiny amount of chromate or dichromate, otherwise you hardly get any chlorate in the long run.




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chornedsnorkack
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[*] posted on 17-4-2020 at 11:11


Chlorinating water allows many possible reactions:
  1. Cl2+H2O<->HCl+HClO
  2. 2Cl2+2H2O<->4HCl+O2
  3. 2Cl2+2H2O<->3HCl+HClO2
  4. 5Cl2+4H2O<->8HCl+2ClO2
  5. 3Cl2+3H2O<->5HCl+HClO3
  6. 4Cl2+4H2O<->7HCl+HClO4

Note that:
  1. All of the reactions produce a strong acid HCl - therefore adding a base or a neutral buffer favours right sides of all 6 above chlorine in headspace gas or molecular solution
  2. All of the reactions decrease number of gas molecules - number 2 and 4 produce less oxygen or chlorine dioxide than consume chlorine, and other 4 have no gas on right. Therefore storing chlorinated water should produce vacuum in headspace

So how do you favour reaction 5 over all the others?

[Edited on 17-4-2020 by chornedsnorkack]
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lordcookies24
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[*] posted on 24-4-2020 at 20:00


Quote: Originally posted by Texium (zts16)  
Quote: Originally posted by tahallium  
Thanks but can bubbling chlorine gas through hot water produce chloric acid?
No! Why would you even think that could work?


Perhaps he confused chloric acid (HClO₃) with hypochlorous acid (HClO). I am hoping that is the case because then he would be right.
Cl₂ + H₂O ⇄ HCl + HClO
This would not work in a hot or bright environment though, as HClO can be subject to fairly exothermic decomposition.

If you truly did mean chloric acid, it can be made through a reaction of sulfuric acid with barium chlorate.

Ba(ClO₃)₂ + H₂SO₄ → 2HClO₃ + BaSO₄

Barium sulfate (BaSO₄) is insoluble in water and alcohols but soluble in sulfuric acid, so make sure to get your stoichiometry right and it will be easy to get rid of.
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[*] posted on 24-4-2020 at 23:03


Quote: Originally posted by lordcookies24  
Quote: Originally posted by Texium (zts16)  
Quote: Originally posted by tahallium  
Thanks but can bubbling chlorine gas through hot water produce chloric acid?
No! Why would you even think that could work?


Perhaps he confused chloric acid (HClO₃) with hypochlorous acid (HClO). I am hoping that is the case because then he would be right.
Cl₂ + H₂O ⇄ HCl + HClO
This would not work in a hot or bright environment though, as HClO can be subject to fairly exothermic decomposition.

HClO has 2 common decomposition reactions:
1)2HClO > 2HCl+O2
2)3HClO > 2HCl+HClO3
As you see, the second reaction combined with HClO formation gives the net result
3Cl2+3H2O > 5HCl+HClO3
So, how do you get that reaction, and not O2?
Chlorates are known to be formed in hot water, but does this also apply in acidic pH? Any other conditions that would favour HClO3 over both O2 and HClO?
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clearly_not_atara
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[*] posted on 25-4-2020 at 08:48


Quote: Originally posted by chornedsnorkack  

So, how do you get that reaction, and not O2?
Chlorates are known to be formed in hot water, but does this also apply in acidic pH?

Again, no. As I stated earlier, alkaline conditions are necessary for the formation of chlorine oxyanions. You can't pour out a bucket of water and watch it flow up a hill. Cf. the Pourbaix diagram:

https://www.nrcresearchpress.com/doi/pdfplus/10.1139/v85-155

Chlorine is more stable than any chlorine oxoacid, and it obviously has higher entropy. Therefore, the reaction Cl2 + H2O >> HClOx is not spontaneous under any (reasonable) conditions. So for instance in the above paper, perchlorate solution has an equilibrium pressure of O2 of twenty atmospheres at pH 12.5!

At sufficiently high pressures the reaction 2Cl2 + 7O2 >> 2(Cl2O7) may become spontaneous, but these are not "reasonable conditions" by any stretch of the imagination.

Attachment: karlsson2016.pdf (5.5MB)
This file has been downloaded 31 times





[Edited on 04-20-1969 by clearly_not_atara]
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chornedsnorkack
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[*] posted on 25-4-2020 at 11:23


Quote: Originally posted by clearly_not_atara  
Quote: Originally posted by chornedsnorkack  

So, how do you get that reaction, and not O2?
Chlorates are known to be formed in hot water, but does this also apply in acidic pH?

Again, no. As I stated earlier, alkaline conditions are necessary for the formation of chlorine oxyanions. You can't pour out a bucket of water and watch it flow up a hill. Cf. the Pourbaix diagram:

https://www.nrcresearchpress.com/doi/pdfplus/10.1139/v85-155

Chlorine is more stable than any chlorine oxoacid, and it obviously has higher entropy. Therefore, the reaction Cl2 + H2O >> HClOx is not spontaneous under any (reasonable) conditions. So for instance in the above paper, perchlorate solution has an equilibrium pressure of O2 of twenty atmospheres at pH 12.5!

At sufficiently high pressures the reaction 2Cl2 + 7O2 >> 2(Cl2O7) may become spontaneous, but these are not "reasonable conditions" by any stretch of the imagination.

At which concentration?
Simple cold chlorinated water is described as having 7,3 g/l dissolved chlorine total at 20 Celsius, 1000 mbar total headspace pressure meaning about 50 mbar water vapour and 950 mbar chlorine in headspace.
It is further described as about 70 % molecular dissolved chlorine and 30 % hydrolyzed. Meaning that of the total 100 mmol/l dissolved chlorine, 70 mmol/l is molecular solution and 30 mmol/l does react by exactly
Cl2+H2O<>HCl+HClO
It is not 1000 mmol/l, sure, but it is reasonable conditions to get 30 mmol/l HClO. And 30 mmol/l HCl lowers pH to 1,5 in absence of bases/buffers.
If you also equilibrate the reaction
3Cl2+3H2O<>5HCl+HClO3
what would be the total concentration of hydrolyzed chlorine besides the unaltered 70 mmol/l molecular dissolved chlorine?

[Edited on 25-4-2020 by chornedsnorkack]
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clearly_not_atara
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[*] posted on 25-4-2020 at 13:34


How many words do your people have for zero? The equilibrium species is never chlorate, it is perchlorate, and equilibriation is glacially slow at reasonable temperatures and pressures. You cannot produce chloric acid at low pH. Even if you could, there is no good reason to. Even if there were, the standard methods are safer.

The true equilibrium is 2 Cl2 + (many) H2O >> 4 H+ + 4 Cl- + O2(g). Only at pH 14 does the equilibrium pressure of O2 drop below atmospheric.

"What if...?" No.

Quote:
At which concentration?

Read the paper.

[Edited on 25-4-2020 by clearly_not_atara]




[Edited on 04-20-1969 by clearly_not_atara]
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