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Author: Subject: Electrolysis of aqueous NaCl for producing HCl acid
sakshaug007
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[*] posted on 15-3-2009 at 22:49
Electrolysis of aqueous NaCl for producing HCl acid


Hello Everyone,

I'm new to the forum and I am wanting to attempt at home synthesis of hydrochloric acid. My technique of choice is to electrolyze a saturated solution of NaCl to produce H2 and Cl2 gas I was then going to combust the gases to form HCl which would be bubbled through distilled water to form the acid. My question, is there a safer way of forming the HCl gas rather than combustion, i.e., using a solid state catalyst? My HSC program indicates the reaction between H2 and Cl2 is thermodynamically favorable (positive entropy, negative enthalpy and gibbs free energy) at room temperature, does that mean they will react spontaneously in the presence of one another? Any advice is appreciated. Thank you.
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Saerynide
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[*] posted on 15-3-2009 at 23:55


They react in the presence of UV light too, but it will also explode violently...

And remember, spontaneously doesn't say anything about the rate. If I recall correctly, the rate at room temp is very very very slow.

[Edited on 3/16/2009 by Saerynide]




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hissingnoise
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[*] posted on 16-3-2009 at 06:02


Cl2 and H2 react smoothly enough in normal lighting.
But why not just buy HCl, sakshaug007, or use the old H2SO4/NaCl route. . .
Electrolysis is the harder way!
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sakshaug007
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[*] posted on 16-3-2009 at 07:45


Quote:
Originally posted by Saerynide
They react in the presence of UV light too, but it will also explode violently...

And remember, spontaneously doesn't say anything about the rate. If I recall correctly, the rate at room temp is very very very slow.

[Edited on 3/16/2009 by Saerynide]



I did consider using UV light as a catalyst in fact my idea was to seal off the electrolysis cell from any source of light and only have the bubbling container be exposed to the UV, that way it would react in situ with the water and would not be as violent since it would only be small bubbles at any given instant. Thanks for the thermodynamics rate correction, I knew there was something wrong with my theory.
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sakshaug007
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[*] posted on 16-3-2009 at 07:50


Quote:
Originally posted by hissingnoise
Cl2 and H2 react smoothly enough in normal lighting.
But why not just buy HCl, sakshaug007, or use the old H2SO4/NaCl route. . .
Electrolysis is the harder way!



I was interested in devising a setup that would be a relatively easy production of HCl using at home items. I actually don't have any H2SO4 and I wasn't sure on the purity of OTC muriatic acid (HCl). Thanks for your reply.
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hissingnoise
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[*] posted on 16-3-2009 at 10:11


If you can get H2SO4 you'll make the task easier---even draincleaner should work. . .
HCl gas should be led into water by an inverted funnel to prevent suck-back caused by its great solubility.
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Tacho
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[*] posted on 16-3-2009 at 14:16


Please, take a look:

http://www.sciencemadness.org/talk/viewthread.php?tid=2154&a...

Good luck.
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kclo4
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[*] posted on 16-3-2009 at 14:41


Most all stores that have any swimming pool supply will be selling Sodium Bisulfate, which acts like sulfuric acid in these types of reactions. Adding NaCl to Sodium Bisulfate will result in sodium sulfate and Hydrogen chloride of if in solution, Hydrochloric acid.



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hissingnoise
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[*] posted on 16-3-2009 at 15:11


The downside to the bisulphate route is the head required---not as good as H2SO4, but better and more efficient than electrolysis.
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sakshaug007
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[*] posted on 16-3-2009 at 15:24


Alright thanks for all your help, I will no longer attempt HCl synthesis via electrolysis. By the way what are the temperatures required for the NaHSO4/NaCl reaction for HCl? Is it self sustaining?

Thanks
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hissingnoise
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[*] posted on 17-3-2009 at 06:47


I haven't tried it but I assume dry distillation with intimately mixed reactants is the way to go.
The bisulphate proton needs a fair bit of a push!
I can't recall temperatures but I don't think they're attained in solution. . .
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