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Author: Subject: Using solid Sodium Hypochlorite to generate Chlorine?
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[*] posted on 7-4-2018 at 06:11
Using solid Sodium Hypochlorite to generate Chlorine?


Sodium Hypochlorite reacts with HCl to form Cl2 and Tablesalt.
NaClO + 2 HCl ----> Cl2 + H2O + NaCl

I have seen this reaction used to genrate Chlorine from HCl and Bleach(6% NaClO).
But now i found pool chlorinator an Amazon claiming to be 96% "Aktivchlor" with is a German term usualy used to refer to NaClO.

(https://www.amazon.de/Chlor-Multitabs-200g-mit-Aktivchlor/dp...)

Is 96% a realistic number?
Can Solid pool chlorinator be used in a Chlorine generator or is it advisable to use 10% NaClO solution?
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[*] posted on 7-4-2018 at 06:28


I doubt that it is NaOCl. More likely Ca(OCl)2.
NaOCl.5H2O is (apparently) a greenish solid that melts at 18C. I have never heard of it being sold like that as a commercial product.
Calcium hypochlorite however is commonly sold as a pool-shock oxidiser. And yes, you can produce Cl2 from it. I believe that trichloroisocyanuric acid is considered better, but the difference is probably marginal.




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[*] posted on 7-4-2018 at 06:46


You should be able to use solid hypochlorites with HCl considering it has plenty of water, the reaction proceeds via the hypochlorous acid intermediate, in solution, which is unstable.

OCl- + H+ -> HOCl
HOCl + HCl <-> Cl2 + H2O, where the equilibrium is driven to the right by removal of chlorine as a gas.




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[*] posted on 7-4-2018 at 08:43


Can even do without the HCl, like use NaHSO4. Also, using Damp Rid a source of CaCl2. Try the follow path:

Add CaCl2 to NaOCl and place in a vessel filled with CO2 (from, say, the action of vinegar on baking soda). Shake forming a white precipitate (CaCO3) in a solution of dilute HOCl. Performs this many times, can cool to isolate out the HOCl and store for short periods as hypochlorous acid is not stable especially to light or contaminants. Reactions:

2 NaOCl + CaCl2 = Ca(OCl)2 + 2 NaCl

Ca(OCl)2 + H2CO3 --> CaCO3 (s) + 2 HOCl

To possible create chlorine gas, try a hot drip consisting of dilute cheap H2O2 and NaCl (the drip solution is preheated with added NaCl to help salt out the chlorine gas that will be mixed with evolving oxygen gas).

HOCl + H2O2 --> HCl + H2O + O2 (g)

HCl + HOCl = Cl2 + H2O

So, HCl appears to have been replaced with CO2, CaCl2, H2O2 and NaCl.
---------------------------------------------

Another path not using H2O2, add preheated aluminum foil to a vessel containing a solid piece of copper metal, the HOCl from above and a some amount of NaCl serving as an electroyte. Jump start the reaction in a microwave oven for 30 seconds. Repeat till reaction appears to be vigorous. Add more NaCl to help salt out the Cl2.

The chemistry follows the so called bleach battery creating here Al(OH)3 as a white precipitate and chlorine (see https://www.sciencemadness.org/whisper/viewthread.php?tid=30... ).

An alternate path to HOCl for the bleach battery for chlorine generation, which is less expensive than the CO2/CaCl2 plus NaOCl path is by the action of vinegar/NaCl plus NaOCl. However, in the battery chamber, do not allow prolong (days) exposure of acetate to residual chlorine as a nasty organochlorine compound may form, per my experience.

[Edited on 8-4-2018 by AJKOER]
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[*] posted on 8-4-2018 at 23:06


Solid NaOCl is pretty much unobtainium. It TECHNICALLY exists (and even that is a crystallohydrate, not an anhydrous salt), but it's too hard to isolate and store and has little or no benefits over the solution to be useful.



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[*] posted on 9-4-2018 at 00:03


Thank for that, Ave. That's pretty much what I thought but I could not recall a redference.



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[*] posted on 9-4-2018 at 00:21


Quote: Originally posted by ave369  
Solid NaOCl is pretty much unobtainium. It TECHNICALLY exists (and even that is a crystallohydrate, not an anhydrous salt), but it's too hard to isolate and store and has little or no benefits over the solution to be useful.


Nonsense. It is commercially available (via chemical supply) and appears to be not too difficult to make if one was so inclined. It has massive benefits over the solution in terms of stability on storage though.

https://pubs.acs.org/doi/abs/10.1021/acs.oprd.7b00288

[Edited on 9-4-2018 by DJF90]
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[*] posted on 9-4-2018 at 00:51


Sure about the sodium salt? I know about the 5-hydrate, but according to my literature that is quite unstable and has no practical importance. This contradicts the link you provide. I never found a commercial supplier for the solid salt and if it really is available I would love to have some.



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[*] posted on 9-4-2018 at 01:21


@woelen: Yes, the pentahydrate was shown to retain 99% of its active chlorine over 1 year, IIRC. This directly contrasts the stability of the solution. TCI are the only commercial supplier I am aware of as yet, but it is a recently developed product so it may take a while to filter through to the mainstream.

http://www.tcichemicals.com/eshop/en/gb/commodity/S0939/

The OPRD paper has a experimental section detailing its preparation though, and it looks simple enough.

[Edited on 9-4-2018 by DJF90]
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[*] posted on 9-4-2018 at 01:29


Here is a PDF of the OPRD paper mentioned above by DJF90.



Attachment: acs.oprd.7b00288.pdf (2.5MB)
This file has been downloaded 484 times





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[*] posted on 9-4-2018 at 04:37


Quote: Originally posted by DJF90  
and appears to be not too difficult to make if one was so inclined.
[Edited on 9-4-2018 by DJF90]


How exactly do you do that? You can't crystallize it out of the solution by boiling, it will decompose. You can't precipitate it by common ion effect, the Na salt is too soluble for that. Please provide proof of that statement.




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[*] posted on 9-4-2018 at 05:26


If you cared to read the paper (it is open access, and also attached as .pdf by Hexavalent above, so there is no excuse for being ignorant/lazy), there is an experimental section that provides details:
Quote:
Chlorine gas is added to a 45−48% NaOH solution to prepare a highly concentrated NaOCl solution. After removing of the precipitated NaCl by filtration, the filtrate is cooled to around 12 °C to precipitate the NaOCl·5H2O crystals, which are collected by centrifugal filtration.

Note this was developed for industrial production, hence the centrifugal filtration. I'm sure good ol' vacuum filtration will work fine for a lab scale prep.
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[*] posted on 9-4-2018 at 06:23


Was it ever done by anyone in the lab?



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[*] posted on 9-4-2018 at 06:36


Interesting info. This certainly may be something I will try on a small scale. Leading chlorine through 50% NaOH until no more chlorine is absorbed should not be too difficult.

What I find remarkable is that if it is so simple as described, that this salt is not available already for 50 years or so. The common hypochlorite Ca(OCl)2.2H2O is available already for tens of years at low cost, why not the solid NaClO.5H2O? Why does the entire world still use the crappy unstable low concentration bleach if such a great solid and highly concentrated compound can be made easily at low cost? Having solid NaClO would be much more pleasant than Ca(ClO)2. The latter always is a nuisance with all those calcium ions, producing insoluble stuff and turbid liquids.




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[*] posted on 9-4-2018 at 06:37


The original posters link is to 5kg of 200g tablets of Activchlor. However, these look more like the standard 200g tablets of chlorinator that you get in the UK and they are designed for slow-release dispensers. For this purpose they need to be sparingly soluble so I suspect that they are TCCA and not the very soluble sodium hypochlorite pentahydrate. "Activchlor" looks like a trade name and the owner of such a trade name an apply it to pretty much anything they want so just because they sell sodium hypochlorite solution under he same name means very little.
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[*] posted on 9-4-2018 at 10:15


I had a look into the solid NaClO.5H2O, you can purchase from TCI.

The stuff is expensive ($55 per 500 grams) and it must be stored at low temperatures (< 0 C). When stored in a freezer, then you lose hypochlorite at a rate of appr. 1% per year, which indeed is quite well. But you have to store the material in a freezer. Not everyone wants that or can assure storage under such conditions all the time. I myself don't have cold storage, so this chemical is not something I can keep around.

At 27 C it "melts" in its own water of crystallization and it also quickly decomposes when this happens. It forms a concentrated solution. Dissolved NaClO only is stable up to 13% or so at room temperature.




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[*] posted on 9-4-2018 at 19:00


Aktivchlor is usually TCCA. If you want to create chlorine gas you should look for "Langzeit Chlortabletten". The product you linked is TCCA. Source: https://www.poolladen.de/poolchemie/desinfektion/festchlor-organisch/chlortabletten/chlor-l-tabletten-10-0-kg-200g-tablette
Its in german but since you linked a german product i assume you can und erstand it.

[Edited on 10-4-2018 by Cardie]
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[*] posted on 10-4-2018 at 05:16


Here is an extract from an old chemistry journal courtesy of Atomistry.com (see http://sodium.atomistry.com/sodium_hypochlorite.html ):

"Solutions containing a high percentage of sodium hydroxide yield a concentrated solution of the hypochlorite, provided additional alkali is introduced periodically to maintain the concentration of the solution. Under these conditions, the sodium chloride produced by the reaction is precipitated, and a solution of hypochlorite obtained which reacts with hydrochloric acid to yield up to 49.2 grams of chlorine per 100 c.c. of solution. The temperature of the reaction should be maintained below 27° C., but the very concentrated solutions obtained lack stability, the hypochlorite changing to a mixture of chloride and chlorate, one molecule being oxidized at the expense of another. From very concentrated solution sodium hypochlorite separates as a solid hydrate, its composition approximating more nearly to that of the heptahydrate than the hexahydrate. It is a very unstable, hygroscopic substance, but dehydration in a current of dry air at reduced pressure converts it into the solid anhydrous hypochlorite, which melts about 45° C. It is less hygroscopic than the hydrate, and contains 40 to 60 per cent, of available chlorine.

When the turbid liquid formed by heating the heptahydrate at 20° C. is cooled slowly to the ordinary temperature, large greenish-yellow, very deliquescent crystals of the pentahydrate are formed. They melt at 27° C., and are stable at ordinary temperature in absence of air. "

Reading the above carefully does not suggest a very easily prep to me for sodium hypochlorite pentahydrate (with needed periodic additions of NaOH, temperature no more than 27° C, possible disproportionation to chlorate, heating any collected heptahydrate at 20° C and then cooling slowly to RT) which possibly explains the cost issue.

Interestingly also, dehydration of the heptahydrate in a current of dry air at reduced pressure can converts it into the solid anhydrous hypochlorite. The latter per a source (see http://www.oxy.com/OurBusinesses/Chemicals/ResponsibleCare/D... ), to quote:

"Anhydrous sodium hypochlorite is very explosive."

Another source (see "Bretherick, Vol 1 at https://www.monash.edu/__data/assets/pdf_file/0007/568150/br... ) to quote:

"The anhydrous solid obtained by desiccation of the pentahydrate will decompose violently on heating or friction [1,2]."

which adds a friction hazard also.
-------------------------------------------------------------------

Found a more recent reference on NaOCl.5H2O, see https://pubs.acs.org/doi/full/10.1021/acs.oprd.7b00288
--------------------------------------------------------------------
Some may be interested in other safe solid hypochlorites, which are more easy to prepare and have been employed in commercial products!

See, for example, one of my prior threads on "Preparation/Exploration of Mg(ClO)2.2Mg(OH)2, a Solid Hypochlorite", at
https://www.sciencemadness.org/whisper/viewthread.php?tid=26... .

[Edited on 10-4-2018 by AJKOER]
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