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Author: Subject: H2O + Cl2 ==> HCl + HClO?
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[*] posted on 7-5-2010 at 23:36
H2O + Cl2 ==> HCl + HClO?


I was looking for the ways of making hydrochloric acid at home and I found this:

http://www.ucc.ie/academic/chem/dolchem/html/elem/elem017.html

Quote:

Reaction of Chlorine with Water

When Chlorine Water (i.e. a solution of chlorine gas in water) in a flask inverted in a basin of the same liquid is exposed to bright sunlight, the Chlorine is decomposed and a solution of Hydrochloric Acid remains.
H2O + Cl2 ==> HCl + HClO
The Hydrochloric Acid, HClO, is not very stable and the solution readily decomcomposes, especially when exposed to sunlight, yielding Oxygen.
2 HClO ==> 2 HCl + O2


Is this really true? If it is, I could do the electrolysis of salt water in the bottle and then expose the bottle to the sunlight to make HCl and HClO mixture, after that HClO would decompose to HCl and O2 so that could be one of the easier ways to get the HCl at home.
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[*] posted on 7-5-2010 at 23:55


I guess ill speak even though a quick UTSE could have worked.

Feed CL2 into H2O and you get CL2 + H2O <=> HCl + HOCl

Check wiki,,,,,, the worlds, most used, most hated chemistry site........(Sheeeshhh Hipocracy comes to mind)......

Equilibrium see! You want the right half of the equation with half of the product reduced to HCl in a pefect world.

If you need HCl you must find a means to destroy the hypochlorus acid(HOCl). Heat may do so but im unsure. Then just distill your Azeotropic HCl around like 27% right folks? My memory sucks so look into that azeotrope concentration and all the better ways expressed thru the forum to concentrate it higher.





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[*] posted on 8-5-2010 at 00:36


You have to lead the chlorine away from the salt solution you are electrolysing, because that will be forming NaOH in it; letting the Cl react with that solution gives NaCl and NaOCl, eventually bringing you back to your starting point.

It will take at least 26.8 amp-hours to make 1 mole of 'Cl', a half mole of Cl2 (about 11 liters at STP). In theory that would give you 1 mole of HCl after all the HOCl has decomposed, and 11 liters of O2. This would result in 1 liter of 1 M hydrochloric acid, or about 125 ml of 20% constant boiling HCl, or 85 ml of concentrated hydrochloric acid equivalent to the technical concentrated hydrochloric acid found in hardware stores.

And remember that chlorine will attack and dissolves most metals you'd use as electrodes.

It's possible, but maybe not practical.
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[*] posted on 8-5-2010 at 01:45


Well, according to the quoted text, HOCl should decompose into O2 and HCl (it is unstable). So it looks I should just wait for it to decompose.

I was thinking about electrolysis of it in the bottle, so some chlorine would still get out of the bottle with hydrogen. I would connect the pipe to the bottle and the chlorine and hydrogen would be getting through the pipe. I would then bubble them into water to make chlorine water (hydrogen would escape from the water at the same time. After that I would get chlorine water in the bottle which exposed to the sunlight would form HCl.

For electrodes I would use graphite from pencils, so no reactions would happen.
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[*] posted on 8-5-2010 at 06:04


Probably not a practical way to make HCl. Heating table salt and swimming pool pHdown (NaHSO4) will produce HCl gas, which is easily abosrbed in H2O to make aqueous HCl. Pardon me for sidetracking this thread by suggesting something that will actually work.
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[*] posted on 8-5-2010 at 06:11


In your plan the H will tend to carry away some Cl2. Look up the solubility of Cl2 in water, and remember that to get 1 mole of HCl gas you are making 11+ liters of Cl2 and 11+ liters of H2.

Here's a chart to help http://www.engineeringtoolbox.com/gases-solubility-water-d_1...

At 10 C you'll need around 3,5 liters of water, and it will be close to being completely saturated with Cl2 meaning the later portions will dissolve with difficuly and that the H2 will carry much of that portion off. Of course if the current you use is low, the Cl2 will be reacting with the water fast enough the its concentration will not be as great as if yoy just took it from a ank of Cl2.

Do consider that H2 + Cl2 is a ... rather reactive mixture, and that light can initiate that reaction.


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[*] posted on 6-12-2011 at 19:14


With respect to the decomposition versus the disproportionation of HClO solution into HClO3, I believe some may find very interesting what is reported in the literature with respect to very dilute solutions and the power of diffused sunlight and my actual observations last winter.

From "A treatise on chemistry", Volume 1 By Henry Enfield Roscoe, Carl Schorlemmer, page 192:

"Saturated chlorine water gives off chlorine freely on exposure to the air, and bleaches organic colouring matters. When exposed to direct sunlight it is, if sufficiently dilute, gradually converted into hydrochloric acid with evolution of oxygen:

Cl2 + 2H20 = 4HCl + 02.

It has been proposed to employ this reaction in measuring the chemical action of light, but the decomposition is not sufficiently regular for this purpose ; thus Pedler (1) has shown that a solution containing 1 molecule of chlorine to 64 of water undergoes no appreciable alteration during two months' exposure to tropical sunlight, whilst more dilute solutions undergo more or less decomposition, as shown in the following table

Mols. H20 for 1 mol. Cl2 / Percentage of Cl2 acting on water.

64 no action
88 29%
130 46%
140 29%
412 78%

In the case of more dilute solutions, the reaction in sunlight appears to take place almost completely in accordance with the above equation, except in so far as small quantities of chloric acid are formed. In diffused daylight, however, a considerable quantity of the latter acid is obtained, so that in this case the reactions are probably those put forward by Popper (2):

Cl2 + H20 = HCl + HClO

8HCl0 = 2HCl03 + 6HCl + 02

Under certain conditions, however, sunlight brings about the reverse change, causing the formation of free chlorine from a mixture of hydrogen chloride and oxygen (see p. 200).

1 Journ. Chem. Soc. 1890, 57, 613. 1 Annalen, 1885, 227, 161 "

Now, last winter I left out in the sun (partially open to the air) some fresh dilute HClO in a thick transparent glass flower vase (diffused light?), which I further re-diluted (actually intending to discard hence the diluting to save the pumbling!). After two weeks, I noticed that the solution developed a much stronger chlorine like smell (so much for my effort to dilute!). Passing NH3 near the top produced a cloud of NH4Cl. Upon discarding the solution down the drain in an old shower, I noticed that where I splashed some on the shower floor, an intense bleaching or other chemical action occurred (HClO3?). I now suspect that the diffused sunlight on the dilute HClO produced HCl and HClO3, the latter acid being so strong as to account for the smell and bleaching/chemical reaction. The cold temperature and dilution may have assisted in keeping gases dissolved in the solution and preserving the HClO to be acted upon by the sunlight.

More precisely with respect to the strong smell, HClO3 in the presence of HCl can liberate Chlorine dioxide:

6 ClO2 + 3 H2O <==> 5 HClO3 + HCl

or Chlorine and Chlorine dioxide by the equilibrium reaction:

8 HOCl <==> 4 H2O + 2 ClO2 + 3 Cl2

It is also possible, that no HClO3 was created and the decomposition of some HClO into HCl upon exposure to direct sunlight reacted with more HClO to form Cl2.

For those who have searched the literature on the web, this bipolar observation (referring to very dilute HClO under diffused light vs. dilute and direct sunlight impact on HClO) may explain why some authors note the creation with sunlight of HClO3 and others cite just the decomposition reaction.


[Edited on 7-12-2011 by AJKOER]
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[*] posted on 7-12-2011 at 16:59


I agree with everything posted, but the way you wrote the reaction is very misleading,

Quote: Originally posted by AJKOER  
More precisely with respect to the strong smell, HClO3 in the presence of HCl can liberate Chlorine dioxide:

6 ClO2 + 3 H2O <==> 5 HClO3 + HCl



Better, I think, to write

5 HClO3 + HCl --> 6 ClO2 + 3 H2O

I am not saying that equilibrium does not exist, but it is essentialy a very "one-way" reaction, assuming there are no hydroxide ions to shift the reaction in the opposite direction.


Quote: Originally posted by AJKOER  
It is also possible, that no HClO3 was created and the decomposition of some HClO into HCl upon exposure to direct sunlight reacted with more HClO to form Cl2.


No doubt.

(2)HClO + light --> (2)HCl + O2
HCl + HOCl <==> H2O + Cl2

Quote: Originally posted by AJKOER  

For those who have searched the literature on the web, this bipolar observation (referring to very dilute HClO under diffused light vs. dilute and direct sunlight impact on HClO) may explain why some authors note the creation with sunlight of HClO3 and others cite just the decomposition reaction.


Very much agreed. I think HOCl, even highly dilute in the dark, will decompose in storage

The net equation could probably be described
(5)HOCl + light --> HClO3 + (2)H2O + (2)Cl2
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[*] posted on 13-12-2011 at 06:45


Quote: Originally posted by AndersHoveland  
I agree with everything posted, but the way you wrote the reaction is very misleading,

Quote: Originally posted by AJKOER  
More precisely with respect to the strong smell, HClO3 in the presence of HCl can liberate Chlorine dioxide:

6 ClO2 + 3 H2O <==> 5 HClO3 + HCl




Better, I think, to write

5 HClO3 + HCl --> 6 ClO2 + 3 H2O

I am not saying that equilibrium does not exist, but it is essentialy a very "one-way" reaction, assuming there are no hydroxide ions to shift the reaction in the opposite direction.


AndersHoveland, I agree with you and that's one reason in another recent tread, I posted the catalyst (in the current context Chlorine) that can accelerate the reaction as I and Mellor (see page 288) have presented it.

Reference: "A comprehensive treatise on inorganic and theoretical chemistry", Volume 2 By Joseph William Mellor

"In the presence of chlorine, the reaction progresses: Cl02 + 1/2Cl2 + H20=HCl03 + HCl, with the side reactions: 6Cl02 + 3H20 = 5HCl03 + HCl, and 3Cl2+3H20=HCl03+5HCl. At 60° another reaction : Cl02=1/2Cl2+02, sets in. Consequently, the decomposition of aq. soln. of chlorine dioxide is very complex, for there are (i) 2Cl02=Cl2+202, which is accelerated by raising the temp, or exposure to sunlight; (ii) 6Cl02 + 3H20=5HCl03+HCl, which is accelerated by the presence of chlorides or by platinum; (iii) 2Cl02+1/2C12+2H20=2HCl03 +2HCl, which is accelerated by chlorine; (iv) 3Cl2 + 3H20 = HCl03 + 5HCl, which is accelerated by platinum or chlorine dioxide; and (v) 2Cl2+2H20=4HCl+02, which is accelerated by light."

LINK:
http://books.google.com/books?pg=PA275&lpg=PA275&dq=...


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[*] posted on 13-12-2011 at 08:24


Quote: Originally posted by AJKOER  

"In the presence of chlorine, the reaction progresses: Cl02 + 1/2Cl2 + H20=HCl03 + HCl, with the side reactions: 6Cl02 + 3H20 = 5HCl03 + HCl, and 3Cl2+3H20=HCl03+5HCl.
or exposure to sunlight; (ii) 6Cl02 + 3H20=5HCl03+HCl,

Reference: "A comprehensive treatise on inorganic and theoretical chemistry", Volume 2 By Joseph William Mellor


I am not disagreeing with the researcher's findings, but am suggesting that his interrpretation may not be entirely correct. As HCl actually reacts with chlorate, I do not see how any reaction would favor the formation of those two products together. I think there may be other equations that more accurately reflect the reaction. That source is very old, at a time before modern methods of analysis, and it is obvious that there is plenty of speculation in that book. That was just how chemistry was done then.


Quote:

At 60° another reaction : Cl02=1/2Cl2+02, sets in. Consequently, the decomposition of aq. soln. of chlorine dioxide is very complex, for there are (i) 2Cl02=Cl2+202, which is accelerated by raising the temp,


This certainly is an interesting insight. 60°C is also the temperature that hydrogen peroxide is able to much more rapidly oxidize NH4OH or hydrochloric acid. Could there be a connection? I find it unusual because this is a relatively low temperature for activation energy compared to almost all other reactions.

Quote:

A. D. White also found that aluminum is slowly attacked by hypochlorous acid, and the resulting aluminum hypochlorite immediately decomposes into aluminum hydroxide, oxygen, and chlorine.


The reaction is probably more complex than this. I would think that if Al(OCl)3 was actually formed, it would decompose into Cl2 and ClO2, not necessarily O2.

[Edited on 13-12-2011 by AndersHoveland]
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[*] posted on 13-12-2011 at 12:49


Having dissolved Al in HClO, the unstable Al(OCl)3 appears to quickly decompose with exceptional bleaching power (O2), depositing an insoluble white salt ( Al(OH)3 which may dissolve as the solution becomes more acidic), and clearly there is Cl2 released (odor) and, upon closing the reaction chamber, some chlorine water is formed.

With time and diffused light, the HCl created from the Chlorine water could slowly react with Al(OH)3 moving the Chlorine water equilibrium reaction to the right:

Cl2 + H2O <--> HClO + HCl

3 HCl + Al(OH)3 <--> AlCl3 + 3 H2O

and

3 HClO ---Diffused Light---> 2 HCl + HClO3

and per previously cited reactions, ClO2 could form but, not I suspect, as an immediate product on the decomposition of the Aluminum hypochlorite here.

Interestingly, the only references I could find on Al(OCl)3 are old. The lack of complete understanding of the reaction chain may explain the absence. Old textbooks I find are very good on reporting actual observations, and I agree, are more speculative/inaccurate on precise chemical details.

Here is another dated description of the reaction:

"The aluminum bleach is the most important of these, and its method of preparation is typical of the method used in the preparation of all the others.
The aluminum bleach consists of a solution of a mixture of aluminum chloride and hypochlorite that is made by treating a solution of calcium bleach with aluminum sulphate; the calcium sulphate separates out and the aluminum compounds are left in solution. The aluminum hypochlorite is very unstable and is only made as needed. It is so very unstable that it decomposes on the fiber without the use of acid, and the aluminum compound left is antiseptic, so that it not only does not need to be washed out, but in many cases it is a decided advantage to leave it on the bleached material. For example, when used to bleach paper stock the aluminum chloride prevents fermentation when the stock is stored."

http://books.google.com/books?pg=RA3-PA36&lpg=RA3-PA36&a...


[Edited on 13-12-2011 by AJKOER]
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[*] posted on 14-12-2011 at 16:35


I find it interesting that H2O2 will only reduce chlorate if a >40% concentrated sulfuric acid is added, but that hydrochloric acid will reduce chlorate at any concentration.

[Edited on 15-12-2011 by AndersHoveland]
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