When Cl is added to water, is there any way to control the ratios of ClHO to HCL produced?
Is there an upper limit for concentration in solution, I've looked on google, but can't find much apart from pool chlorine suppliers etc.BromicAcid - 1-6-2004 at 17:36
1.46 g Cl2 soluble in 100 ml of H2O Really I don't see the point of controlling the equilibrium here.
Cl2(g) + H2O(l) <----> HCl(aq) + HOCl(aq)
HOCl is fairly unstable when concentrated beyond a few % and the reverse reaction will occur readily:
HOCl(aq) + Cl-(aq) + H+(aq) ----> Cl2(g) + H2O(l)
That is the principle behind adding H2SO4 to Ca(OCl)2 to produce chlorine, it is a reliable reaction. So I'm sorry to say even if you could
shift the equilibrium it still wouldn't make much of a difference at only a few precent.
But you do have a point in that it is useful. The oxidation potential of HOCl (1.64V) is much greater then just the hypochlorite anion (.89V).
However to make a solution of HOCl maximally concentrated you could rely on the insolubility of CaSO4 and dissolve Ca(OCl)2 in H2O and add H2SO4 and
lucky for you the precipitates are somewhat more manageable due to their low volume.
Additionally if you did not have access to Ca(OCl)2 you could add Cl2 to a solution of Ca(OH)2 and the CaCl2 will stay in solution whereas the
Ca(OCl)2 will precipitate out.
Bubbling Cl2 thought concentrated hot alkali will oxidise many things. High temperatures favor the formation of chlorate. But you are not favoring
the formation of anything in that NaCl is also produced in equal molar amounts I believe.
If you were to post more information on what your final goal to the equilibrium shift will be maybe I can be of further assistance. BTW the maximum
concentration of NaOCl in solution I believe corresponds to the formula NaOCl*7H2O beyond that it decomposes rapidly but that's from
memory.
So in conclusion:
Quote:
When Cl is added to water, is there any way to control the ratios of ClHO to HCL produced?
Not to any significant extent, and regardless HOCl is fairly unstable, however it is possible to seperate the anions with the aid of calcium salts.
Quote:
Is there an upper limit for concentration in solution
1.46 g Cl2 soluble in 100 ml H2O, but in terms of HCl produced, the azeotrope is 20% and concentrated solutions are 38%. As for HOCl, my Condensed
Chemical Dictionary says it is only stable in very dilute solutions.
Hmmmm.... actually I just researched a bit and HOCl can actually reach concentrations of 25% and decompose only slowly when pure. However it
decomposes very rapidly when even small amounts of acid are present. No wonder I never got concentrated solutions when adding H2SO4 to Ca(OCl)2.
Mind you that even when something like NaOCl is in a basified solution because HOCl is a really weak acid it still exists atleast partially as HOCl
because of the equilibrium.
So by adding the NaOH, you eliminated the HCl immediately by neutralizing it and the HOCl is slightly eliminated and therefore you destroy the
right-hand side of the equation and the Cl2 just keeps dissolving, this is all just coming together in my head right now, is that why Cl2 dissolves so
fast in NaOH(aq)?
Yes, I know this post ranted a lot.
[Edited on 6/2/2004 by BromicAcid]Lestat - 1-6-2004 at 17:46
I was thinking more of using it to make sodium hypochlorite, bleach is just too impure for my purposes
[Edit]
I'm after using it as a start point for chloroform mostly, and NCl3
[Edited on 2-6-2004 by Lestat]
[Edit]
Phosgene is made by oxidation of chloroform, would KMnO4 be sufficient for making phosgene from chloroform, or is a catalyst necessary?
[Edited on 2-6-2004 by Lestat]BromicAcid - 1-6-2004 at 18:00
The manufacture of sodium hypochlorite is going to yield the same impurities present in the industrial manufacture, although you may be able to make a
stronger solution and therefore eliminate your volume of waste. In addition the reaction will yield the same impurities regardless of the grade of
hypochlorite. Therefore the only sure way to guarantee the purity of CHCl3 in my opinon is to distill it with some anhydrous salt. My very old
method was to distill in a volumetric flask on it's side into a florence flask on it's side in an ice bath. However don't forget
Organikum's method involving electrolysis of a solution of acetone with NaCl in water, much better then working with chlorine gas. And
doesn't NCl3 require chlorine gas, which would of course yield less impurities then NaOCl.
As for phosgene I believe iron is one catalyst and it is the reaction between water and chloroform that is the nuisance reaction that causes this in
the presence of light. I don't know about purposely making phosgene from chloroform with an oxidant.
NCl3 Thread
And of course there is a are a couple threads about chloroform. However phosgene has not been as frequently discuessed.
[Edited on 6/2/2004 by BromicAcid]Lestat - 1-6-2004 at 18:07
From what I've read, NCl3 is supposed to be be able to be made by passing ammonia gas through hypochlorite bleach, but when i did it, there was a
lot of foam in the jar of bleach, and just the smallest trace of yellow liquid which failed to go off even when touched with a flame.
As for chloroform, I once mixed nail varnish remover with bleach in an attempt to make it, the bleach was too impure for any decent yield, but i could
definatley smell chloroform.
Maybe using pure calcium hypo would give a much better yield?Proteios - 1-6-2004 at 18:14
Quote:
Originally posted by Lestat
I was thinking more of using it to make sodium hypochlorite, bleach is just too impure for my purposes
[Edited on 2-6-2004 by Lestat]
LOL.... I was waiting for you to come back Lestat...... and wasting no time in posting junk again....... but ill humour you for the moment......
Here are a few retorical questions to help you on your way.
1) do you know what bleach is?
2) do you know how bleach is made?
3) Do you understand the kinetic and equilibria associated with 'bleach'
4) given the production method, what do you think the likely composition of bleach is?
And finally, an most pertenently.....what impurity is it that makes bleach unsuitable for you reaction? (which you still havent given us any details
on)
*EDIT*
Thanks for the P2P message lestat......
as i suspected, you dont know how bleach is made........ look it up!
secondly.....
Quote:
I was perhaps wrong in referring to bleach as too impure, not concentrated enough would be more accurate.
A point that you failed to inform Bromic of before he wasted all that time typing in now... redundant info.
A point that you failed to correct.
If you are stuggling with the difference between concentration and purity you should only be reading this forum, and not posting on it. For the sake
of everyone else... think before you post.
For you sake.... look up the hazard sheets on NCl3 and CHCl3.
[Edited on 2-6-2004 by Proteios]vulture - 2-6-2004 at 06:48
PM from lestat:
Quote:
Greetings. I apologise for my previous conduct, I was just new to boards in general, I never meant to postwhore but I realise i sounded like an
arsehole.
BUT, I don't see any significant change in your attitude.
Consider this as your last warning.unionised - 2-6-2004 at 13:49
The answer to the question, as first asked ie can you change the ratio of HCl to HClO is no.
For each chlorine oxidised to HClO one gets reduced to HCl. There will always be the same number of molecules of each formed (though the HClO can
decompose which will worsen things from your point of view).Proteios - 2-6-2004 at 15:00
from memory... and unchecked......
hydrolysis/dissolution of Cl2O in water.... gets you the nearest thing to HOCl... although the stuff disproportianates with time.BromicAcid - 2-6-2004 at 17:57
Yes, chlorine monoxide is the acid anhydride of hypochloric acid, however I don't believe it is very soluble in water and it undergoes a bit of
self oxidizing if memory serves not only furnishing HOCl on dissolusion in water but also HClO2 and possibly HClO3, more so the last one if heat is
present. But if the solution is kept fairly cool and the addition slow and controlled it should readily furnish the HOCl desired, however chlorine
oxides are notoriously unstable, the most stable being dichlorine heptoxide which explodes from shock, heat, organic matter and other things and this
is the most stable.
[Edited on 6/3/2004 by BromicAcid]
Hypochlorous Acid HClO
kazaa81 - 5-12-2004 at 13:11
Hallo to all,
I'm interesting to know how to SIMPLY make hypochlorous acid and their salts.
I've tried by adding acetic acid (9%) to 5% NaClO but, chlorine like smell has becomed, due to the production of Cl2 and not of HClO.
Any info. is welcomed.
Thanks for helpBromicAcid - 5-12-2004 at 13:28
Hypochlorous Acid almost the same title, this should be there. The chlorine was produced by decomposition of hypochlorous acid. You had
hypochlorous acid, but it's unstable for a number of reasons. Very pure solutions can be concentrated to nearly 20% with a heat lamp though.
They decompose at room temperature but can be made to stay around for a while if kept at lower temperature after they're made.
Thanks
kazaa81 - 5-12-2004 at 13:43
Thanks very much, Bromic,
but I've previously read the theard you've post and this seem to me incomplete.
I'm searching for simple methods for making HClO and using it to make their salts...many thanks indeed.
Or the likewise reaction with barium hypochlorite. Solution must be kept cold.
Normally:
Cl2 + H2O <---> HOCl + HCl
Only a small amount of Cl2 solvated in H2O at any time, but a base will shift the equilibrium, hypochlorites usually less soluble then chlorides and
crystallizing out:
Cl2 + 2MOH ---> MCl(aq) + MOCl(s) + H2O
The hypochlorite obtained by filtration.
Chlorine oxide is the anhydride of HOCl and solvates in water according to:
Cl2O + H2O ---> 2HOCl
But then again chlorine oxide is not noteably stable, which is good otherwise acidified solutoins of HOCl might produce it in quantities that could
prove to be an explosion hazard.
Your best bet is the insitu preparatoin of HOCl and simultaneous reaction with a metal oxide/hydroxide slurry, and filtration/washing the obtained
solid. This is also the same procedure for organic hypochlorites. Many hypochlorites lack stability, e.g., NaOCl decomposes if you take it beyond
NaOCl*6H2O and nearly all of them decompose enough over time to where solutions of them should be restandardized with every use. Remember reaction
mixture should be kept cool <10C otherwise ClO2- ClO3- formation might become overzealous.
Quote:
I've previously read the theard you've post and this seem to me incomplete.
Regardless, your question should have just been posted to that thread, it would not surprise me if these threads end up merged.
[Edited on 12/5/2004 by BromicAcid]neutrino - 5-12-2004 at 16:07
If you make any HClO, make sure that your solution is free of chloride (Cl-) ions. Otherwise, this happens:
Remember that commercial bleach is made by this reaction (in the other direction with NaOH added to neutralize the HClO formed), so it will always
have a lot of chloride in it.
