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Author: Subject: Chlorine
BromicAcid
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[*] posted on 25-7-2004 at 05:11


Quote:

Also, if anyone knows the reaction between sodium bisulfate and calcium hypochlorite, what is it?


An educated guess:

2NaHSO4 + Ca(OCl)2 ----> CaSO4 + Na2SO4 + H2O + Cl2 + 1/2O2

Cl2 and 1/2O2, sounds like chlorine monoxide might be decomposing in there somewhere.....




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guy
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[*] posted on 25-7-2004 at 14:00


Quote:

Using a glass jar with a screw on lid, I filled it up with water, then added an equimolar (approximately) quantity of each reactant, quickly screwing the lid on. Chlorine gas bubbled vigorously, and the solution became quite warm. The vapor above the solution became quite green.


The green liquid could be liquified chlorine. Chlorine is easily liquified under pressure.
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[*] posted on 25-7-2004 at 14:14


Have you quoted the bit you thought you had?

[Edited on 25-7-2004 by unionised]
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[*] posted on 25-7-2004 at 22:09


Must have missed it


Quote:

The solution itself seperated into two layers. A gunky, white precipitate settled to the bottom, leaving a pale green (and CLEAR) supernatent.


precipitate -> CaSO4

green --> Cl2 (liquid)

[Edited on 26-7-2004 by guy]
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Saerynide
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[*] posted on 25-7-2004 at 23:24


Liquid chlorine?? Are you sure? :o Thats really freaky... Imagine that bottle broke or leaked :S



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[*] posted on 25-7-2004 at 23:32


I dont know for sure but this site says that chlorine is easily liquified under pressure

http://www.ucc.ie/ucc/depts/chem/dolchem/html/elem/elem017.html
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[*] posted on 27-7-2004 at 06:42


CaSO4 as the white precipitate makes sense, but two things seem to suggest that the liquid was not liquid Cl2. First, the liquid faded in color from green to red-orange over the course of 30 minutes, even while still under pressure. Second, when the jar was opened, the liquid did not start boiling, and evaporation so far has been approximately the rate of water.
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[*] posted on 27-7-2004 at 09:41


You may have formed the hydrate, substituting low temperature by pressure.

This will decompose slowly - heating speeds the process up.

This might be a very good way to clean Cl2 btw! The bisulfate/hypochlorite reaction as the FeSO4/hypochlorite suffers from the oxygen produced what makes this method unusable for most organic chlorinations and many inorganic chlorinations too.

The hydrate formation is a known way to high-purity Cl2.

Interesting anyways!




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[*] posted on 27-7-2004 at 12:38


Has anyone ever seen liquefied Cl2 btw? Does anyone have a pic of it or know where to find one? Ive searced for hours :(



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[*] posted on 27-7-2004 at 12:41


Yes and no, I have seen cyliders containing liquid Cl2, but I was quite happy that there was a layer of steel between me and it.
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[*] posted on 27-7-2004 at 13:03


I had a chem book that had the original apparatus in which chlorine was liquefied in and it showed liquid chlorine in it. It was like a distillation setup, a receiving flask connected to a distilling flask with a condenser but it was all one solid piece. One side was heated with a Bunsen burner and it caused the pressure to build and condense the chlorine on the other side of the vessel. The liquid looked just like chlorine gas itself, just a little darker.

Also, whereas Vulture has seen cylinders full of Cl2(l) I seen train cars full of Cl2(l) that pass by my area every day.




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Organikum
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[*] posted on 27-7-2004 at 18:49


If I remember right you need either 6atm pressure at 20°C, or -35°C to liquidify chlorine.

I guess the apparatus worked as told with a compression-decompression + cooling cycle. The condensor will nevertheless have to be cooled by some good cooling mixture for to reach the wanted -35°C. The compression-decompression speeds things mainly up as I believe.
Propane shoud do the trick for cooling - but who wants liquid chlorine? The hydrate should do nicely for intermediate storage purposes.

(ok, I admit, I would want liquid chlorine as soon I find a useable bottle to store it...)




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[*] posted on 27-7-2004 at 21:26


The apparatus was nothing more than a bent glass tube heated on one end and cooled in an ice-salt bath on the other. The hydrate was added and the tube was sealed. This is not an experiment on my list of things to do.
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[*] posted on 28-7-2004 at 03:11


Hmmm... 6atm isn't much, however, Im not mad enough to try it :P

Bromic, do you still have that chem book? If you do, can you please scan that pic? :)




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[*] posted on 29-7-2004 at 09:01


Can anybody provide a reaction for the hcl with hypochlorite method? I have seen five so far on this thread and do not know which one to accept. I read that Cl2 is generated when the pH falls to lower than 6, so I would assume that the whole process relies on the decomposition of hypochlorous acid:

4HOCl ---> 2Cl2 + 2H2O + 1O2

or perhaps:

2HCl + 1NaOCl ---> 1NaCl + 1Cl2+ 1H2O

I believe Organikum posted the first one, which is correct? I really hope the second one! Also what is this equimolar amounts of NaCl with hypochlorite? What is the reaction for this? I hope at least one of these does not produce oxygen.

[Edited on 29-7-2004 by Mendeleev]

[Edited on 29-7-2004 by Mendeleev]




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[*] posted on 29-7-2004 at 12:28


Just an educated guess.
1)HCl +NaOCl --> HOCl + NaCl
2)HOCl +HCl --> H2O + Cl2
So overall
2HCl +NaOCl--> H2O +Cl2 +NaCl

So I think the second one is correct as I have never heard of oxygen being produced.

[Edited on 29-7-2004 by rogue chemist]
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[*] posted on 29-7-2004 at 13:27


NaOCl in water exists as an equilibrium of NaOCl, HOCl and Cl2.

So actually both equatations are ok. The predominant reaction is the one outlined by rougue chemist, the minor but existing reaction is the HOCl decomposition. Oxygen is unavoidable when making Cl2 from bleach, some HCl is also an unavoidable byproduct. The HCl can get scavenged by bubbling through water though. Oxygen is harder to get rid of.

Therefor bleach is often not to be regarded a suitable way to produce Cl2.




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Mendeleev
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[*] posted on 7-8-2004 at 20:36


Can chlorine be suitably dried using a calcium chloride filled U-tube? If not, is 93% sulfuric acid cocentrated enough to do the job?



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[*] posted on 7-8-2004 at 20:50


Calcium chloride doesn't work very well. A wash bottle (preferably two) with 98% sulfuric acid would be suitable to dry the chlorine.

As for generating chlorine, keep things simple:

Dip HCl on potassium permanganate,

8HCl + 2KMnO4 --> 3Cl2 + 2MnO2 + 4H2O + 2KCl

or on calcium hypochlorite:

2HCl + Ca(OCl)2 --> CaCl2·H2O + Cl2.

The second method requires less HCl and is cheaper, thanks to the low price of calcium hypochlorite (in comparison to potassium permanganate).

Quote:
Theoretic: Reacting hypochlorites with acids would get you HClO and not Cl2. Mix in an equimolar amount of NaCl to your hypochlorite and use twice as much acid, that will work.

A 25% solution of HClO decomposes immediately at 0°C. If this reaction forms HClO, it will be in a much higher concentration. How much nanoseconds do you think that the HClO is going to last before undergoing decomposition?

[Edited on 8/8/2004 by Sarevok]




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[*] posted on 8-8-2004 at 21:36


I realize that chlorine can be dried using 98% sulfuric acid, but I was curious if it could be done using lower concentrations such as 93% or 90%. And why wouldn't a CaCl2 U-tube work?

[Edited on 9-8-2004 by Mendeleev]




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[*] posted on 9-8-2004 at 00:10


Calcium chloride does dry it but not completely.

And as for sulfuric acid, I am not sure but just boil it down to 98%
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[*] posted on 9-8-2004 at 02:03
Re: Large-Scale Chlorine Production


By far the most satisfactory large-scale industrial method for producing chlorine is electrolysis. Electrolysis of brine results in some evolution of chlorine, but most of it remains in solution as an alkaline solution of NaOCl, Cl2O. and Cl2, used as household bleach. For production of dry gaseous chlorine, electrolysis of a molten alkali or alkaline earth metal chloride is used. The chlorine produced and collected at the anode is usually a byproduct of the production of the metal at the cathode, most often Mg (obtained from molten MgCl2, extracted from the sea or deposits of carnallite) which is widely used in alloys with aluminium.

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[*] posted on 9-8-2004 at 10:18
Chlorine production


You do not need molten salts for copious Chlorine production from NaCl electrolysis. You add muriatic acid to a saturated NaCl solution. The solutions goes cloudy but if you add just enough H2O to clarifiy it you are OK. It can be dried through 1 or 2 calcium chloride stages. I am in the process of building a Cl generator. I will post pics when ready.

[Edited on 8/9/2004 by chloric1]




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[*] posted on 9-8-2004 at 11:24


Do this chloric1!
I am interested how it works out for you. I had no good luck trying this. Corroding electrodes (carbon from batteries), slow and unsatisfying. But I want to try it again as soon I get my hands on good and cheap carbon plates. Multiple cells seem to be a must though.

Real dry chlorine calls for 2 stages of conc. H2SO4 after all I read and after my personal experiences too. (valid also for REAL dry HCl)




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[*] posted on 13-8-2004 at 22:20
Chlorine production


I have had sucess producing chlorine via electrolysis you can read about it here:

http://thecratermaker0.tripod.com/chlorgen.htm

The chlorine produced by electrolysis is contaminated with oxygen however and unless you need a really large amount of it I would suggest just using a hypochlorite reduction of some sort to make it.

I also posed a scan of some interesting chlorine reactions up on the page

chloric1 i have NEVER seen any precipitate form when concentrated HCl is added to saturated NaCl - could it have been some impurity?
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