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Chlorine

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Organikum - 3-1-2004 at 13:05

Dont say I am nuts, but I searched and found no thread dedicated to this useful compound.

Lets open the competition:
We want chlorine - dry - clean - and in serious amounts from all over the world freely available starting chemicals.

Extra bonus if no H2SO4 is needed for drying! And pleaze not the stoneold KmNO4/MnO2 recipes.....
New ways! Transmutation rulez!

[Edited on 4-15-2004 by Polverone]

DDTea - 3-1-2004 at 13:20

Chlorine is easy and cheap to produce in the lab, but is not the easiest to dry without H2SO4. Normally, I would suggest NaCl + H2SO4 + KMnO4...but, we're talking innovation now.

Of course, you could simply add NaOCl to HCl...that's the way I've always generated Chlorine (my KMnO4 is too precious for this). Alternately, you could electrolyse HCl.

These would be the cheapest ways to actually produce Cl2, and any other way seems like it would be unnecessarily costly. Drying the gas is another issue.

BromicAcid - 3-1-2004 at 14:21

In place of sodium hypochlorite I usually use calcium hypochlorite. Also I have made chlorine gas by the reaction of HCl(aq) with H2O2(aq). Too bad both of those methods produce chlorine contamainated with water. As does the electrolysis of salt water. You could always do electrolysis of a molten chloride but the temperatures would be unacceptable most of the time, on the plus side you really wouldn't have to worry about a water contaminate.

molten chloride

Organikum - 3-1-2004 at 14:51

The Cl2 produced this way is told be rather impure - this states at least my old chem-engineering book.

I read about electrolysis of NaCl saturated with conc. HCl. Anybody tried this?

All electrolytic methods suffer from the electrode problem: Who can afford platinium electrodes or platimium coated titanium electrodes, I not.
Graphite/coal works, but for how long do the electrodes stand the chlorine attack and how does this attack contaminate the Cl2?

HCl and H2O2: Concentrations, temperatures? Pleaze...?

trichloroisocyanuric acid

Organikum - 3-1-2004 at 14:55

Any way to kick the chloride (dry preferrred) from this OTC compound? This would be nice - isocyanuric acid is a very useful compound and if this works one can get two goodies by one shot.....

(ok, the one goodie is a nastie goodie...)

unionised - 3-1-2004 at 15:46

Deacon's reaction;
4 HCl+ O2 --> 2 Cl2 +2 H2O
Needs a hot catalyst IIRC.

Cl2 from TCCA

Polverone - 3-1-2004 at 20:17

Not dry, unfortunately! But I stumbled across a patent some time ago - GB 1401120 - that stated chlorine gas could be produced by mixing trichloroisocyanuric acid with sodium chloride solution and then heating or simply reducing pressure over the solution. It would still need drying.

I mentioned in another thread that according to a journal article I read, dry Cl2 could be produced by heating dry calcium hypochlorite with cobalt salts (or cobalt/iron salts) as a catalyst.

Anhydrous CuCl2 will release dry Cl2 when heated (probably not useful on a very large scale).

[Edited on 7-1-2004 by Polverone]

Pyrovus - 5-1-2004 at 00:52

Not completely on topic, but the method of heating the copper (II) salt also works to produce fluorine. According to Reference Book of Inorganic Chemistry (Wendell M Latimer, Joel H Hildebrand) "the cupric halides decompose according to the equation: 2 CuX2 -> 2CuX + X2. Cupric fluoride decomposes around 500°C, and the chloride and bromide at somewhat lower temperatures." Unfortunately the book doesn't say what at rate this reaction occurs, so it might not necessarily be terrible useful.

KABOOOM(pyrojustforfun) - 5-1-2004 at 20:08

Quote:

Cupric fluoride decomposes around 500°C, and the chloride and bromide at somewhat lower temperatures

but I'm pretty sure CuCl2 decomposes @ 993°C!

Theoretic - 14-1-2004 at 05:30

Decomposition of nitrogen trichloride...

If_6_was_9 - 24-1-2004 at 12:54

Bleaching powder is Ca(OCl)2. We had a big jar of Ca(OCl)2 that was used for swimming pools. That was over 20 years ago and I haven't seen any Ca(OCl)2 in stores that sell pool supplies in the last 20 years. The pool "chlorine" sold now is usually trichlorotriazinetrione.

bleaching powder

Organikum - 24-1-2004 at 13:13

MERCK:
Improperly called " chloride of lime" or "calcium oxychloride". A relatively unstable chlorine carrier in solid form; a complex chemical compd of indefinite composition, presumably consisting of varying proportions of Ca(OCl)2 , CaCl2 , Ca(OH)2 and H2O in its molecular structure. Maximum available chlorine content approaches 39%. Commercial products usually range between 24% and 37% of available chlorine.

Available at every better sorted drugstore where I live.

Geomancer - 24-1-2004 at 13:26

IIRC, bleaching powder has other crap in it, too. The 'net seems to think calcium chloride. I have a big jug of stuff in my basement that claims to be 68% hypochlorite. It's called "Shock It". I bought it a couple of years ago, but I think it should still be available in the US. I suspect the impurities are the hydroxide and chloride, but I'm not sure. It generates chlorine well enough, though. The same company makes an enhanced product with (I think) a 75% concentration.

Mumbles - 24-1-2004 at 13:44

76% actually i believe. I've heard that the other 24%/32% was water in the hypochlorite's crystal structure. It wouldn't suprise me if there was Calcium Chloride and Hydroxide in it though. I am assuming it is made the same way that NaOCl(caustic soda with Chlorine bubbled in).

I know for a fact it's still available in the US. I saw it at the hardware store no more than 3 hours ago. It's made by HTH. "Shock-it" is the normal 68% and "Super Shock-it" is the enhanced one.

Hey Polverone

chloric1 - 25-1-2004 at 09:49

Quote:

Originally posted by Polverone

I mentioned in another thread that according to a journal article I read, dry Cl2 could be produced by heating dry calcium hypochlorite with cobalt salts (or cobalt/iron salts) as a catalyst.

[Edited on 1-4-2004 by Polverone]

I have seen this before and I was wondering what conditions are necessary for the catalyst. Preparation,ratios, activation what type of substrate, if any, was the catalyst deposited on. I have a small amount of cobalt sulfate and I could precipitate Co(OH)3 with hypochlorite then warm to 300C to dry to oxide. Then I would soak in Ferric Nitrate with tarce nitric acid and roast it with a large propane torch. This would deposit ferric oxide on cobalt oxide. I could also use both nitates on diatamaceous earth and perform the same roasting. Calcium hypochlorite can be bought for $10-$15 for 5 pounds in spring and summer and I could use this easy method. Also, if you could list the reference.

THanks

Polverone - 25-1-2004 at 11:45

It was a homogeneous catalyst. There was no special preparation that I recall. I originally read it in an English translation of The Journal of Applied Chemistry of the USSR, and it was from an issue from the 1950s or early 1960s. Not very specific, I know! I shall try to locate the article again.

[Edited on 1-25-2004 by Polverone]

Cl2 from TCCA

kryss - 25-1-2004 at 12:05

the beauty of using di-chloro- isocyanurates is that they generate Chlorine in Situ, avoiding all the hassle you get with handling Chlorine gas.So eg if its NCl3 your interested in making, you'd mix the DCI with a slightly ammonium salt, and hopefully get your product. IF you make up a solution of it in water over days it'll start smelling of chloramine as it reacts with ammonia in the air so can see no reason why this wouldnt work.

Theoretic - 26-1-2004 at 05:14

There's no ammonia in air.

Organikum - 26-1-2004 at 05:48

And TCCA is TRI-chloro-isocyanuric acid not DI-chloro btw.....

further adventures with chlorine

Organikum - 12-4-2004 at 08:05

After doing some exoeriments with bleaching powder and HCl and H2SO4 for producing chlorine with very unsatisfying results btw. I changed to electrolysis of HCl - with even more dissatisfaction.

-> Rolling Stones: I can get no....

F**k!

The chlorine was meant to be used for organic chlorinations (toluene...) and also for anorganic uses (Sn -> SnCl4 etc.).
Anyways, chlorine produced by the above named ways works lousy - or not at all.

Huh?
Thats exactly what I thought.
After exploiting the hypothesis that all chemistry is a fraud - there are many good reasons to believe this, I started over with research and found out.....

.....found out that all above named ways produce chlorine admixed with oxygen/hydrogen. And at least in organic chlorinations already small amounts of oxygen will fuck up the reaction. Also in inorganic chemistry the impurities seem to be at least "unfavorable".

My next try is to make chlorine from the electrolysis of zinc-chloride (actually my first idea - how funny...) which is told to produce chlorine free from oxygen and hydrogen.
We will see.
And I will tell.

Remark: Also impure chlorine is not healthy at all! Cough, cough.......

Proteios - 13-4-2004 at 16:04

pah.... lightweights!

before i turned professional.... by far the most fearsome Cl generator i made was.....

1/4 fill a wine making bottle (demijorn) with FeSO4 (cheap, freely available (by the kg!) from garden centres), gurgle common domestic bleach onto this (NaOCl soln). Cl comes off like the clappers. Get yourself a bung and some tubing, a little CaCl2 from drying and u got it made. Cl hardened all the tubing i ever got hold of as a kid. As for how it works..... meh.... never really did figure it out.... but the Fe II goes from green 2 brown..... so there something else going on other than the acidity of the Fe.

basically this outclasses any other Cl gerenator i ever made is the sheer volume of Cl that can be made, the convenience, and the controlability.

Happy hunting!

[Edited on 14-4-2004 by Proteios]

[Edited on 14-4-2004 by Proteios]

t_Pyro - 13-4-2004 at 18:03

Proteios: The reaction you stated most probably is a redox reaction between Fe2+ ions and Cl atoms in the +1 oxidation state. Fe2+ is oxidised to Fe3+, and Cl+ reduced to Cl. 2Cl->Cl2.

My choice of preparation of chlorine would be the electrolysis of concentrated NaCl solution, drying the Cl2 by passing through a drying tower packed with fused calcium chloride, and collect the gas by the downward displacement of oil, or some other non-reactive liquid. The choice of oil would also be important: if it contains too many double bonds, you might end up chlorinating the oil. Maybe benzene could also be used instead...

Organikum - 13-4-2004 at 22:55

I choose the ZnCl2 electrolysis as it produces quite pure Zn which is useful for me in other reactions and - last not least - NaCl electrolysis produces chlorine which is NOT free from oxygen/hydrogen. (after what I read on this subject).

The FeSO4 + bleach is nonsense.
The NaCl electrolysis would work but not for my purposes.

[Edited on 14-4-2004 by Organikum]

Organikum - 14-4-2004 at 00:59

Bleaching powder + water + FeSO4:

Doesnt even SMELL like chlorine.

The same bleaching powder + water + H2SO4 produced an instant strong chlorine stench, as did HCl.

Dear Proteios, only teenagers call others "lightweights" and speak like "before I went professional" - those real in the know dont have to show off like this.

Have a nice day, be a little more careful in what you post in future - somebody might get hurt following your "tips".

Organikum - 14-4-2004 at 02:33

Ok. Retried it with liquid bleach and some chlorine stench was produced this time. (from FeSO4 and bleach) but no serious amounts. I guess this is the most lousy way to produce impure chlorine known. There was no chlorine visibly evolved - ya know this green/yellow stuff.

Somehow I believe the iron might love to react with the chlorine - doesnt it?

Nick F - 14-4-2004 at 03:49

FeCl3 is available cheaply for etching PCB's, do you think this might have a use? Maybe heat would work?

Proteios - 14-4-2004 at 06:35

Lol.... so the test for purity is stench and colour....

as for your inability to get it to work..... well it worked 10 years ago when i was playin with this stuff, and working well. I would suggest tired old bleach but the fact that its gone brown suggest some reaction, but from memory this works best with adding a little bleach to a lot of FeSO4 (i seem to remember you get the most out at about a 1:1 ratio)

As for the FeII FeIII oxidation.... well yeah, like i said, i never really figured out what was goin on.... true alchemiy

Plus a suggestion... if you wanna make serious amount of Cl... dont do it in a test tube.... 5L bottles with 1/2 kg of FeSO4 work pretty well.

Im not shittin ya. It works big time.

[Edited on 14-4-2004 by Proteios]

Proteios - 14-4-2004 at 06:45

Quote:

Originally posted by Organikum
Dear Proteios, only teenagers call others "lightweights" and speak like "before I went professional" - those real in the know dont have to show off like this.

meh... the words were spoken in jest.....

Before gettin a degree in chem. and before gettin a phd in chem. I was really into this... 'what can i make from household chemicals stuff'..... really into it....that why i like this forum.....

!!mega-nostaligia!!

Nick F - 14-4-2004 at 07:50

"Before gettin a degree in chem. and before gettin a phd in chem."

Haha, there he goes again...
You've got a chemistry PhD and can't work out what's going on in that reaction?

Proteios - 14-4-2004 at 08:27

Quote:

Originally posted by Nick F

Haha, there he goes again...
You've got a chemistry PhD and can't work out what's going on in that reaction?

nope.... but i dont hold my manhood weak until someone else does!

The first response to the reaction was

'FeSO4 + bleach is nonsense'

it then improved to..... 'a poor Cl generator'.... im sure eventually we will get back to the only practical way to quickly make large controlable amounts of Cl from domestic products.

:cool:

!Happy huniting! :cool:

[Edited on 14-4-2004 by Proteios]

Nick F - 14-4-2004 at 09:13

Edit: Arrghhh, just ignore that, spent too much time thinking about copper...

[Edited on 15-4-2004 by Nick F]

Z-Row - 14-4-2004 at 12:58

Perhaps the sodium ion is grabing the sulfate ion so you have a iron(II) ion floating around which grabs an oxygen and the rest is liberated as gas.

4NaOCl + 2FeSO4 --> 2Cl2 + 2Na2SO4 + 2FeO + O2

but I'm such an amateur!

[Edit: where'd the sodium hydroxide come from?]

[Edited on 14-4-2004 by Z-Row]

Geomancer - 14-4-2004 at 13:21

Proteios: I don't have a good feeling about your proposed method. Liquid bleach is a lousy chlorine source. Solid hypochlorite should do better. You suggest a 1:1 ratio. 1:1 what?
I suspect the effect you observed was acid disproportionation of the hypochlorite combined with oxidation of the Fe(II) to Fe(III). Organikum wasn't testing for purity with his eyes and nose, just for the presence of chlorine gas. Any reaction producing the amounts of clorine you describe would produce a green cloud.

Organikum: I find it interesting that NaCl and ZnCl2 should give different products. Any reason for this? Also, are you sure that hypochlorite gives O2 or H2? I can see H2O and CO2, but these are more easily removed. Chlorine suitable for producing SnCl4 can allegedly be produced by oxidizing HCl with KMnO4 and drying the gas with H2SO4. MnO2 may work too.

Proteios - 14-4-2004 at 15:38

Quote:

Originally posted by Geomancer
Proteios: I don't have a good feeling about your proposed method. Liquid bleach is a lousy chlorine source. Solid hypochlorite should do better. You suggest a 1:1 ratio. 1:1 what?

This isnt jedi nite school......your feelings really dont enter into it..... it is demonstrable.

As for the ratios..... anything u choose....mass, volume....the principal point that i was making is that you dont have catalytic amount of FeSO4 and drown it in bleach.... from memory the reaction goes best when the FeSO4/ bleach mix is a thick goopy brown mess.

As for bleach being a lousy source of Cl....well yes n no. Yes in terms of avail. Cl per mass..... no in terms of available Cl per dollar freely available at the hardware store. Plus i think bleach is one of the purer chemical available as a domestic product.

Compared to electrolysis the current method is king. Electrolysis has a propensity to eat the electrodes in Aq. Electrolysisin molten salts is a pain in the ass. Speed of generation from electrolysis is a joke, plus the gas is tricky to collect.

ok.... this is my hack at the chemistry.....

on the left of the eqn. you know u have FeSO4 n NaOCl..... the brown stuff you get is a hydrate of one of the higher Iron oxides, or im no judge....The Na and SO4 ion remain in solution start to finish and do nothing.
FeII => FeIII +e
HOCl + e => OH- +0.5 Cl2

NaOCl ~ HOCl + NaOH

QED (?)

[Edited on 14-4-2004 by Proteios]

t_Pyro - 14-4-2004 at 19:11

Electrolysis of an aqueous solution of any ionic chloride would result in the formation of chlorine at the anode, provided the solution is concentated, and the anode inert. The evolution of O2 is due to two reasons: The decomposition of HOCl in situ, and due to a competing de-electronation of OH- ions. The former can be counteracted by using a high current density so that the chlorine evolved has little time to react with the water, and also the % evolution of O2 is less. The latter can be counteracted by ensuring that the chloride ion has a much higher concentration than the hydroxyl ions- in short, make sure that the solution is concentrated.

As for the FeSO4 HOCl bit:
Fe2+ -> Fe3+ Eo = -0.77V
2HOCl + 2e -> Cl2 Eo=+1.63V
Hence, the oxidation of ferrous to ferric, along with the reduction of chlorine should be a feasible reaction, at least on pen and paper.

Organikum - 15-4-2004 at 09:15

First I apologize for my mistake in believing that FeSO4 + bleach wont work at all.

Here I want to show some simple calculations:
- 100ml saturated ZnCl2/H2O solution contain about 235gram chlorine
- 100ml bleach 10% contain about 5gram chlorine.

So if the electrolytic cell runs down to a concentration of 50% - whats no problem at all without getting serious amounts of oxygen, the cell will produce 100gram+ chlorine. (I took into calculation some chlorine going into solution - all this is done for a worst case electrolysis) For producing 100gram chlorine from bleach 10% and FeSO4 you will need at least 2 liters of bleach - but stop - dont forget the chlorine which goes into solution in the masses of water here! So either you use much more bleach - more water - no sense, or you boil the aqueous solution and get wet wet wet chlorine which you wont dry with ease or some CaCl2.

- Carbon rods from batteries are free.
- Electricity - oh my....
- Old computer PS´s are free and everywhere.
- The zincoxide which I use is asscheap and if I would like to some HCl would be all whats needed to make new ZnCl2 from the zinc gained.
- An assembly of five jugs produces a steady stream of chlorine - all over all 1kg chlorine - from five jugs a 200ml. Absolute minimum.
- And it is fast. I will post the exact amount of chlorine per timeunit and cell after the weekend.

4,5 liters of 10% bleach equal 100ml ZnCl2 in H2O (saturated solution) - ever thought on this? :cool:

Could a moderator be so nice and edit my stupid dyslexic typo in the first post of this thread?
CHLORINE
not chloride of course.
thanks.
ORG

Proteios - 15-4-2004 at 09:58

Quote:

Originally posted by Organikum
Here I want to show some simple calculations:
- 100ml saturated ZnCl2/H2O solution contain about 235gram chlorine
- 100ml bleach 10% contain about 5gram chlorine.

ORG

1L about 50g of Cl2... about 20L of gas. Gurgling a L of bleach onto FeSO4 alsmost instantly gets you 15L of Cl..... thats enough the make most rooms intolerable. To gereate 20L of gas by electrolysis would take ..... time. As for the Cl being wet.... the water content is essentially that of the vapour pressure of water (ca 20 mg/L)....I think youll find the biggerst prob. is the redisollution of the Cl2 into the water from the bleach, which i would guess will rob you of 1/4 of your yield.
The kit i used was a 5L bottle, adding maybe as much as 2-3L bleach.....making maybe 50 L Cl2 (by these nos.).... thats by most reckonings a pretty fearsome Cl generator.

I dont doubt a reasonable electrolysis kit could be rigged... all the electrolysis kits i ever rigged worked but were impractical.... they just never made the Cl quick enough to be useful (like i say... i was a teeny at the time)

chemical Cl2 react with water... once thats happened you cant get it back by distillation. However....it dissolves in water 2 give HCl n HOCL.... the latter will react again with FeSO4. I think the best guess of a 25 % loss to the water/FeSO4/NaOCl soln is as good as any. Bleach is cheap. FeSO4 is cheap. For me... even when poor, the losses were acceptable

[Edited on 15-4-2004 by Proteios]

[Edited on 15-4-2004 by Proteios]

Organikum - 15-4-2004 at 10:11

I guess you just dont take in account the planned use of the chlorine. I by no way plan to create a wargas which makes rooms inhabitable, but I want to produce an reagent for organic and inorganic reactions. Therefor I count my chlorine in gram not in liters as it are grams to count (moles would be more correct) which are going to react. Liters is no useful unit for me here. Also I want an steady and clean stream of chlorine over a longer time.

You might read the first posts in this thread to understand what this is about - no offense we are just talking about different things I believe and this might help to clear this up.

Example. For making 250gram SnCl4 I need at least 150gram chlorine and 250gram SnCl4 is not so much at all - Friedel-Crafts reactions require rather huge amounts of metalchlorides to work. (just an example - not more)

Organikum - 15-4-2004 at 10:30

The solubility of chlorine in water is about 0,1mole per liter at 25°C so roughly 3,5 gram per liter water - basic conditions boost the solubility though.

The production of oxygen in metalchloride electrolysis is dependant on the ph - under acidic conditions almost no oxygen will be formed not so under neutral or basic conditions. Thats why NaCl isnt favorable as I believe - because of the NaOh formation raising the ph.
The presence of oxygen is not problematic for the Sn to SnCl4 reaction but it is for the toluene chlorination. ZnCl2 was choosen because of the availability of zincoxide and because I have uses for the electrolytic pure zinc lateron. I will probably also try ironchloride another day for the electrolytic iron. It is just always nice in my eyes to have double features....

Proteios - 15-4-2004 at 10:33

Quote:

Originally posted by Organikum
Dont say I am nuts, but I searched and found no thread dedicated to this useful compound.

Lets open the competition:
We want chlorine - dry - clean - and in serious amounts from all over the world freely available starting chemicals.

sure np Organikum.... I reckon u r right bout the cross purpose thing.. but from the first post FeSO4 is certainly an obscure an effective way of producing large cheap fast amoutns of Cl... V obscure!

As for war gases.... gas attacks (WW1)typically consisted of tons..... the quatities we r talking here arnt even close 2 war gas amounts.....

As for units.... g/L moles/ (bar) are all sensible units used by chemist who regularly deal with gases.

I guess if you just want small amounts of cheap, controllable, pure Cl... the simplest answer is: buy a cylinder.

Organikum - 15-4-2004 at 11:47

Looking in the catalog of big chemical suppliers I see that I "can" buy everything there I ever thought of and much more. Not regarding the fact that this is of course plain hypothetical as here where I live it is at least "not easy" to buy such things, it is against what I am after. I am after to produce what I want from scratch - scratch here in the meaning of common OTC compounds - just for to do it. It gives me a good feeling and makes the police stay away from kicking in my door and asking unfriendly questions. I had this already - I dont need this anymore rest assured.

No cylinder. Ok?

Btw. could you Proteios try to write some more understandable english - this is an very international board and it is not so easy to understand what you are wanting to express. It´s not like you are talking to your favorite drugdealer at the corner of your ghetto ya know?

Proteios - 15-4-2004 at 16:24

when i was a teeny doin this stuff, I got pally with a pharmasist.....he was kinda happy to order the stuff from the mainstream chem. suppliers....albeit...nothing too lethal or illegal....10% profit. Just a thought.....

Still one of my happier memories is making metalic bismuth from BiCO3.... which the pharmasist had loads of as some sort of stomach settling thing/antisceptic.

Good luck with the synthesis

bogus poster - 16-4-2004 at 00:06

Might oven drying the FeSO4 (dehydrating) be helpful if one wants to make Cl2 from bleach and FeSO4. Also can the bleach be concentrated by distilling some water out - the bleach of commerce available here is only 5% concentrated?
And finally: How pure is the Cl2 produced this way? It will contain water - that was told before, but if - what else?

t_Pyro - 16-4-2004 at 09:29

Organikum, I didn't understand what you meant by "metachlorides" in relation to Friedel-crafts reactions. Were you talking about the alkylation and acylation reactions using organic chlorides?

The impurities in Chlorine made from bleach will primarily be due to oxygen, as stated previously.

I see now why zinc chloride would be prefered. The decrease in ph of the solution would be counteracted by the precipitation of the zinc hydroxide, hence preventing the de-electronation of the hydroxyl ions. Sodium hydroxide, however, is highly soluble, and would therefore cause oxygen impurities. Just a thought, but does chlorine react with alkaline pyrogallol?

[Edited on 16-4-2004 by t_Pyro]

Organikum - 16-4-2004 at 10:34

Sorry - it was a typo - it should read "metal-chlorides" referring to AlCl3/SnCl4/FeCl3 to name some of the Lewis-acids commonly used in Friedel-Crafts alkylations.

The use of FeSO4 seems to me on a second thought not so bad at all - not with bleach alone but as suggested with calciumhypochlorite aka Pool-Shock and FeSO4 - this would solve the problem of the masses of bleach needed - worth a try.

Proteois: To make you happy, I was wrong. So I have to thank you for a astonishing easy and at least to me new way to produce chlorine.
Thanks.
(this new way is much more valuable as the damage done to my poor ego - I blame it on vulture and thats it for me

)

A thought: Might the addition of an acid - HCl or H2SO4 suppress the the oxygen formation at least in parts? Hm.

Somehow I really lack this "golden touch" these days.....
...nevermind.

[Edited on 16-4-2004 by Organikum]

Proteios - 16-4-2004 at 21:39

Quote:

Originally posted by Organikum
The use of FeSO4 seems to me on a second thought not so bad at all - not with bleach alone but as suggested with calciumhypochlorite aka Pool-Shock and FeSO4 - this would solve the problem of the masses of bleach needed - worth a try.

no problem

just happy to pass on the knowledge

on the above idea with the bleach powder..... looks reasonable, but may run into problems with solubiltiy/transport/mixing issues.

Ca(OCl)2 20g/100ml
CaSO4, low sol. CaOH2, v. low sol.
Fe2O3 insol.

300g FeSO4, should produce 70g Cl2. 24L.
This will produce an awful lot of insoluble stuff.... I hope that doesnt affect things too much 4 you!

Good luck n Happy Huntin!

[Edited on 17-4-2004 by Proteios]

Organikum - 17-4-2004 at 12:52

here we go....

FeSO4 straight from the box admixed with some dishwasher NaCl for to prevent clogging. It was filled into a special reaction vessel made from NYLON which was produced in my secret underground laboratories after my specifications......

(yeah - looks like a Coke-PET bottle I know... thats just a coverup... ehem..)

This is the complete setup for production of Cl2 - without drying up to now, which will be tried first by CaCl2 - if thus isnt satisfying I will go for conc. H2SO4 shudder

The also in my underground labs produced "pressure bleach injector" (looks like a bug sprayer I know...) will be filled with bleach, pressurized and will inject the bleach into the high-tekk reactor vessel where it is supposed to produce chlorine there.

First try is straight after the book - as Proteios suggested - next try will be made with adding some calcium hypochlorite to the FeSO4 to spice the reaction a little up.

Ok.
Some last wishes before I gonna kill myself/go to jail/get lynched by my neighbors ?

The_Davster - 17-4-2004 at 13:14

Why not just add some calcium hypochlorite straight into the bleach to increase the hypochlorite concentration, instead of adding it into the iron sulfate. If you add the solid calcium hypochlorite to the solid iron sulfate, and if there is any water present in either reactant or the coke bo...erm...nylon reactant vessel you will have chlorine produced without you being able to controll the rate.

Organikum - 17-4-2004 at 14:07

I will make a test before by mixing some hypochlorite with FeSO4 in a testtube to see what happens - or better how fast and violent this is going to happen.

Saerynide - 18-4-2004 at 04:28

Quote:

Originally posted by Organikum
Ok.
Some last wishes before I gonna kill myself/go to jail/get lynched by my neighbors ?

If you go to jail, be sure to tell us how the food is

I am a fish - 18-4-2004 at 05:32

Quote:

Originally posted by Organikum
FeSO4 straight from the box...

FeSO4 is pale green. It looks like yours has been oxidised into the Fe(III) state.

[Edited to fix typo]

[Edited on 18-4-2004 by I am a fish]

Organikum - 18-4-2004 at 07:42

Partially decomposed to ferric sulfate and this hydrolyzed - thats possible and probable, but it worked well so it will have been still mostly FeSO4. I have some other boxes where the FeSO4 is - as you told - of this pale greenish color. As this was just a "proof of principle and setup" it didnt matter, but you are right. I think I will post another picture after the next where the color is "right" - not to stirr confusion.

FeII

chloric1 - 18-4-2004 at 08:22

If you have Battery electrolyte(H2SO4) you can moisten the FeSO4 with it and add roofing nails to regenerate the green color if you so wish. Also, a little free H2SO4 will limit Cl2 solubility in water and may help to expel this. Just a thought. may not be practical if your time is as precious as mine.

Proteios - 18-4-2004 at 14:57

Cl2 w NaBr was a good way of making Br2 (just collecting in vial in iced water). However Cl hardened all the tubing i ever got hold of (either chlorination of the tubing, or the plasticiser), Br really eats tubing, presumably due to it propensity to condense as a liquid on almost everything.

just another happy memory

Chlorine gas production

ApprenticeCook - 19-4-2004 at 04:27

Addition funnel on RBF via a rubber tubing T-joint. Add HCl (~30%) to KMnO4 in the RBF while heating and stirring at high speed, the tube which is connected to the T-joint has a CaCl2 inline dryer then is vented to the system that requires the now pretty much anhydrous chlorine gas.
Check Rhodiums site:
https://www.rhodium.ws/chemistry/eleusis/chlorine.html
https://www.rhodium.ws/chemistry/equipment/inline.gas.dryer....

Organikum - 19-4-2004 at 04:41

Please read the first post in this thread.

BromicAcid - 30-4-2004 at 14:58

This is just a spew of information, some new, some old.

From Preparative Inorganic Chemistry
MnO2*xH2O [~100g] + 4HCl [145.88 ] ---> MnCl2 + (x+2)H2O + Cl2 [70.91g]

Concentrated, air-free hydrochloric acid (d 1.16) is added dropwise to precipitated hydrated manganese dioxide (e.g., the 86% pure commercially available material) [Note: PbO2, BaO2, KMnO4, are also listed] in a flask equipped with a dropping funnel and a gas outlet tube. The gas formation may be regulated by moderate heating.

The chlorine thus formed is passed through water then H2SO4 then CaO then P2O5 and liquefied in a receiver cooled with Dry Ice-Acetone bath.

[Yeah I know you didn't want to pass though H2SO4 and you wanted an easy prep and this was already covered]

Electrolysis of an NaCl solution with HCl in an electrolytic cell described by Bodenstein and Pohl. The oxygen content of the Cl2 produced in this manner is .01%.

From Complete Treatise on Inorganic and Theoretical Chemistry

Extremely pure Cl2 can be prepared via thermal decomposition of AuCl3.

To remove the last traces of HCl from Cl2 it is recommended to bubble though a CuSO4 solution.

You can bubble Cl2 into cold water to form the hydrate which keeps well below 9 degrees and in the dark. Heating slightly furnishes chlorine with minimal impurities.

The action of potassium or sodium chlorate on hot concentrated HCl leads to chlorine formation but if the temperature is too low then chlorine oxides are the favored product which may explode.

Heating a slurry of a chloride and a nitrate in concentrated H2SO4 furnishes nitrogen oxides and chlorine gas. The resultant mixture of gasses is passed though H2SO4 where the nitrogen oxides are retained.

Action of HCl(aq) on Ca(OCl)2 can occur with or without a binding agent for the calcium hypochlorite. Heating the reaction increases the O2 concentration of the output gasses.

Organikum - 1-5-2004 at 00:41

Thanks Bromic!

"You can bubble Cl2 into cold water to form the hydrate which keeps well below 9 degrees and in the dark. Heating slightly furnishes chlorine with minimal impurities."

This sounds like a very feasible idea for storage of Cl2 without compressing it.

A remark on the HCl and hypochlorite procedure: The use of dil. H2SO4 (battery acid) instead of HCl is advised as the sulfuric acid retains the water.

I meanwhile overcame my aversion of using conc. H2SO4 for the drying of the Cl2 - it is not so bad as it looked like and gives superior results compared to CaCl2 for example.

I am a fish - 1-5-2004 at 01:35

Quote:

Originally posted by Organikum
Thanks Bromic!

"You can bubble Cl2 into cold water to form the hydrate which keeps well below 9 degrees and in the dark. Heating slightly furnishes chlorine with minimal impurities."

This sounds like a very feasible idea for storage of Cl2 without compressing it.

Though perhaps not the most dependable way of storing it:

1. Damn! There's a power cut.
2. Why has everything near the fridge gone green?
3. What's that odour? It smells like...
4. Arrrrgh! No! Cough! Help Me! Cough! Aaaaaccckkk!!!!

Theoretic - 1-5-2004 at 09:52

When iron is fused with KNO3 (I think NaNO3 could also work) iron ferrate is formed like so:

2KNO3 + Fe => K2FeO4 + 2NO

which could be reacted with HCl like so:

2K2FeO4 + 10HCl => 3Cl2 + 4KCl + 2Fe(OH)3 + 2H2O.

NO could be recycled and turned back into nitric acid, although the process could present technical difficulties.

DDTea - 14-5-2004 at 12:13

Just skimming through this thread again, I have not seen this mentioned--this was discussed in my Chemistry textbook, and I thought it was very neat. Have you noticed that if you add even a little bit of HCl to Hypochlorite solutions, you get a flood of Cl2 gas?

This is because Cl2 is much more soluble in basic solutions than in acidic solutions (as you know, the NaOCl would be formed)... So, lowering the pH causes the Cl2 to suddenly fall out of solution.

So, here's what I propose: dissolve as much CaOCl2 as possible in Warm water. Then the addition of a little bit of a strong acid, e.g.: HCl (since we don't want to use H2SO4, and HCl is cheaper anyway). The resulting Chlorine would then be led into a separate container, immersed in dry ice or alternately, sealed and placed in a freezer over night to freeze out the water.

Cl2 gas production.

Prince_Lucifer - 14-5-2004 at 20:52

That is a very clever idea Samosa, certainly worth investigating further

I wonder if passing the generated Cl2 gas through Conc. H2SO4 would help remove H2O molecules?
Org, with the FeSO4/10%(aq)Ca(OCl)2 procedure, have you had a chance to incorporate an inline drying tube yet?
I would be interested to hear your results, and also your opinion on whether either method listed above would be a suitable Cl2 donor for toluene chlorination reactions?!
I realise what im about to say is off topic, but it is related. How adversely would Cl2 with an EXCESS of H2O, affect toluene chlorination attempts?
Good work guys, theres no shortage of smart cookies lurking around

Organikum - 15-5-2004 at 04:57

I am not sure how water will affect the chlorination of toluene - it doesnt hinder acetone chlorination, thus I know.

But oxygen which is often a byproduct of chlorine production WILL block the chlorination of toluene.

And thats the pont of my search:
A STEADY stream of anhydrous and oxygen-free chlorine.

Dropping HCl or H2SO4 onto bleach/bleaching powder/ hypochlorite will produce chlorine. This chlorine wont be evolved in a steady stream and it is wet and probably not free from oxygen.

- I tried H2SO4 and hypochlorite and bleach and FeSO4. I encountered big problems with suckback and the flow was anything but steady.
- I tried the elctrolysis of zincchloride with a graphite and an Al electrode (to deposit the Zn there). This sucks.
- I tried the electrolysis of NaCl/HCl with graphite electrodes. This work best from all tried up to now but the graphite gets eaten fairly fast from the aqueous HCl.
- I will try now the electrolysis of zincchloride with graphite electrodes and hope the Zn will separate in flakes and not deposit on the rod and shorten the cell as it happened with Al (which got also eaten up...)

I am very sure now that chlorine production for organic chlorinations is best to be done by an electrolytic method. Which electrolytic method being best I dont know by now.

Perhaps I am just an idiot, but every chemical method tried led to a severe chlorine intoxication (I have no fumehood) and ruined the reaction intended by suckback/clogging of the inlet tube and/or sudden outbursts of Cl2 blowing the shit through the condensor - all chlorinations are highly exothermic.
This is no fun at all and dangerous too. BzCl is a strong lachrymator and blowing the stuff all over you is all but funny or healthy.
The electrolytic method I am working on now is - as told - most promising. It produced fair results and gave a controllable reaction. And my lungs still worked afterwards.

Of course I will post my final results here and actually I am sureI will be able to present an "save and easy to build" electrolytic chlorine generator soon.

Up to then - patience please.

Proteios - 15-5-2004 at 09:44

solutions to suck back.
1)suck back arrestors can be purchased from suppliers.... they generally are not that good, and i would have doubts about their CL resistance

2) run with a carrier gas. The simplest is just attach an air pump to the CL generator. However if you are really keen to lose the O you can just use a N or Ar cylinder. If gas consumption is too high then you can run on close circuit inert gas. However you will need good plumbing and a good pump (chem. resistance wise). Carrier gas will also remove many of the problems of the exothermic reaction do to the lower conc. of Cl in the gas.

3) Cl liquifies low (ca -35). Some commerical freezers can get this low. Stage 1 liquify Cl. Stage 2 connect Cl ampoule to kit, and regulate Cl flow with a hair dryer.(caution... with vapour pressure! There is a pressure explosion hazard here (pressures in excess of 6 bar may be expected), but as long as the hazard is recognised, and accounted for, there should be no problem.)

4) convert Cl to bromine that is much easier to store. I dunno what you are trying to do with these clorides, but for most organic reactions bromine will work just as well. again use hair dryer to regulate flow.

[Edited on 15-5-2004 by Proteios]

Ahh, Chlorine...So Deadly, So Green

Chemtastic - 26-6-2004 at 21:41

In this thread:
http://www.sciencemadness.org/talk/viewthread.php?tid=606
Haggis posted that:

Quote:

There are many methods for generating chlorine gas, but common bleach and sodium bisulfate does the trick.

I've seen both in decently pure and concentrated solutions (12.5% NaClO to beat 5% bleach) at my local NAMCO pool store.

I'm still not sure about the reaction though...

NaClO + NaHSO4 --> ???

Anyway, isn't "wetness" of Cl2 just a physical property? If there's no chemical interaction, why not just stick your chlorine in any freezer for drying...I would think any H2O vapor would form a layer of ice on the container bottom, especially with really cold temperatures...

Liquefying the Cl2 would work too, but if you haven't any industrial refrigerators in your neighborhood, a dry ice enclosure would probably work (what's that, -80C?). Finally, just quickly transfer the liquid to a pressure-safe container, like they do with the butane in BIC lighters.

I'd still like to know exactly how sodium hypochlorite and sodium bisulfate react, as, unless someone has a warning against it, I plan to try it.

EDIT: If this works, I'll be somewhat less upset that NAMCO switched from HCl to NaHSO4 for their ph-Decrease.

[Edited on 27-6-2004 by Chemtastic]

Chemtastic - 28-6-2004 at 11:44

Today, I ended up trying a mixture of acetic acid (as 5% vinegar) and sodium hypochlorite (as 6% bleach) for the production of chlorine. I guess the reaction would go something like this:

4H+ + 2OCl- --> Cl2 + 2H2O??

I started off with about 50mL of each solution, the vinegar being clear and the bleach a pale yellow (is NaOCl solution naturally yellow, because i always see it in stores as such?). Upon mixing, there was no bubbling to indicate that anything was happening. However, the solution quickly changed from pale yellow to a moderately deep amber. I also smelled the chlorine gas odor from about 3-5 feet away.

After the solution had sat on the deck (doing this outside) for about 5 minutes, the entire mixture had turned a shade of bright pink? What caused this transition, and was it seperate from the initial change? Why was one so fast and the other so slow?

My hypothesis is that the initial reaction rapidly produced most of the possible chlorine gas, but most of this became dissolved in the solution, giving it it's amber hue. Then, over the next five minutes, the HOCl in solution decomposed in the light (though it was cloudy...) to purely dilute HCl. However, I don't know what would account for the colors.

Baking soda was added to the final pink mixture, and it did react fairly vigorously, but this could have been due to an excess of vinegar in the original solution...i only used equal volumes...nothing fancy...

Does anyone have any ideas what could have been in the solution. particularly what caused the colors and why none of the gas supposedly produced was released as bubbles?

Saerynide - 28-6-2004 at 12:01

Ahhh.... vinegar and bleach.

That was the lie I used to explain to the doctor how I got chlorine poisoned

I said I was cleaning the bathroom

Esplosivo - 28-6-2004 at 12:08

The Cl2 produced could have reacted directly with some of your vinegar forming chloroacids. The pinkish colour you stated seems strange to me - most probably impurities. Oh btw, Cl2 is soluble in water, I think heating the solution will reduce its solubility and therefore form Cl2 gas.

hodges - 28-6-2004 at 16:05

Bleach also contains some sodium hydroxide, and vinegar is a weak concentration of acetic acid, so its possible it did nothing more than neutralize the NaOH. Try adding much more vinegar than bleach and see if you get Cl2 bubbles then.

Electrolysis with a lead electrode

trilobite - 28-6-2004 at 17:52

I believe chlorine will attack a lead electrode, but what will happen after that? Will the PbCl2 formed stick to the electrode or will it drop to the bottom of the electrolytic cell, and if the former happens, will the PbCl2 layer grow until the cell resistance is so large that the electrolysis stops?

Generating chlorine gas at a steady rate iinterests me too but so does electrolytic halogenation (bromine included) and platinum eletrodes aren't my cup of tea either.

Yes, this thread is about new methods, but I'd also like to ask what is wrong with the MnO2 method? H2SO4 dropped to MnO2 and NaCl maybe? That's the old way of making bromine, using NaBr/KBr instead. Being outdoors is just great for avoiding the worst case scenarios but often there is no electricity available.

Saerynide - 28-6-2004 at 20:39

Quote:

Being outdoors is just great for avoiding the worst case scenarios but often there is no electricity available

Extension cord!!!!

ballzofsteel - 28-6-2004 at 21:38

Why not just drip conc Hcl onto your TCCA?

Theoretic - 29-6-2004 at 06:46

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.

Organikum - 29-6-2004 at 10:09

For electrolysis where Cl or HCl is formed or electrolysed you have to use graphite electrodes, for electrolysis with sulfuric acid lead the material of choice.

Hypochlorites and mineral acids produces very well chlorine. It is well referenced in the literature and was tried by myself with HCl and H2SO4, to use the acids diluted is preferred for to avoid heating up.

TCCA and HCl will produce chlorine also. The use of 20% HCl avoids that to much HCl is expelled together with the chlorine.

There is nothing wrong with the MnO2/HCl method if one has lots of MnO2 and HCl available. It is volumetric not very effective for you need 4 mole HCl to get one mole Cl2. With bleach or TCCA you get 1 mole Cl2 for 1 mole HCl, this makes your setup more compact - important so you want bigger amounts of Cl2.

[Edited on 29-6-2004 by Organikum]

sulfuric

ballzofsteel - 29-6-2004 at 18:05

Alternatively,you could mix salt with your TCCA and drip sulfuric onto it.
Wouldnt this eliminate the need for drying.
The H2SO4 would release chlorine and absorb water,whilst forming HCl at the same time,which in turn would release more Cl2 and so on.

sulfuric

ballzofsteel - 29-6-2004 at 18:08

Alternatively,you could mix salt with your TCCA and drip sulfuric onto it.
Wouldnt this eliminate the need for drying.
The H2SO4 would release chlorine and absorb water,whilst forming HCl at the same time,which in turn would release more Cl2 and so on.

guy - 29-6-2004 at 22:39

can MgSO4 be used for drying chlorine?

The_Davster - 29-6-2004 at 23:21

If it is anhydrous and not the heptahydrate, I would assume so.

Organikum - 30-6-2004 at 01:51

When you drop conc. H2SO4 onto a mixture of NaCl and TCCA you will get a mixture of Cl2 and HCl gas.
This idea was lately dicussed at the HIVE brought up by a busy bee who has obviously fallen in love with TCCA.

There is no practical way to come around the drying steps in chlorine production I believe. Usually CaCl2 is used as first step and H2SO4 as second to dry chlorine. MgSO4 or CaCl2 alone wont get the Cl2 real dry - of course it depends on what you want to do with it if this suffices or not.

Theoretic - 30-6-2004 at 09:11

"Hypochlorites and mineral acids produces very well chlorine."
How?
2NaClO + H2SO4 => Na2SO4 + 2HClO.
If hydrochloric acid is used, it will work.
Another possibility is if HClO decomposes and the resulting HCl reacts with more HClO. But that would use twice as much hypochlorite and contaminate the chlorine produced with oxygen.

Organikum - 30-6-2004 at 09:55

Theoretic you suggested:
2NaClO + H2SO4 => Na2SO4 + 2HClO

May I suggest:
4NaClO + 2H2SO4 => 2Na2SO4 + 2H2O +2Cl2
edit: I corrected the equatation after trilobite told me that I forgot something...

The prove is in the pudding. Add some H2SO4 to bleach and you will see it by yourself.
I dont want to annoy you by quoting 19th century textbooks where I discovered this method first times ago....

[Edited on 1-7-2004 by Organikum]

Decomposition of hypochlorous acid

trilobite - 30-6-2004 at 12:04

The first equation is correct in the sense that hypochlorous acid is indeed a species existing in acidic aqueous solutions of hypochlorite, but that doesn't mean hypochlorous acid is stable. The latter equation isn't balanced correctly, two chlorine atoms are missing from the right side. The truth is that hypochlorite disproportionates in acidic solutions to chlorine and perchlorate as hypochlorous acid gets oxidised by hypochlorite ions. Here are the half -reactions and the whole equation for you.

Code:

12ClO- + 12H+ + 12e-  --->  6Cl2 + 12H2O           E= 1.63V
2ClO- + 6H2O          --->  2ClO4- + 12H+ + 12e-   E=-1.38V
-----------------------------------------------------------
14ClO- + 12H+         --->  6Cl2 + 2ClO4- + 6H2O   E= 0.25V

Or in other words, so that the role of hypochlorous acid becomes obvious:

Code:

12HClO- + 12e-  --->  6Cl2 + 12H2O           E= 1.63V
2ClO- + 6H2O    --->  2ClO4- + 12H+ + 12e-   E=-1.38V
-----------------------------------------------------
12HClO + 2ClO-  --->  6Cl2 + 2ClO4- + 6H2O   E= 0.25V

Also, when chlorine is dissolved in water or aqueous sodium hydroxide, the following happens

Code:

Cl2 + H2O <---> HOCl + HCl

So the function of the sodium hydroxide is in fact solvate the hypochlorous acid by neutralizing the acids shifting the equilibrium to right and preventing the decomposition. However, no perchlorate would be formed in those conditions as those chloride ions are left in the solution.

Code:

2ClO4- + 12Cl- + 8H+ ---> 7Cl2 + 8H2O

Thus, if one wanted to make chlorine from pure solid hypochlorites, he might want to add some chloride salts too.

[Edited on 1-7-2004 by trilobite]

Chemtastic - 30-6-2004 at 17:27

Quote:

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.

So you would recommend 1 mole of Ca(OCl)2 to every 1 mole of NaCl to every 2 moles of the acid used, Theoretic?

Also, no one ever answered my question about drying the Cl2 produced. Wouldn't it be possible by either lowering the temperature below 0C and removing the ice or lowering the temperature below -35C and removing the liquid Cl2?

unionised - 1-7-2004 at 12:08

Chlorine will tend to form the hydrate under those conditions (ie cold and wet).

MgSO4, CaCl2or H2SO4 or quite a lot of other things can be used to dry chlorine.

Organikum - 1-7-2004 at 12:39

Just use HCl and TCCA or a hypochlorite. Wetten your calcium hypochlorite before dropping diluted HCl on it. Dropping concentrated HCl will liberate shitloads of HCl with the Cl2.

Hmmm...

Chemtastic - 1-7-2004 at 13:31

I never knew that gases had hydrated forms...

Proteios - 2-7-2004 at 05:05

loads of gases form hydrates....clatherates.....methane, SO2, all the noble gases. The conditions are usually kinda quirky, but gases forming hydrates is not uncommon. SO2 is somewhat different in that it will both react with water, and form clatherates. I dunno about Cl2, but see no real problem in this forming clatherates too.

Theoretic - 2-7-2004 at 10:56

"May I suggest:
4NaClO + 2H2SO4 => 2Na2SO4 + 2H2O +2Cl2
edit: I corrected the equatation after trilobite told me that I forgot something..."

Oh no. You've corrected it and it's still unbalanced! Two oxygen atoms are missing from the right side of the equation.

"The prove is in the pudding. Add some H2SO4 to bleach and you will see it by yourself."

True, this is because bleach also contains NaCl.

"I dont want to annoy you by quoting 19th century textbooks where I discovered this method first times ago...."

Well, the textbooks apparently said so because the disproportionation of HClO by hypochlorite happens.

Hold on... Perchloric acid!?! Can someone please provide details? Rate, optimal temperature, side reactions?

"So you would recommend 1 mole of Ca(OCl)2 to every 1 mole of NaCl to every 2 moles of the acid used, Theoretic?"

What I meant is 1 mole of chloride ions to 1 mole of hypochlorite ions. 1 mole of Ca(ClO)2 compound, two moles of NaCl and one mole of H2SO4.

"Also, no one ever answered my question about drying the Cl2 produced. Wouldn't it be possible by either lowering the temperature below 0C and removing the ice or lowering the temperature below -35C and removing the liquid Cl2?"

Well, you could use the CaSO4/ CaCl2/Na2SO4/NaCl mixture you made when producing chlorine

. You could use anhydrous CaSO4 or CaCl2.

[Edited on 3-7-2004 by Theoretic]

hodges - 2-7-2004 at 16:12

Quote:

Originally posted by Theoretic
"May I suggest:
4NaClO + 2H2SO4 => 2Na2SO4 + 2H2O +2Cl2
edit: I corrected the equatation after trilobite told me that I forgot something..."

Oh no. You've corrected it and it's still unbalanced! Two oxygen atoms are missing from the right side of the equation.

I've seen a discussion of this reaction in Usenet before, although I don't remember where. As I recall, some O2 is also produced by the reaction, as well as the Cl2. Commercial bleach is made by bubbling chlorine through NaOH and thus contains NaCl as well as NaOCl. Thus when you acidify regular bleach you don't get any oxygen, because of the extra chloride. Also I don't believe NaOCl is very stable by itself so in practice its usually going to be mixed with NaCl.

kryss - 3-7-2004 at 01:24

Trilobite your getting mixed up hypochlorite disproportioates into chlorate not perchlorate! Hence if you heat a given quantity of bleach you get one third of its equivalent as chlorate, the rest as chloride.

Then if you go on to melt the solid chlorate if further disproportates into perchlorate and chloride.

Theoretic - 3-7-2004 at 13:16

Kryss, it is true that hypochlorites disproportionate into chloride and chlorate BY THEMSELVES, but molecular hypochlorous acid has properties WAY different from the hypochlorite ion, so that reaction has the right to differ.

S.C. Wack - 3-7-2004 at 14:13

Not much different:
5HClO = HClO3 + 4HCl + O2 or
3HClO = HClO3 + 2HCl and
HCl + HClO = H2O + Cl2
2HClO = 2HCl + O2
6HClO + NaCl = NaClO3 + Cl2 + 3H2O

A quote from Cotton and Wilkinson:
"In general, the chemistry of these acids [halogen oxo acids] and their salts is very complicated."

Yes hypochlorous acid can be made from bleaching powder with HNO3, HCl, and best with boric acid, but there are problems as you see. And not in equimolar amounts, even in very dilute solution: 4NaClO + 4HCl = 4NaCl + 2Cl2 + O2 + 2H2O

It looks like CO2 is a strong enough acid to give Cl: 2Ca(ClO)2 + 2CO2 = 2CaCO3 + 2Cl2 + O2.

chemoleo - 3-7-2004 at 20:20

On that note - ever wondered why calcium hypochlorite still smells of chlorine? Is it because of the CO2 reacting with it? I bet the only way to get Ca(OCl)2 stable is to keep it under vacuum....

Pyrovus - 4-7-2004 at 01:43

Commerical Ca(OCl)2 generally comes in a hydrated form, so the chlorine smell might come partly from interaction between the water of hydration and the hypochlorite ions:
OCl- + H2O <-> HOCl + OH-
With subsequent decomposition of the HOCl.

kryss - 4-7-2004 at 15:40

Kryss, it is true that hypochlorites disproportionate into chloride and chlorate BY THEMSELVES, but molecular hypochlorous acid has properties WAY different from the hypochlorite ion, so that reaction has the right to differ.

Ordinary bleach is stabler in alkaline solution but slowly goes off, especially in sunlight.

Hypochlorite solution as in pure NaOCL is less stable and more light senstive.

I think this is all tied into the amount of HClO present as it is light sensitive (UV) and heat sensitive.

I think 3OCl ->ClO3 + 2Cl-

Although you lose chlorine you keep all its oxidising power.

For HCLO its different:

HClO -> HCl + 0.5 O2
and
HClO + HCl ->Cl2 +H2O

I have read that an azetrope can be distilled off under vacuum, think you need to distill HgO with OCl- - might be Chlorine monoxide though.

Michal - 11-7-2004 at 04:26

I have just recieved my order form the pottery store, and I wanted to try and make Cl2.

I used MnO2 and HCl 10% sol.

I added a little bit of MnO2 in a reaction tube, and added the HCl sol.

But nothing happend, there was nog smell of chlorine at all

What do you people think is the problem ?
Is 10% HCl to weak to produce chlorine from MnO2 ?
Maby I added to mutch HCl beqause it was a HCl/MnO2 solution, not a powder with a few drops of HCl.
Also, I used my MnO2 straight from the bag, maby it has to be activated with a weak HCl solution ?

If you use HCl it will oxidise the Cl- to Cl2, so would'nt NaCl solution work ?

Thx for your time to read these questions

Edit: Hurray, finaly there was a smell of chlorine, but not really strong (if I smelled 5cm away from the reaction tube) it smelled the same like 5% NaOCl solution.

When heated the smell became stronger

Will KMnO4 work better in producing chlorine ? But I will propably stick to my MnO2 beqause it's dirt cheap -> 1kg = 1,92 euro

[Edited on 11-7-2004 by Michal]

Reverend Necroticus Rex - 11-7-2004 at 11:33

I never had much success using MnO2 to release Cl from HCl, KMnO4 does a brilliant job of liberating chlorine, so much so that it's a good idea to be careful as to how much you add, I make ALL my chlorine gas this way

Organikum - 12-7-2004 at 09:09

MnO2 and HCl produce chlorine mainly upon soft heating. This is actually the advantage of the method, you can get a constant stream of chlorine this way and have no instant Cl2 outbreak upon addition.
The HCl added to the MnO2 should bo between 15% and 20% in concentration. To low a concentration can produce explosions in the worst case (its not very probable but it has happened). More than 20% will give you lots of HCl-gas togehter with the Cl2 as the Cl2 produced will drive it out. The MnO2 used has to be wettened with water before putting it in the reaction vessel.

ORG

Acid-Hypochlorite

Chemtastic - 24-7-2004 at 10:47

As recommended to me a few weeks ago, I finally got around to trying the method on Meglomania's website, using calcium hypochlorite and sodium bisulfate. It worked REALLY well, even at room temperature and without too much of either reactant. The biggest benefit was finding both reactants in crystalline form. This meant they could be mixed with no reaction, forming the deadliest "just add water" mix I've ever used.

I did get an interesting repeat of results. 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 solution itself seperated into two layers. A gunky, white precipitate settled to the bottom, leaving a pale green (and CLEAR) supernatent. I don't know what the white gunk is, but could it be Ca(OH)2? The green fluid I think is Cl2 dissolved in water. This was the same color as the vinegar and bleach experiment from before, and it shared another similarity. After some time in the sunlight, it similarly changed in color from pale green to pale red. What color is HClO in solution, since Cl2 + H2O --> HClO + HCl, and HCl is clear? Also, if anyone knows the reaction between sodium bisulfate and calcium hypochlorite, what is it?

guy - 24-7-2004 at 20:57

The white precipitate is probably Calcium sulfate which is only moderately soluble.

[Edited on 25-7-2004 by guy]

[Edited on 25-7-2004 by guy]

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