Sciencemadness Discussion Board

Chlorine

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BromicAcid - 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.....

guy - 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.

unionised - 25-7-2004 at 14:14

Have you quoted the bit you thought you had?

[Edited on 25-7-2004 by unionised]

guy - 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]

Saerynide - 25-7-2004 at 23:24

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

guy - 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

Chemtastic - 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.

Organikum - 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!

Saerynide - 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 :(

unionised - 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.

BromicAcid - 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.

Organikum - 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...)

S.C. Wack - 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.

Saerynide - 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? :)

Mendeleev - 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]

The_Davster - 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]

Organikum - 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.

Mendeleev - 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?

Sarevok - 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]

Mendeleev - 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]

guy - 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%

Re: Large-Scale Chlorine Production

JohnWW - 9-8-2004 at 02:03

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.

John W.

Chlorine production

chloric1 - 9-8-2004 at 10:18

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]

Organikum - 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)

Chlorine production

Pyrophoric - 13-8-2004 at 22:20

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?

darkflame89 - 14-8-2004 at 00:38

That precipitate mentioned must have been the NaCl precipitating out of the solution due to the ultra high concentration of chloride ions. You could try again with a higher concentration of NaCl or a higher concentration of HCl(which i doubt).

Pyrophoric - 14-8-2004 at 05:59

When i used the words "i have NEVER seen" i was trying to imply that although i thought it theoretically possible for a precipitate to form, my observations supported that it were not probable (at least with common salt and acid).

The salt solution i was using was saturated at room temperature (crystals forming on the surface) and my HCL measures sg1.1+-.01 which corresponds to ~20-24%HCl.

Saturation at 20*C NaCl is 35.9g per 100g H2O or 35.9% this corresponds to .614mol per 100g H2O of Cl- at saturation.

My HCl at an approx concentration of 22% would give a chloride concentration of .603mol per 100g H2O and so irrespective of the amount added could never bring the concentration of chloride ions above the saturation value of .614mol per 100g H2O! :cool:

Of course it might be possible to precipitate NaCl if using a higher concentration of acid - not likely to be obtained OTC (at least in my area) unless purchased from lab supplies!

Chlorine?

chloric1 - 21-8-2004 at 11:46

Not sure I want to call it a complete success but I am running the generator now. The output is not nearly as I hoped it would be. I am using 5 liters of liquid which is now been adjusted to mostly HCl. Secondly, I feel I do not have complete gastight seals so i baught some better adhesive. I will try it out later this afternoon and wait til it cures(24 hours). I definately notice the odor but it must be small amounts because it does not even burn my eyes! I am running at 8.3 amps for crying out loud:mad::o! I will give it 20 more minutes to overcome the volume of liquid.

Could use unglazed pottery as a diaphragm to concentrate chloride ions at the anode. Most pottery would need a run with epson salts to fill some pores with Mg(OH)2. This may help me control the CL2 and channel it where I need it. Sure would be nice to be able to make homemade hypochlorites and chlorinate chloroform etc. etc:D

[Edited on 8/21/2004 by chloric1]

Nevermind

chloric1 - 21-8-2004 at 12:41

Project abandoned! Power supply shorted out!:mad::mad::mad:!! That was a $100 supply! I thought it had current overload protection. I will have to buy another supply. So, no electrolysis or electronic expereiments for at LEAST a month maybe more. :(:(

I finding out that chemistry is more satisfying in books than in real life! Might just chunk all my gear and stick to the books only! I will just sign up for a chemistry program at the University and experience that way. I surrender:(:(

JohnWW - 21-8-2004 at 13:06

Either way, that sounds expensive.

John W.

Tell me about it

chloric1 - 21-8-2004 at 13:52

I would try to sell what I could. I just cant dump nay of my chems.

As far as chlorine goes I think I will just set up ground glassware and use hydrochloric acid with permanganate. I only have about 50% of my ground glassware and I really can't use it yet. IT cost so much if I break something, I am down for weeks or even months:o! I maybe just want to go to a University or work in a lab that way I can damage someone elses property:D

Organikum - 22-8-2004 at 02:09

KMnO4 + HCl, MnO2 + HCl, TCCA + HCl, these work just fine, for reasons of availability I prefer the latter method.

Not dripping the HCl ON, but running it to the bottom is advised with TCCA.

Pitch is a perfect sealant for these gasgenerators. PLASTILIN works well as a quickfix - good to have some by hand I learned.

Chlorine

mick - 22-8-2004 at 09:28

I am sure the comment was only in p****d off mod, but please do not go to university to damage someone elses equipment because other people, especially technicians, also use the equipment and if it is left damaged they can be in danger.
One example of many is someone boiled the 2 gallon acetone recovery still dry out of hours (I plug it into a standard time so after 11/2 hours it turns of if I have forgotten about it and only ever run it to just below half full). They made an attempt to clean some of the black stuff in there and refilled it to set it up for another run. It was obvious thing were not right, after emptying the flask you could see the bottom of the flask had a ring of cracks and looking like a shattered windscreen. Luckily they did not do me for harrisment and threatening behavior after I explained to them against the wall that
1) If you do not know what you are doing do not do it
2) They put me in serious danger
3) There was enough solvent there, if it caught fire, to trash the whole of the labs.
4) The most important was, if the labs were trashed I would be out of a job. There is no way that the place would be refitted as labs because of all rules, regs and costs.

Mendeleev - 26-8-2004 at 18:52

I realize Cl2 is toxic, but just how much. Say about 100 L of Cl2 at 1 atm blows out of an apparatus and is dispersed equally throughout a .5 acre lot how dangerous is it to breath the air there, or is it dispersed too much to be dangerous? I'm not expecting to blow out 100 L of chlorine, but just hypothetically.

BromicAcid - 26-8-2004 at 19:23

I would say that would be quite dangerous. I should convert to ppm and compare to the accepted threshold but I know that would be dangerous.

A liter or two over a .2 sq acre area is excessive as I've had 'spills' this size, the clouds kill the grass at this level in areas and some trees, it will hurt your eyes, make you cough, and possibly cause permanate damage if you are not quick to act.

But can someone give a more specific answer. 100 L over .5 acres would seriously seem like trouble to me.

Mendeleev - 26-8-2004 at 20:36

I did some very rough calculations and if you have a 30m x 30m plot and 1 mol (22.4 L) Cl2 is distributed equally through a height of 2 m, that would be roughly 40 mg/cubic meter. Which would make 1 L only 2 mg/ cubic meter. I don't know how harmful 2 mg/cubic meter. If your property was half the size I calculated then it would be around 4 mg/cubic meter. Are you sure you only had one liter of this stuff, 4 mg/cubic meter doesn't sound that bad, as I read the lowest lethal concentration (LCLO) in humans was 2530 mg/cubic meter, although I may be mistaken. Here's here's the msds with Othe LCLO msds:

http://ptcl.chem.ox.ac.uk/MSDS/CH/chlorine.html

[Edited on 27-8-2004 by Mendeleev]

Organikum - 26-8-2004 at 22:28

Chlorine is toxic - yes, but it is to handle easily with good ventilation. A room in the basement without window would be a bad place though to do these experiments.

To kill oneself with chlorine is a hard task, one runs away before it kills you - the stuff hurts.
In WWI chlorine could only be successfully used as a war-gas because they were shooting with machine-guns above so nobody could get up and out of the trenches.

Most gas is evolved during the disassembling of the generator/drying setup. A additional inlet on the generator where a aquarium-airpump or similar can be connected and the system purged before disassemblation circumvents this problem and is advised.

The positive point of chlorines toxity/smell is that you know soon when you have a leak. Have some pitch or PLASTILIN or clay by hand and solve the problem in seconds. Forget hot-glue or other glues and silicone.

Chlorine is easily scavenged OVER lye. All in all it is good to work with, better than with dry HCl. Of course the setup needs some testing and trying - these things never work perfect on first shot.

Contemplation

chloric1 - 19-9-2004 at 10:33

Well, I have gotten over my hissy fit and had some time to contemplate. I realize I will have to redisgn my apparatus. I may go smaller scale and try a divided cell so I don't have an electrical short. Next time I will use a porous ceramic as a membrane. Aluminum oxide is chemically inert but suffers alkaline attack I think. Hmmm! most ceramics do have some action with alkalies and phosphoric acid. What to choose?:o:o

I had a nice idea typed up....

stygian - 19-9-2004 at 19:54

but it seems my username expired :( Here we go again.

Consider the following.

2H3BO3 + 3Ca(OCl)2 ---> 3Cl2 + 3H2O + Ca3(BO3)2 (solubility anyone?)

One mixes the powdered chems well, puts into standard 2L gas generator, and initiates the reaction with a bit of water. This way should have a much higher Cl2:reactant volume ratio than methods using HCl (max ~37%?)

Organikum - 26-9-2004 at 12:46

The best method I is TCCA + dil. HCl.
Using the tabs for swimmingpool desinfection as is, is great.

No heating up.
No suckback.
The tabs dissolve not too fast, so everything stays under control.

Ahhh,goverment takes it away and industry gives better things back.

Love it.

ORG :D

Marvin - 26-9-2004 at 13:22

stygian,

A good attempt, but if you count oxygens you'll see they dont balence.

The reason is that you are trying to turn chlorine in oxidation state +1 in hypochlorite into oxidation state 0 in chlorine without a reducing agent. The only way it would work would be to partially decompose the hypochlorite into chloride but then all you are really doing is making HCl a different way.

Using Ca(OCl)Cl , or a mixture of hypochlorite and salt in some water and adding sulphuric acid could work, but then you have the heating problems Organikum was talking about. Volume of gas generator shouldnt be much of an issue unless you are making massive amounts.

Organikum - 26-9-2004 at 13:39

Hypochlorites (bleaching powder) and dil. H2SO4 (battery acid 38%) works fine.

Its just that 1 mol hypochlorite provides just one mol Cl2, but TCCA and HCl provide 3 mol Cl2 per 1 mol TCCA and 3 mol HCl.
And the slowl dissolving tablets make TCCA so very nice.

Chlorine

MadHatter - 26-9-2004 at 13:52

Chlorine is toxic - no doubt about that ! Some of it evolves during my
electrolysis of chlorides into chlorates and subsequently perchlorates.
I live in an apartment(flat) and can easily vent the excess Cl2 into the
atmosphere without problems. I can still smell it and at most it's an
irritant ! Just be sensible about how you handle this 1. As for producing
Cl2 for use from electrolysis, you should
fashion a U-shaped apparatus that will not allow the Cl2 to re-combine
with the brine. If you're using NaCl as the chlorine source you should
get chlorine on 1 electrode and a solution of NaOH on the other.

garage chemist - 26-9-2004 at 14:27

Organikum, the reaction between TCCA and an acid sounds VERY interesting as Calcium Hypochlorite is not available at pool stores where I live. (I've got one kilogram of chlorinated lime, but I have no supplier for it, so I don't want to use it up).
TCCA however can be got quite easily.

Could you write down the equation TCCA + HCl? I know almost nothing about the chemistry of TCCA, only that it somehow slowly produces chlorine or hypochlorite in the water of swimming pools.
Perhaps we should start a thread about the reactions and chemistry of TCCA?

I just don't feel well producing chlorine with a reaction I can't formulate. ;)

No problem...

Organikum - 26-9-2004 at 16:03

TCCA and HCl:

C3Cl3N3O3 + 3HCl = 3Cl2 + C3H3N3O3

TCCA: 232,14 g/mol
HCl: 31% muriatic acid contains about 10 mole Cl per liter (5 mole Cl2)

So take at least 350 ml HCl (31%), better more and dilute it to 15-20%. Add this to 250g TCCA and you will get netto/effective after drying and losses 3 mole Cl2 (70,91 g/mol), thats over 200 gram.

When the HCl is icecold outa the freezer then you can dump the acid over the TCCA, as the evolution of the Cl2 is not sooo fast.

Hey, thats the hit, no shit! :D

garage chemist - 27-9-2004 at 07:40

Thanks a lot, Organikum!

Now I will buy some TCCA when I'm at "The Home store" next time.

However, HCl is only sold at max. 25% concentration where I live and its rather expensive.
I'm sure that sodium hydrogen sulfate will also work to produce Cl2 from TCCA because it's an acidic salt and its solution in water acts like a mixture of dilute sulfuric acid + sodium sulfate.

The reaction should then go as follows:

2 C3N3O3Cl3 + 6 NaHSO4 ----->
2 C3N3O3H3 + 3 Cl2 + 3 Na2SO4

Is this right?
When yes, I now have access to virtually unlimited quantities of Cl2 for my production of chloral hydrate. :cool:

Another idea, perhaps not so related to the topic:
In pool water, TCCA reacts with water to form hypochlorous acid, which serves as the disinfectant.

C3N3O3Cl3 + 3 H2O ----> C3N3O3H3 +
3 HClO

And now my idea: Could this reavtion somehow be used to convert acetone to chloroform using the haloform reaction?
If yes, we all would have a much better synthesis for chloroform which doesn't use such excessive amounts of bleach! :o

garage chemist - 29-9-2004 at 06:12

Now I've got a kilo of TCCA and it makes lots of greenish bubbles when it's mixed with HCl. :)

However, it doesn't produce chlorine when NaHSO4 soln is added! :o I don't know the reason, but there is only a slight evolution of colorless gas.
I've got no idea what this is!
So my equation with NaHSO4 is wrong.

But, I think that when NaHSO4 soln is mixed with NaCl soln, the resulting soln contains both H3O+ and Cl- Ions and should therefore also work for producing Cl2 from TCCA.

Organikum - 29-9-2004 at 06:18

TCCA is practically inert to sulfuric acid in any concentration.
You can of course make HCl in any of the known ways and vent this to the TCCA.

The concentration of the HCl doesnt matter, 25% is ok as written in my last post. HCl shouldnt be expensive anywhere as it is actually a waste product, look around you for sure will be able to localize a cheaper source.

garage chemist - 29-9-2004 at 08:00

I can get HCl cheaper and in higher concentration, but only via mail order, and there are always the shipping costs.
I prefer using the 25% HCl as it is available from the Home Store.

I have posted the equilibrium reaction of TCCA with water before.
When NaOH is added to TCCA, would the equilibrium be shifted to the right (formation of Hypochlorite)?
I'm going to try this out!
It would be great if this is an alternative source of hypochlorite!

Theoretic - 29-9-2004 at 08:58

How about aqua regia? The "aqua regia " equilibrium gives you chlorine, like so:
3HCl + HNO3 <=> NOCl + 2H2O + Cl2.
Heating aqua regia should convert most of HCl to Cl2. To get more bang for your buck, use an excess of whichever acid is more available, this will shift the equilibrium to the right.

garage chemist - 29-9-2004 at 09:33

Half an hour ago I've mixed TCCA and NaOH in water while keeping the mixture cool, then I added a few drops of acetone and the mixture became cloudy.

The I smelled the reaction mix and was greeted with the lovely smell of chloroform!
:cool: :cool: :cool:
A new, high- yielding OTC method of making CHCl3!
No large amounts of bleach needed!

TCCA contains LOTS of chlorine and this reaction converts it ALL to chloroform!

I will elaborate a complete synthesis of chloroform from TCCA.
If anyone's interested, I can write it down when I'm ready.

UhhKaipShaltaBlet - 17-2-2005 at 07:26

Excuse me of my ignorace ,but what's "TCCA"?

neutrino - 17-2-2005 at 14:01

Trichloroisocyanuric acid, a common pool chlorinating chemical.

chloric1 - 18-2-2005 at 03:06

Quote:
Originally posted by garage chemist
Half an hour ago I've mixed TCCA and NaOH in water while keeping the mixture cool, then I added a few drops of acetone and the mixture became cloudy.

The I smelled the reaction mix and was greeted with the lovely smell of chloroform!
:cool: :cool: :cool:
A new, high- yielding OTC method of making CHCl3!
No large amounts of bleach needed!

TCCA contains LOTS of chlorine and this reaction converts it ALL to chloroform!

I will elaborate a complete synthesis of chloroform from TCCA.
If anyone's interested, I can write it down when I'm ready.


That is awesome! I would like to try the TCCA & NaOH mixture on ammonia mixed with acetone or methylethyl ketone. A little gelatin wouldn't hurt for cautionary measure. I just need to find TCCA in a container small enough to be economical. How well does TCCA dissolve in water and in alkali anyway?

garage chemist - 18-2-2005 at 03:48

I'm sorry, I wasn't able to produce significant amounts of chloroform from TCCA and, to be honest, I stopped the research because a supplier started carrying chloroform and I bought 500ml for 10€.
That's more than I will ever use...

But it's still possible to make lots of chloroform from TCCA by adding HCl and bubbling the resulting chlorine gas into ice-cold dilute NaOH. This produces NaOCl in a concentration depending on the starting concentration of NaOH. Up to 20% NaOCl are possible. Upon reacting this with acetone, chloroform will be had in high yield.
This method converts ALL chlorine from the TCCA into chloroform.

Making NaOCl from chlorine + NaOH would also be highly advisable for hydrazine production, because my chloroform attempts directly from TCCA failed and this indicates insufficient NaOCl concentration in the TCCA + NaOH mixture.
Don't try to make hydrazine with that!

Instead, build a fume hood / get a gas mask / work outside when its windy, produce chlorine from TCCA and HCl and bubble it into cold NaOH. The resulting NaOCl solution is most likely much purer than the storebought pool chlorinator, this can improve yields with hydrazine production considerably.

Wait a minute!

chloric1 - 18-2-2005 at 03:50

Now I know pool suppliers also sell sodium dichloroisocyanurate. Would this be your reagent from TCCA and NaOH? Or maybe its not that simple. Maybe the alkali hydrolysizes the TCCA to release free NaOCl. I am not familiar with the oxidizing species involved here though. I once tried a Sodium Dichlor oxidation of sodium bromide and only got a yellow precipitate.:( No free bromine. Anyone a pool expert out there? I would like to hear some input on the stabilized chlorinators

Electrolytic Chlorination

Eclectic - 18-2-2005 at 05:21

Is anyone familiar with a method of producing chloroform by electrolysing aqueous CaCl2 and acetone? I saw this a few years ago but I'm having trouble locating the details.

[Edited on 18-2-2005 by Eclectic]

frogfot - 18-2-2005 at 05:48

GC, what was the yield of chloroform?

Lately I've been in demand of big amounts of cheap chlorine, in synth of things like S2Cl2. Since calcium hypochlorite here is 40$/2 kg, TCCA even more, KMnO4 is 11$/0,75 kg and MnO2 is 8$/kg..

...therefore I wanna finally try the electrolytical way.. Since I like mspaint here I've drawed the theoretical setup I'll use. It suppose to work continiously.

http://img.photobucket.com/albums/v113/frogfot/chlorineelapp...

I've tried to keep it as simple as possible and doable with OTC parts.
So the electrolyte in anodic cell will be continiously refilled by KCl soln so it will have slightly higher level of liquid than in cathodic cell.
The membrane will be composed of glass wool (glassfiber blankets). I dunno if some filler/binder will be needed.. probably yes..
Important thing is to keep appropriate current dencity above membrane, so it wouldn't boil from inside (had some bad experience with salt bridges..).

9V would be probably enough.

Everything seems to be quite easy to make. Anyone see some obvious mistakes in the design?

[Edited on 18-2-2005 by frogfot]

[Edited on 18-2-2005 by frogfot]

Esplosivo - 18-2-2005 at 07:31

Nice method to produce chloroform, nice and cheap. The chloroform though being organic would probably result in the remains of hydrolysis of TCCA, or excess TCCA being dissolved in it though. Distillation might be enough to seperate them, but the organic products of hydrolysis might be as volatile or nearly as chloroform and share the same boiling point (although I don't think so). Am I correct in assuming that the chloroform synthesized in this way will contain a lot of contaminants?

[Edited on 18-2-2005 by Esplosivo]

garage chemist - 18-2-2005 at 08:13

I only got some white precipitate and no liquid chloroform. Actually, the smell of chloroform disappeared rapidly from the mix...
You can experiment if you want, maybe you've got more luck than me.

However, I would convert the TCCA into NaOCl via chlorine (from TCCA + HCl) and NaOH and then add acetone to this. This reaction is much cleaner, the produced chlororform doesn't need to be distilled and the yields are much higher, anyway.

@ frogfot:
Your cell looks good. But the chlorine production will be slow. To make S2Cl2, you would have to run it a long time and keep the sulfur molten the whole time.
Or you could condense the chlorine into a liquid with dry ice. If you can get dry ice, this would be ideal. Then you can let the chlorine boil off at the required rate.

The MnO2 seems cheap though. Or is it still too expensive for you?

[Edited on 18-2-2005 by garage chemist]

Boiling point

chloric1 - 18-2-2005 at 10:16

You are right about liquifying chlorine at dry ice temperatures. Actually -30F should work. But, I should mention that the plain ice will not have enough contact surfaces with your receiver to chill it. You need to add it to a nonfreezing liquid until the liquids temperature approaches that of the ice itself. Try a nonflammable solvent like trichloroethylene. Anywho just realized I have no way to measure extreme cold myself so I need to find me a voltmeter with a thermocouple!:o Off to Jameco I go!

frogfot - 18-2-2005 at 12:02

Liquifying chlorine is a bit scary. I do smelly things in homemade fume hood that could vent an acccidental leakage from a normal chlorine generator.. but accident with liquid chlorine would be too much for it..

As for speed, at 10A it would give me 1,16 ml Cl2/s (at 100% efficiency) which is suitable in most experiments..

[Edited on 18-2-2005 by frogfot]

evil_lurker - 11-5-2005 at 16:18

Anyone ever tried to build a DeNora cell?

Basically it uses brine as an electrolyte, graphite as the anode, and mercury as the anode.

When electricity is passed thru the cell, Chlorine gas is formed at the anode, and Sodium is evolved at the cathode. The sodium forms an amalgam with the mercury, which is then drained off and placed into water. The sodium reverts into sodium hydroxide and the mercury is reloaded into the cell.

The reaction is:
Hg
(Cathodic reduction) Na+ + e- ------> Na (Amalgam)


2Na (amalgam) + 2H20 ------> H2^ + 2NaOH

12AX7 - 11-5-2005 at 16:55

I've always wondered how you keep the Na from reacting with the electrolyte; or does the voltage tend to keep that "sealed in"?

Aside from a small contaminant of mercury, it makes very high purity lye industrially, so I've read. Seems to me it would take a lot of mercury and be kind of complicated in a continuous process (not necessary of course, but with only 2% Na (also as read), almost a necessity) though.

Given this thread is about the other product, it might work for that purpose, just loading the mercury with sodium until it starts fizzing or something.

Tim

garage chemist - 19-5-2005 at 05:13

@ Organikum:
I have been wondering about the exact amount of chlorine which is produced when a given amount of TCCA is reacted with an excess of HCl.
Today I tried to measure the gas volume, but the chlorine layer above the water was shrinking before my eyes as it dissolved to make sterile water.
But I also found out that the reaction indeed goes to completion, as I used 3,5mol of HCl on 1mol of TCCA and boiled to expell all chlorine, and on further addition of HCl, no more chlorine was produced on boiling.

In your "Practical production of chlorine" thread, you said that 6mol of Cl2 are produced from 2 mol of TCCA (I saw that you took in account that the commercial material is only 92% TCCA and therefore the "molar mass" of this product is 250 g/mol).

But in your downloadable calculations on TCCA, you stated that only 2mol of Cl2 are produced from 1mol of TCCA!

What is true now? Am I missing something?

[Edited on 19-5-2005 by garage chemist]

Preliminary workup

chloric1 - 5-6-2005 at 06:56

Recently purhcased two 8 oz giant 3 inch TCCA tablets and ground them up with my high tech granulator(ball ping hammer+ziplock freezer bag):D. Anywho, I took about 2 or 3 grams in a small beaker and added 50 ml of 31.45% HCl all at once. It definately made chlorine but the reagents bubbled to a surprising volume. Two thoughts I have; First dilute my muratic acid to exactly one half with water,Secondly slowly drip the HCl on the solid with extremely low heat. I wanted to see if it was possible to generate a gentle yet consistant flow of chlorine. Any creative thoughts welcome. Maybe a Kipp generator setup? :o

garage chemist - 5-6-2005 at 07:29

The concentration of the HCl is very important.
Someone found out that the reaction of TCCA with concentrated HCl does not go to completion. Reaction with dilute HCl gives much more chlorine and the reaction almost goes to completion.
The TCCA needs lots of water to be able to dissociate into HOCl and cyanuric acid.
The HCl must not be stronger than 15- 20%.

For making chlorine, I use a two- neck 500ml rbf with a pressure- equalized dropping funnel on the middle neck and a hose adapter on the side neck where the chlorine can be withdrawn.
All joints must be greased and secured with wire, otherwise they may get loose without you noticing until the chlorine gets into your nose.

The TCCA is coarsely granulated, not powdered (to prevent excessive foaming) and wetted with a little water.

When all HCl has been added, the mix must be heated to complete the reaction.

Thank you

chloric1 - 5-6-2005 at 09:11

I admit I have been out of the loop as far as practical hands on chemistry goes for awhile. I have a three month old daughter and am going through a career hiccup. But, Everything now will find its place;) I am going to stop being an armchair chemist and get going soon.

Buy the way, if I add water to my TCCA will it help to coalulate it again and give me a satisfactory Cl2 stream or is it too late for my powder? I do have a few larger particles but 80% is probably 50 or 60 mesh or finer. Ahh, I will play around a little more with it today and tomorrow.:cool:

By the way, a little off topic but still chlorine related. I have a pickel jar I will to use as an electrolytic cell/battery and I am sterilzing with sodium hypochlorite. This morning while doing the TCCA workup, I also added conc HCl to my very dilute hypochlorite(less tahn 1%) to destroy more organics and left it in the sun. I fugured that would be interesting as we are approaching record heat today. I wanted to see with how fast oxygen would be liberated from water in sunlight. What I found was that the gases evolved only when a foriegn object introduced. This is of absolutely of no value except in its academic realizations. But its cool:cool:

[Edited on 6/5/2005 by chloric1]

[Edited on 6/5/2005 by chloric1]

chloric1 - 5-6-2005 at 09:30

Here is a photo that demostrates the photodissocation of chlorine and water

Photodissociation.JPG - 310kB

<!-- bfesser_edit_tag -->[<a href="u2u.php?action=send&username=bfesser">bfesser</a>: reduced image size(s); attached]

[Edited on 9.8.13 by bfesser]

Chlorine gas/liquid color

Analytiker - 20-7-2005 at 19:08

I worked with liquified chlorine with a headspace quite a while back. I had a sight glass installed at the liquid/gas interface.

Green gas, red liquid. Very Christmas-like.

kickflip_333 - 2-8-2005 at 13:05

festive AND fun. :D

Mason_Grand_ANNdrews - 30-7-2006 at 09:15

I have any useful stuff to the chlorine source. If anyone interested, i will make a samll HTML to the subject.

Magpie - 18-5-2007 at 15:13

I've been wanting to try some chlorination reactions using UV light to generate free radicals for some time. I had read through this thread some time ago and decided to try the TCCA + HCl method for generating chlorine. My goal was to produce a steady stream of relatively pure, dry chlorine, with good control. Good control included being able to stop the reaction quickly should something go wrong.

My chlorine generator is a 1 liter suction flask. 18% HCL was to be dropped as required onto coarsely ground TCCA using a 250 mL separatory funnel. This assembly was sitting on a stirrer-hotplate for heating/agitation, if required.

For a drying train I set up two wash bottles in series partly filled with Rooto sulfuric acid (~96%). The inlet dip tube in each bottle was about 4cm below the acid surface. The glass tubing was 6mm OD and was connected together with 6mm ID Tygon (soft PVC) tubing. Following this is a 250mL Erlenmeyer flask as safety bottle in case of suckback. Finally, the Tygon tubing was connected to a 7mm OD gas dispersion tube having a fritted end to produce fine bubbles. This dispersion tube was placed in the chlorination vessel which is a large test tube. In the test tube I had placed some DCM.

Everything seemed tight so I turned on the 500w halogen lamp and the hood fan, and started dropping in HCl, just a few drops at a time. A yellow/green gas could be seen in the generator flask right away. But also right away a fatal flaw in my setup was apparent. Due to downstream resistance an adequate Cl2 pressure had to be built up and maintained in the generator flask. But when this pressure began to build it would vent Cl2 up through the stopcock of the dropping funnel. This can be easily understood when the backpressure is estimated: (2)(4cm)(1.85) = 14.8cm H2O back pressure just to overcome the sulfuric acid heads alone. Then add in the glass frit resistance and the DCM head of about 8cm and I come up with ~23cm H2O of backpressure. This was just about right to overcome the head of HCl in my dropping funnel. So now I see why addition funnels with pressure equalization are useful!

Actually, after reviewing this thread again, it seems like a another absorbant might even be advised. That is, shouldn't the first bottle contain an absorbant to take out the HCl gas that will surely be a contaminant in the Cl2?

It seems to me that the every gas delivery situation can be different. That is, some will require contaminant removal via solid absorbants (ie, CaCl2, mol sieves, etc), some will require absorption through a liquid (ie, H2SO4, NaOH, etc), and some may require some of each. In all cases some amount of driving pressure head will be required. This seems like an important and generic problem that may be encountered by the home chemist. But I find little or no information, guidelines, etc, in any of my books or on the internet.

This is a long story. But I wanted to get it out as an example of problems that can be encountered in gas generation, cleanup, and delivery. I think chlorine generation is a good example of typical problems that must be overome.

Below is a picture of rev1 of my chlorinator. Critique, suggestions, and discussion are welcome.

[Edited on by Magpie]

chlorininator.jpg - 57kB

The_Davster - 18-5-2007 at 15:56

I use acid filled syringes, with the tubing to the generating vessel secured well, to add the acid to the generator vessel. No issues with the pressure anoyances here.

Also just realized I have not worked with chlorine for around a year and a half...weird.

garage chemist - 18-5-2007 at 16:51

There is no need to remove HCl gas from the chlorine stream, as your reaction (chlorination of DCM) produces HCl anyway, like most organic chlorinations do.
Only removal of H2O vapor is necessary (but this is important). Conc. H2SO4 can be used, CaCl2 too.

CaCl2 may be a better choice for you, since it creates far less pressure drop than H2SO4.

While your method of chlorine production sounds good (apart from the issue with the dropping funnel), your method of chlorination has to be improved. For instance, a reflux condenser is absolutely vital if you are chlorinating anything. Those reactions either have to be done at the boiling point of the substance, or they produce a lot of heat. Cooling slows down the reaction too much, so refluxing is necessary.

Or did you just want to create a solution of chlorine in DCM? Then you would have to cool with ice.

I doubt that you will be able to get more than one chlorine attached to the DCM, since I have done chlorinations of things like red P in boiling chloroform and upon recycling of the solvent, no noticeable CCl4 was present.

One last thing: granulated TCCA does not react completely with HCl, since it gets coated with cyanuric acid. You either have to finely powder the TCCA, or you have to heat the gas generator to boiling after all HCl has been added.

[Edited on 19-5-2007 by garage chemist]

evil_lurker - 18-5-2007 at 17:38

Be careful with DCM... from what swim understands from reading patents, the reaction can go screwy fast resulting in a runaway explosion!

It must be moderated, preferably by putting in tubing or what not every .5 to 1cm or so to stop the radicals from going chain too fast.

I could be wrong, but better safe than sorry.

Intergalactic_Captain - 18-5-2007 at 19:23

So how much light is actually necessary to facilitate a chlorination? It seems like everyone here is using something around a 500watt lamp...I'm assuming this is simply because those who are performing these reactions have the fume hood / venting system to perform them in a lab setting. What about the sun, though? I've been doing some reading and it sounds like it's possible, though nobody here seems to have done it...unless I'm missing something.

I'd like to attempt the chlorination of toluene tomorow...I can build the setup (nigger-rigged, of course), but the strongest artificial light I can do is a 100watt (incandescent) bulb...It's supposed to be sunny with minimal cloud cover this weekend though, so I'm planning on using the sun unassisted...Should it fail, I could always rig up a light bulb, though I'd really rather not - my dad took my good extension cord while I was at school and I haven't seen it since.

So, if anyone's done up the toluene chlorination, did you do it in the sun? How well did it work? Also, just how bad is benzyl chloride - I know its a nasty lachrymator and all that jazz, but will sitting a few meters away be too close?

One last question - I'm planning on making benzaldehyde anyway...Is it absolutely necessary to distill off the toluene before the water/bicarb boil? Something about distilling tear gas without a respirator just doesn't sit too well with me. I'm thinking of rigging up a good magnetic stirrer and trying to react it as an emulsion...On second thought, this probably won't work too well if at all, but has anyone ever tried it?


EDIT - I just realized this is the chlorine thread...I know Organikum started it with the intent to use it for benzaldehyde, so I guess it's not too off topic...For relevance, though, I'm planning on dripping H2SO4 into a stoichiometric mix of KMnO4 and NaCl in water. I know there's a danger of Mn2O7, but I'm thinking it'll go to HCl->Cl2 in situ...It's more expensive than HCl, but easier for me to get ahold of the reagents for some odd reason.

[Edited on 5-18-07 by Intergalactic_Captain]

G.i.B. - 18-5-2007 at 23:48

I used the glass tube from an old mariuana growing light as a reaction vessel to chlorinate toluene, and did it in direct sunlight. It worked quite well. I only made about 20-30 ml just for the hell of it, and I am not sure about the purety of my end product. I used a 200 ml syringe to add the acid, so I had no problems with the pressure build-up.

Intergalactic_Captain - 19-5-2007 at 07:58

How bad was the tear-gas effect though? Benzyl chloride is supposed to be quite potent, though not very volatile...Would you have been able to pick up your tube and carry it a short distance without tearing up/etc.?

Magpie - 19-5-2007 at 09:31

Garage chemist says:
Quote:

There is no need to remove HCl gas from the chlorine stream, as your reaction (chlorination of DCM) produces HCl anyway, like most organic chlorinations do.


Um...of, course. That's what I get for not writing out the reaction equations.
Quote:

your method of chlorination has to be improved. For instance, a reflux condenser is absolutely vital if you are chlorinating anything.



Yes, I see your point. I wasn't sure how much heat chlorination of DCM would require. I know from experience that heat is necessary for aliphatic bromination. Also, DCM is so volatile it would be difficult or impossible to contain it without a condenser.

Quote:

I doubt that you will be able to get more than one chlorine attached to the DCM


I am sorry to hear this. My hope was to make CCl4. But CHCl3 would of value also.

Use of the syringe sounds like a clever way to get around the pressure problem. ;)

G.i.B. - 19-5-2007 at 10:41

Intergalactic_Captain. I am sure that you can carry it a short distance, Getting the benzyl chloride out of your reaction vessel and into your storage bottle, is another matter, you do not want to smell too much of it. Gasmask is advised ! And clean anything that comes in contact with it outside, you want nothing in your sink, you can smell it for weeks.

[Edited on 19-5-2007 by G.i.B.]

[Edited on 19-5-2007 by G.i.B.]

Magpie - 3-12-2007 at 20:24

I finally got everything together for another attempt at the free radical chlorination of DCM. This time I had benefit of a pressure equalizing dropping funnel, a gas mixing vessel, and a condenser. Also, instead of two H2SO4 water absorption bottles I used 1 bottle filled with BB size CaCl2 followed by 1 H2SO4 water absorption bottle.

I charged the 500 mL RBF with 55 g of coarsely ground TCCA and the dropping funnel with about 200 mL of 20% HCl. The dip tube in the H2SO4 was about 3cm submerged and I placed about 75 mL of DCM in the gas mixing vessel. The sparge tube in the DCM was also about 3cm submerged. Once chlorine generation was steadily underway I placed a 100w mercury vapor light next to the gas mixing vessel.

I wasn't sure what products I would get from this experiment (if any) but I was hoping for some nominal amount of CCl4. Most importantly, however, I wanted to verify that my new setup could generate dry chlorine in a controlled fashion.

At first things were disapointing. Chlorine generation was irregular and my safety relief (glass stopper in dropping funnel) kept popping out. Then once I had about half of the HCl added things settled down and Cl2 generation was a nice steady 1-2 bubbles/sec as seen in the H2SO4 water absorption bottle. I never did apply any heat to the RBF but did turn up the mag stirrer as the gas generation rate dropped off. After about an hour of steady chlorination I terminated the run without adding the rest of the HCl.

My hood was very useful here, steadily removing the Cl2 as it left the top of the condenser. I can't envision working with Cl2 without a hood. It doesn't take much in your breathing space to really make you uncomfortable. Outside I could barely smell it just under the vent. 25 feet away I couldn't smell it - of course it didn't hurt to have a windy day. :D

Here's a picture of my setup. I built a small ss wire cage for the outlet of the gas tube in the CaCl2 bottle before covering it with the CaCl2 BB's. I don't know if that was necessary but I tried to do everything I could to reduce backpressure in the system.

Any questions, comments or critique of this setup and/or the halogenation, is welcomed.

[Edited on by Magpie]

chlorination.jpg - 55kB

12AX7 - 4-12-2007 at 02:45

Hmm, you should have more shielding on that light, UV isn't great stuff.

Tim

Magpie - 4-12-2007 at 10:51

That's a good comment Tim. I was wondering just how much of a hazard that light is when unshielded. The lamp is unfrosted and is a relacement lamp for the shielded outside garage security lights.

evil_lurker - 4-12-2007 at 12:42

Most mercury vapor lamps have UV shielding in the bulb themselves... if the outer jacket gets broken, then yes there is a problem wiht UV radiation... this is a safety precaution by the manufacturers.

Likewise for most chemical reactions, UV radiation isn't an effective catalyst because pyrex type glassware absorbs quite a bit of UV rays.

Magpie - 4-12-2007 at 13:14

Yes, lurker, I see what you mean. I thought that somehow I was protected but had forgotten why. The lamp I have is a Caster type R. The packaging has this warning:

"WARNING: THIS LAMP CAN CAUSE SERIOUS SKIN BURN AND EYE INFLAMATION FROM SHORTWAVE ULTRAVIOLET RADIATION IF OUTER ENVELOPE OF THE LAMP IS BROKEN OR PUNCTURED. DO NOT USE WHERE PEOPLE WILL REMAIN FOR MORE THAN A FEW MINUTES UNLESS ADEQUATE SHIELDING OR OTHER SAFETY PRECAUTIONS ARE USED. LAMPS THAT WILL AUTOMATICALLY EXTINGUISH WHEN THE OUTER ENVELOPE IS BROKEN OR PUNCTURED ARE COMMERCIALLY AVAILABLE."

When we did free radical brominations in organic chemistry class we used an ordinary but unfrosted incandescent lamp, ~75w. We also just used ordinary test tubes (I assume borosilicate) IIRC. We also increased the temperature from 25C to 50C. Now I'm wondering if the added "hv" actually had any effect?

The_Davster - 4-12-2007 at 13:57

The few UV light sources I have used contained the bulb in quartz(?) and was intended for direct imersion into the reaction vessel.

However I think some UV must get through pyrex. I have prepared compounds that fluoresce under UV, and illuminating them through standard pyrex glass with a broad spectrum UV source did cause fluoresence.

12AX7 - 4-12-2007 at 14:02

As I recall, pyrex will transmit low UV, but not high UV, which only quartz will transmit. Soda-lime glass transmits neither to an appreciable degree.

Tim

Magpie - 4-12-2007 at 14:47

It seems like to make my uv light system effective and safe I will need to:

1. Break off the outer glass envelope on my mercury vapor lamp.

2. Use a quartz reaction vessel.

3. Place a safety shield around the lamp/vessel.

Perhaps if I just made sure that my hood sash was down when the lamp was on I would be protected. My sash is double-pane tempered glass.

not_important - 4-12-2007 at 16:34

Chlorine radical reactions are triggered by light in the 300 to 380 nm range, Pyrex and similar borosilicates transmit down 280 nm or so; there's no need for quartz unless the reaction requires that extra energy in the radicals formed. BTW, bromine absorbs in the range 360 to 510 nm.

Some mercury vapour lamps include shunts within the outer bulb that will open if run in air. Also, if the design of the bulb places the phosphors on the inner surface of the outer bulb; removing that bulb will given you lots of short wavr UV. But there were designs that had the phosphors coating the inner bulb or on a separate envelope, which would not give nearly the increase in UV when the outer bulb was removed.

You can take the chlorination all the way to CCl4, the 4th C-Cl bond is about 10% lower in energy than the others and steric factors come into play, but that just means taking additional time.

evil_lurker - 4-12-2007 at 17:57

Quote:
Originally posted by not_important
Chlorine radical reactions are triggered by light in the 300 to 380 nm range, Pyrex and similar borosilicates transmit down 280 nm or so; there's no need for quartz unless the reaction requires that extra energy in the radicals formed. BTW, bromine absorbs in the range 360 to 510 nm.

Some mercury vapour lamps include shunts within the outer bulb that will open if run in air. Also, if the design of the bulb places the phosphors on the inner surface of the outer bulb; removing that bulb will given you lots of short wavr UV. But there were designs that had the phosphors coating the inner bulb or on a separate envelope, which would not give nearly the increase in UV when the outer bulb was removed.

You can take the chlorination all the way to CCl4, the 4th C-Cl bond is about 10% lower in energy than the others and steric factors come into play, but that just means taking additional time.


Not_important is dead on.. deep blue light is what is needed to generate chlorine radicals, not UV light.

I noticed my yields went way up when chlorinating toluene when I switched over to a 20,000K metal halide vs 500w halogen.

Magpie - 4-12-2007 at 19:43

not_important your post is indeed good news. I'm glad I won't have to remove the outer envelope glass on my lamp, or find a quartz reaction vessel! And also about the wavelength needed for brominations - I guess my instructor knew what he was doing after all.

Below is a picture of my mercury vapor lamp. It is a 100w Caster H38 model, type R, made in China. It is also marked "02 06" which I assume is a date of manufacture. I couldn't find any specs on it but assume it is a "high pressure" lamp. I've been trying to find a spectrum for this type of lamp but haven't really been successful yet.

It does have a shunt w/resistor as you can see but I believe this is just part of the starting circuit using a 3rd electrode.

Evil_lurker I believe you are likely correct in that you have the best OTC lamp for chlorinations. I'm just trying to get a setup that will do the job at hopefully a lower cost. I'll just probably have to run my lamp longer than you would.

[Edited on by Magpie]

mercury vapor lamp.jpg - 75kB

Magpie - 6-12-2007 at 21:57

As described above I attempted to make CCl4 by free radical chlorination of CH2Cl2 (dichloromethane, or DCM). Tonight I did a fractional distillation of the 60mL of product from that chlorination. I used a Hempel reflux column packed with ss scrub pad. Before the distillation I washed the product once with 50mL of dilute aqueous Na2CO3 and then once with 50 mL of just water.

First condensate came over at 38.5C and was cloudy. After about 12 mL had collected the condensate had cleared up and T=39.5C. Then the remainder of the distillate was collected (40 mL) in the range T=39.5 to 40.3C. I took this to be all DCM. (DCM literature bp = 40.2C.)

The few mL left in the pot was a little sour smelling and had a slight yellow tint. It didn't have any chloroform or carbon tet smell at all.

So I have to say that my attempt at producing CCl4 was a total failure. I would be interested in hearing of the experiences of any others who have tried this.

MagicJigPipe - 11-12-2007 at 02:00

100w? I'm no expert in free radical chlorination or electromagnetic waves but... Why does it matter what kind of light source you use as long as it produces enough light/UV in the appropriate wavelengths. And saying that, 100w doesn't seem to be near enough energy even if it was extremely efficient. I mean instead of a 100w mercury vapor lamp, why not just go to home depot and buy a 1000w halogen lamp? It's cheap (except for the energy cost, which still isn't much worse than running a large microwave for a while) I mean, light is light no matter what the source.

Or is this all about maximum efficiency?

Also, c'mon you guys you're talking about UV like its gamma radiation. I would imagine the main risk is to your eyes. So, why not just cover it up, wear UVA/UVB sunglasses when you're working with it and don't let it shine on your skin for extended periods. I don't think that's too lax. Am I wrong?

not_important - 11-12-2007 at 02:35

The chlorination runs as a radical chain, giving thousands of C-Cl bond formations per photon. Plus there's no sense pumping in more energy than is convenient to remove, why overdo things?

If you have a lot of UV around it will be everywhere, not easy to avoid if you're working with it. I've had several chunks of skin cut out by doctors. Friends have corneas shaved. All effects from UV, why add to the total dose you get - for some parts of the body free radical damage is forever.

I've had success with this, but long ago. Starting with methane feed from the gas line and a slow electrolytic Cl2 feed; mercury vapour light. A 10 cm diameter borosilicate reaction tube 30 cm long at on focus of a buffed aluminium sheet elliptical reflector. Cooled collection flask for halogenated products. Ran it continuous for many days, mostly ignoring it, then removing several 100 ml of product for fractionation.

Liquid phase may shorten the radical chains, reducing yield per photon.

Magpie - 11-12-2007 at 10:45

Thank you not_important for providing me the benefit of your experience with chlorinating methane. In trying to explain why my attempt failed, the fact that it was in the liquid phase vs gas phase was one explanation I had thought about. In an old book this reaction was done in the gas phase. But in one sense I had this condition in the gas phase over my liquid phase, i.e., evaporating DCM and escaping chlorine. In the end I rationalized the lack of reaction as due to the low tendency of a zero degree alkyl group to participate in the free radical mechanism.

For MajicJigPipe: I had tried a 400w halogen light but it generated so much heat that that in itself became a problem for the low bp DCM. However, that attempt was quickly aborted for that and other unrelated reasons.

MagicJigPipe - 11-12-2007 at 17:43

Ok, then why is everyone so obsessed with maximum light input, talking about that bright sunlight is best? more photons = more free radicals = faster chlorination, right? Or is that an oversimplification?

And I could understand the heat issue (maybe an extremely effecient reflux condenser could compensate?). But say you could get a fluorescent light that emmitted nearly as much light as a 300w Hg vapor lamp. Wouldn't that be just as good if not better than a 100w Hg vapor?

I guess my whole point is. Why is Hg vapor so great and why should I spend the extra time and effort to get one?

P.S. Nice fume hood magpie. I'm still in the process of purchasing materials for mine.... Some day though.... :(

[Edited on 11-12-2007 by MagicJigPipe]

not_important - 11-12-2007 at 18:46

Quote:
Originally posted by MagicJigPipe
...
I guess my whole point is. Why is Hg vapor so great and why should I spend the extra time and effort to get one?
...


Because as I said earlier, Chlorine radical reactions are triggered by light in the 300 to 380 nm range, with some effectiveness on up to 430 nm or so. The phosphors in a fluorescent lamp convert most of the UV to wavelengths longer than 550 nm

http://commons.wikimedia.org/wiki/Image:Fluorescent_lighting...

mercury spectrum
http://www.lamptech.co.uk/Documents/M3%20Spectra.htm

Magpie - 11-12-2007 at 19:19

Here's some more spectra on the Hg vapor lamps:

http://crystec.com/lamp1.gif

and a transmission curve for borosilicate glass:

http://www.sinclairmfg.com/datasheets/borosilicatecurve.htm

Thanks for the note on the hood. ;)

I think use of a regular reflux setup would have been better than what I used, i.e., keeping everything at a boil would increase any reaction rate.

Today I designed an elliptical reflector for my next use of the Hg vapor lamp. Now I just need to find a cheap source of sheet aluminum.

[Edited on by Magpie]

MagicJigPipe - 11-12-2007 at 20:55

So, what about a tanning bed light since they are specifically designed to emit light in those wavelengths? And you wouldn't have to modify the bulb. Ok, it's time to stop questions and start experimenting.

Arggggg!!! I forgot about my broken reflux condenser and all I have now is a 300mm jacket West condenser. I don't think that will do it.... I wonder if I put a fractioning column between the pot and the condenser, if that would be sufficient. That's going to be really tall though... I'll figure something out.
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