Sciencemadness Discussion Board

Bromine Source and Synthesis

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Antwain - 15-10-2007 at 12:26

@Nicodem- Oh my god. How could I possibly have not known about that. For 5 years I have been miscalculating my KOH requirements, this explains a few things that have happened in the past. How dare they call it an analytical reagent. My most impure home-prepared reagents would never be put in my cupboard with 15% impurity. However it also says >85%, so I don't even know how much water is there. I would estimate that I added an additional 3g of KOH for a total of ~31.5g If not too much HCl came across then that is a good yield. Also, by the end of the distillation I would colourimetrically estimate that there was less than 1/20 atm partial pressure of bromine above the distillation flask @ ~80*C, and probably a lot less than that.

I may reduce the distillate to bromide and then feed that to some more BCDMH, but I will wait for that.

Also does anyone know why the solution is yellow? It went the kind of yellow you get in bleach at first, then to a pale yellow. This makes sense I guess, hypohalite and all, but since bromide and bromate are both white... why is it still yellow after being heated to 90*C for an hour?

Jdurg - 19-10-2007 at 17:03

From my own personal experience, once Bromine has been in something it NEVER goes away. It will permanently stain everything and attack almost everything as well. It's a little devil of an element and it took a great deal of glass and acrylic resin to keep it from leaching out. I've used tubing that appeared very clean and clear, heated it up in a flame and seen liquid bromine leach out from it. There's a good chance that you have very dilute amounts of Br2, or BrCl that is giving that very faint color.

Antwain - 20-10-2007 at 16:41

not in alkaline solution.

I should elaborate... When there was bromine dissolved in the solution and it was orange. Upon addition of more alkali and stirring this colour faded immediately. Further, the disproportination of OBr- takes place much more readily than OCl-, so I do not believe it is possible that after boiling it anything but KBr and KBrO3 should remain....

yet it is yellow :o and I don't know why. But I would bet my last dollar against it being bromine

[Edited on 21-10-2007 by Antwain]

woelen - 21-10-2007 at 04:39

I agree with you it is no bromine, but most likely, it must be some contaminant, which is discolored by the bromine.

I have a very similar experience, but in strong acid. I once made a solution of KBr, very concentrated, and added concentrated H2SO4, in order to obtain HBr (with sulfate and potassium contamination). This liquid was light yellow/brown, due to partial oxidation of bromide to bromine, and subsequent loss of SO2. The yellow/brown color remained.

Later, I took some of this liquid, and added Na2SO3 in order to make it colorless. But to my surprise, it did not become completely colorless, no matter how much Na2SO3 I added, and even when there was a strong pungent smell of SO2, it still remained light yellow. For me, this never was resolved. I don't have the liquid anymore, I used it up in all kinds of experiments (mainly for making drops of bromine with KBrO3).

ProfessorMolly - 20-11-2007 at 00:58

The method i have been using to get my bromine would be bubbling chlorine gas into a fully saturated sodium bromide solution, but the problem i have been having is the yeilds i get are awful, out of several runs and in total from 1lb NaBr i got about 60ml of Br2.

I believe that its because Br2 is soluble in water and when the gas tears it out of the solution it is still being dissolved into the water, and the actual dark bromine you get it just the amount left over that the water can't take any more of. So i have been looking for a different route to get this, and what sparked my interest was the H2O2, H2SO4, and NaBr route, now i understand that the purity of the H2O2 would be 35% but the questions that i had that i could not find answered here were, what ratios should be used (as in amounts of H2O2, H2SO4, and NaBr) and from what i read the NaBr is dissolved into the H2O2 and then H2SO4 is added somewhat slowly and Br2 is released, i jsut wanted to make sure of these things before wasting my time and/or money, also i had to ask, the final product , will it be a layer on the bottom of the solution that i can take off easily, or does it require distilling? its not that i don't have the equipment, its just id rather not, but if i must then i will,

If you guys could help me out with these questions id highly appreciate it, Thanks

roamingnome - 20-11-2007 at 13:16

Ive been trying to generate a useable amount of bromine by electrolysis because it seems so simple. Well the first trial mainly just produced bromine water.

So I had white gas and extracted the orange bromine water and the bromine associated with the white gas. But I came back 2 days later to see that the density of the bromine or something caused it to fall out of the top layer back into the water.

Now im building a custom cell that has a DCM water interface. The electrode will be as close to the dcm as possible and the bromine will fall of the electrode into the clutches of the DCM which can be tapped of the bottom. It will be slow but consistent process…

Anyways im not sure where all your bromine went but a DCM extraction of the water will surely not hurt.

StevenRS - 5-2-2008 at 11:21

I have made bromine before with HCl, H2O2, and NaBr, and have yet found a way to store it. It always leaks out of my glass stoppered flasks! Could I seal it with Teflon tape, maybe?

woelen - 5-2-2008 at 12:30

Storing bromine in a bottle with teflon tape around the rim and a thick teflon liner in the cap works exceptionally well for me. I simply don't smell any bromine near the cap of the bottle.

Below follows a picture of what I mean:



This is the bottle with its cap.



The white cap is a modified cap. These caps contain some foam material, when purchased. I replaced that with a thick teflon liner (more than 1 mm thickness). The teflon tape really must be there. Without the teflon tape, I can smell the bromine, even when the cap is screwed on the bottle very firmly.

Btw, I store this bromine in a plastic bag, with 1 kilo of Na2S2O3.5H2O outside the bag, and another bag around these two chemicals. This Na2S2O3.5H2O works as a safety filler, which absorbs bromine in a controlled manner, just in case the bottle breaks.

[Edited on 5-2-08 by woelen]

StevenRS - 5-2-2008 at 17:05

Hmm... I wonder, can I melt Teflon plumbers tape without it decomposing? That could open up a lot of possibilities....

12AX7 - 5-2-2008 at 19:15

Not really. It can be formed more easily when hot, but it doesn't melt (gooify) like thermoplastics. Consider your nonstick frying pan (the Teflon one, not the cast iron one*), which is pretty solid up to perhaps 400-500°F, then gives off toxic fumes any hotter.

*Any cook who knows anything (which, okay, need not necessarily include any chemistry geek -- but I can always hope!), knows that well oiled and seasoned cast iron pans work just as well as any Teflon-coated pan. Technical progress? Yup -- they figured it out centuries ago! :)

Tim

MagicJigPipe - 5-2-2008 at 22:29

I would like to add that I have attempted to melt PTFE tape before. It "melts" very slightly before decomposing into a brown tinged mass at around 330*C. Definitely not good for molding at all.

I use the technique that woelen mentioned above for pretty much all of my volatile chemicals. It works very well. During the winter time, keeping Br2 outside (in the garage or shop) seems to help, also.

StevenRS - 6-2-2008 at 18:43

I tried to coat a glass stopper in it to achieve a better fit, and in never seemed to work, and if it did, it was single use only. I recently tried melting (or heating) the teflon tape covered stopper, then inserting it into the flask, and it seemed to work a lot better.

skippy - 10-2-2008 at 20:37

I did a little bromine synthesis tonight. I don't have any fancy ground glassware so I made my apparatus from two erlenmeyer flasks connected by tubes and stoppers I cast and lathed out of lead. The stoppers were covered with teflon tape. I put a small cooling jacket on but I didn't end up using it as I did the reaction outside where it is very cold right now ;-). I don't have yields yet, I think I may have lost a bit through the stoppers, and my peroxide was old too so I'm not sure what strength it was, so maybe there was still bromide unoxidized. Or maybe I over did it thinking it was weakened (can peroxide oxidize bromine?).

The bromide was from a liquid spa solution, the sulfuric acid was "liquid lightning" and the the hydrogen peroxide was hydroponic 30%. I slowly added all the peroxide at bottle concentration and didn't have any significant distillate until I applied heat to the large reaction flask. I'll put up some pictures of the product soon - its an nice little bit of bromine.

Nothing too revolutionary about the reaction, but I thought maybe someone would find my construction method interesting. The castings were made in wood molds lined with foil. The stopper molds were made by drilling a hole with a forstner bit into wood. The stoppers were cast then drilled and a rod threaded into them so they could be turned on a lathe until they would fit the flasks. The pipe was made by casting lead around a steel mandrel in a piece of wood I cut a trough in with a router. The steel mandrel was tightly wrapped with a strip of aluminum foil to make removal easy.

The pipes were rasped untill they would fit into holes drilled in the stoppers. The pipes were flared into the stopper holes and then soldered with a tin-lead solder for mechanical strength. No tin solder was visible on the inside of tube stopper interface. I may try this reaction again soon with some fresh peroxide.



[Edited on by skippy]

[Edited on by skippy]

[Edited on by skippy]

bromine.jpg - 76kB

a_bab - 26-2-2008 at 13:35

Bromine for everyone: (found this interesting to share): http://www.perekopbromine.com/main.php?lang=eng&page=pro...

Now these are ALOT of bottles...

Also an idea about safely storing it (as somebody said before, with Na thiosulphate in a metal container instead of perlite)

LSD25 - 1-3-2008 at 22:10

Quote:
Originally posted by BromicAcid
To avoid the cholorine gas all together I usually dissolve my sodium bromide in water and add calcium hypochlorite then once the two are mixed I add a bit of hydrochloric acid. The mix instantly turns red from bromine liberation and if the concentration is high enough you get some rewarding drops of dark red in the bottom of the reaction mixture. Then the bromine can be distilled off of course.


So, with the latest incarnation of the constant battle to keep bromine away from those who want to isolate it being to mix sodium bromide with hypochlorite, all that is necessary is to add HCl to liberate the bromine?

(the mix as sold has a considerable excess of calcium hypochlorite (s), around 750-800g/kg IIRC while the NaBr content is around 150g/Kg).a

You know, I was just about to post a question asking if anyone had any bright ideas on how to separate the two, kinda lucky I read the WHOLE thread, ain't it?

I was actually thinking of feeding it something to oxidise (like aluminium), same way you get iodine from KI when you add too much OCl after HCl... It really is the same thing, aint it?

Just finished distilling...

Bohrium - 8-1-2009 at 16:53

Today I obtained 180g (60ml, an 80% yield) of bromine by distilling 400g of DBDMH with 200ml of water while slowly adding 250ml of concentrated HCl.

Here is the apparatus:



After about 75 ml of acid had been added:



And here is a portion of the final product:



Edit:

DBDMH = 1,3-Dibromo-5,5-dimethylhydantoin



[Edited on 9-1-2009 by Bohrium]

Lambda-Eyde - 9-1-2009 at 03:07

That's a wonderful setup. I really envy you. :P

Is there any specific reason to why you chose to use DBDMH? Is it readily available? It sounds like a rather exotic compound to me. Did I miss anything? :)

[Edited on 9-1-2009 by Lambda-Eyde]

woelen - 9-1-2009 at 03:30

DBDMH is not exotic at all. It is a swimming pool brominating agent and in countries where bromination of swimming pools is allowed this is as common as TCCA for chlorinating swimming pools.

Unfortunately, in NL the bromination of swimming pools is forbidden en for this reason this compound is not available here. Here we only have TCCA, Na-DCCA and calcium hypochlorite, no bromine-containing chemicals.

Did you assess the purity of the bromine? Water and remains of acid easily can be removed from the bromine, but I still think it is rather impure, due to the presence of elementary chlorine. In the process with DBDMH besides formation of bromine, there also is formation of chlorine, which reacts with bromine to BrCl, which dissolves in Br2 exceptionally well and cannot be removed easily by distillation, unlike the water and remains of acid.

[Edited on 9-1-09 by woelen]

Bohrium - 9-1-2009 at 05:11

Quote:
Originally posted by woelen
DBDMH is not exotic at all. It is a swimming pool brominating agent and in countries where bromination of swimming pools is allowed this is as common as TCCA for chlorinating swimming pools.


Exactly.

I was able to obtain 5lbs of it for $14.
Quote:

Did you assess the purity of the bromine?


I was unable to do so yesterday, but I'll measure its density later today. Will the BrCl evaporate out if the bromine is left to stand around 5 degrees C? I don't believe much could have condensed, because my condenser coolant was at around 10 degrees.

[Edited on 9-1-2009 by Bohrium]

woelen - 9-1-2009 at 05:35

I would not simply leave the bromine in an open bottle at 5 C. You will loose way too much bromine as well. BrCl dissolves in bromine exceptionally well and is not easily removed by non-chemicals means.

There is a good way to get rid of the BrCl, but this step requires a second distillation of all the bromine. The BrCl is destroyed chemically instead of physically removed.

Prepare a solution of NaBr or KBr. In your situation with appr. 200 grams of bromine I would take 25 ml of water, dissolve 2 gram of KBr or NaBr in the water and add this solution to your bromine. Shake well and then let the bromine settle at the bottom. If you do not have NaBr or KBr, then you can sarifice some of your current bromine to make a solution of a bromide (see end of post).

Separate the bromine from the aqueous layer as good as possible. If you have a separating funnel then that is most conveniently used, otherwise pipette away the bromine from under the aqueous layer.
Next, take 10 ml of concentrated sulphuric acid and add the bromine to this. Shake well again. Any water, left over in the bromine, now is bound to the sulphuric acid.

Distill this liquid and what goes over is very pure and dry bromine.

The treatment with KBr/NaBr destroys any BrCl: BrCl + Br(-) --> Br2 + Cl(-)
The Cl(-) goes in the aqueous layer.

The treatment in turn with conc. H2SO4 is done in order to easily separate from the water. Without the acid, you cannot get your Br2 free of water in a single distill. At 50 C or so, also quite some water vapor comes over as well. But if you separate as much as possible of the water using a funnel, then the rest can be tightly bound by the sulphuric acid. Your H2SO4/Br2 mix now can be separated very easily, no water will come over below 150 C, while the Br2 goes over at 60 C or so. You can distill over all bromine and you stop when only sulphuric acid remains left. Do not continue the distillation to also get the bromine, dissolved in the H2SO4. That last little bit must be regarded as lost. It will only be a small amount, I think less than 1 gram if you use 10 ml of concentrated acid.

----------------------------------

In case you have no bromide: Take 10 ml of household ammonia (5% NH3) with NO additives. Mix with 15 ml of water. This makes 25 ml of 2% ammonia, To this, slowly add 1 ml of bromine while swirling (which is excess of bromine, but you mix with your bromine anyway, so that is OK).
This liquid gives a very clean quantitative reaction with bromine:

8NH3 + 3Br2 --> 6NH4Br + N2

You really have no side reactions. Any BrCl in your bromine immediately reacts to chloride and bromine, so your solution will contain ammonium ions and bromide ions and a small amount of chloride ions. This can be usded as your solution of NaBr or KBr.

[Edited on 9-1-09 by woelen]

Jor - 9-1-2009 at 08:23

I'm most impressed by your fume hood! Did you buy it or build it yourself? If yes, that's a very neat sash and baffles!

pHzero - 29-5-2009 at 07:53

Sorry to revive an old topic (I don't know how people feel about that here, you get told off for it on some forums) but wouldnt an easy way to make a high yield of bromine be to bubble HBr (produced by KBr+H2SO4) through a solution of BCDMH or DBDMH?
You can produce bromine by oxidising HBr, or by reducing the HOBr produced by BCDMH, so I'm surprised that noone's suggested oxidising/reducing one with the other?
Surely HBr+HOBr--->H2O+Br2?

I happen to have some KBr, and I could get BCDMH quite easily.

woelen - 29-5-2009 at 11:32

Yes, this can be done, but it is quite a hassle. It is much simpler to simply dissolve the KBr, add sufficient acid and then add the BCDMH (or any other suitable oxidizer). Bromine drops out and can easily be distilled (see above).

Bubbling HBr through a liquid containing BCDMH has no advantage over the procedure, given above. You also need to distill bromine.

pHzero - 29-5-2009 at 11:54

Quote: Originally posted by woelen  
Yes, this can be done, but it is quite a hassle. It is much simpler to simply dissolve the KBr, add sufficient acid and then add the BCDMH (or any other suitable oxidizer). Bromine drops out and can easily be distilled (see above).

Bubbling HBr through a liquid containing BCDMH has no advantage over the procedure, given above. You also need to distill bromine.


Ah right - I did consider that, but then I thought maybe the K+ or SO4 2- might somehow interfere with the reaction, eg the dissociation of the BCDMH into HOBr.

But still, thats basically what I meant, using one to oxidise/reduce the other, rather than introducing external oxidisers like H2O2

len1 - 7-11-2009 at 02:56

A bottle of Br2 containing some BrCl in solution, will essentially rid itself of the BrCl if let to vent to the atmosphere. BrCl equilibrium is quickly attained, so the task is to rid of the BrCl while minimizing loss and in fact gaining Br2 in the bottle due to dissociation of the former.

BrCl boils at 5C, taking this and assuming the vent is an efficient reflux at this temperature, (ie time scale of reflux much greater than 2BrCl -> Br2 + Cl2 equib constant ) the bromine partial pressure is about 0.07. The equib constant is about 0.14

pCl2 pBr2/pBrCl^2 = 0.14

so

pCl2/(0.93-pCl2)^2 = 2 -> pCl2 = 0.45, pBrCl = 0.48

Say there is initially 1 mol Br2 and 1 mol BrCl in the bottle (obviously instantaneosly since this mixture is not in equib).

The 1 mol of atomic chlorine will be lost as 1/3 mole of Cl2 and 1/3 mol of BrCl. In addition about 1/20 mole Br2 will be lost from the bottle. 2/3 mol BrCl would have been converted to 1/3 mol Br2, so the bottle now contains

1 + 1/3 - 1/20 about 1.28mol Br2.

We have gained 0.28 mol.

Experimental test gave me almost spot on this result.

[Edited on 7-11-2009 by len1]

len1 - 7-11-2009 at 14:01

As a corollary to this

Venting BrCl to the atmosphere in an efficient 0-5C reflux will convert max 56% to Br2, and lose 44% to air.

Thus if one uses BCDMH as the bromine source with HCl, essentially all bromine will come out as BrCl and in a simple distillation setup aimed at liquifying the bromine little yield will be obtained. Setting up one efficient reflux on the distillation products will yield about 50-55% of the bromine. Using a two stage reflux setup, the second stage must be colder than the first so as to trap most BrCl while the first stage is active and pass the Cl2, followed by vent to air, will net about 75% of the bromine.

len1 - 8-11-2009 at 01:11

Further on this topic, I got interested in the maximum yield of bromine available from BCDMH. The question is not without merit, as due to the large market for this chemical it has the cheapest cost per bromine atom compared to say bromide or bromate salts. Surprisingly cost a factor not just for the amateur, professionals are expected to economise too these days.

Coproportionation of BCDMH and say NaBr using HCl has been discussed up this thread, but as the bromine comes 3:1 from the bromide, its essentially comes at the same cost as the bromide. Using BCDMH as per my chlorine generator and with reflux collumn to decompose the BrCl gave a 50% yield with one collumn, still cheaper than bromide, but can one do better?

Yes, Na2S2O5 also discussed up this thread is a large market reducing agent and therefore cheap, it can outperform the HCl method, but only with some fine tuning.

A balanced equation between the Na2S2O5 and BCDMH was written up this thread, but real life is more complicated.

The following reductions of the (BCDMH==RBrCl) are all possible

a) 2RBrCl -> 2RClH + Br2
b) 2RBrCl -> 2RBrH + Cl2
c) 2RBrCl -> 2RH2 + Br2 + Cl2
d) Br2 -> 2Br-
e) Cl2 -> 2Cl-

Which are important is all a matter of the equilibrium constant AND the kinetics, the latter is hard to come by for BCDMH, which lies outside the class of classical chemicals. Could be measured - and this I did just enough to get a workable method. The formula up the thread corresponds to c) and e) occuring together, I find a) accounts for ~70 of the reduction processes, with d coming second. This is somewhat surprising as chlorine is the more easily reduced species, and must be a matter of kineics rather than thermodynamics. It is also clear that d) can not occur unless either b) and c) do not occur, or e) has completed.

So the main reaction is

4RBrCl + Na2S2O5 + 3H2O -> 2NaHSO4 + 2Br2 + 4RClH

1/4 of the moles of metabisulphite are required, which is very efficient.

Take 120 gms (0.5 mole) BCDMH place in a 1 liter three neck flask with thermometer, distilation head, droping funnel. The distilation head leads to a west condenser at 30-45 degrees fed with 0-10C water from a recirculating pump. Add 40-50mls H2O and mix by shaking to a thick slurry. Add 10ml conc HCl. Place 80ml metabisulphite solution made by dissolving 160gms Na2S2O5 in water to 500ml in funnel. Adjust drip rate so the drops are just too fast to count (about 3 per second). The entire flask will soon turn a deep red and the contents will boil, with temperature rising to 83-87C. The addition rate is such that bromine essentially boils off as soon as its formed. After all reducer has been added wait for boiling to subside and turn on heater raising temperature to about 97-99C until gas in flask loses red colour (becomes orange due to BrCl). Yield 30gms bromine phase, + about 1 gm in water phase. Note that insufficient reducer added for anything but a) to occur substantially.

Interestingly small amount of HCl raises yield somewhat.

What happened to the 25% Br that didnt come out? Most is still bound to the hydantoin although more than stoichiometric amount of metabisulphite has been added needed to reduce it. If one keeps adding bisulphite and heating brown-yellow fumes of BrCl appear, a tell tale sign that the remaining a) now occurs together with b) (that is c), and to recover the remaining bromine takes a bit more work.

[Edited on 8-11-2009 by len1]

len1 - 8-11-2009 at 13:57

And Ive got more to say on this.

After ALL evolution of bromine had ceased, as indicated by lightening of vapour in flask I added 100ml conc HCl with the expectation of a BrCl flux due to 25% bromine remaining according to previous experiment. I got just a Cl2 flux! No Br2 condensed in my second stage BrCl reflux reactor. This eliminates possibility of Br- in solution AND Br remaining in the hydentoin complex. So I revise my earlier claim of a 75% yield, its closer to 100%, difference being due to BCDMH containing DCDMH, which happens to be cheaper than the former.

So conclusion is that metabisulphite releases Cl from above compound with difficulty, it mostly reduces bromine.

And so we have a method. Reduce BCDMH with bisulphite as per above to release the bromine, then reduce with HCl to release the chlorine. And so you have used it to its outmost.

[Edited on 8-11-2009 by len1]

Polverone - 8-11-2009 at 16:38

I take it from the course of events above that your BCDMH source did not disclose its substantial content of DCDMH. Was this just an over the counter spa/pool disinfection product? I have seen products based on this compound that claim various percentages of the active ingredient on the label, but have not seen a disclosure of the balance of composition.

len1 - 8-11-2009 at 22:10

The BCDMH I used is intended as a dissinfectant for the consumer market, and on a per bromine basis appeared to be a factor of 2 cheaper than bromide from the cheapest chemical supplier, which got me interested. The label states BCDMH as the active ingredient, but more than that, it claims 650gm/kg available bromine and 280gm/kg chlorine. Thats pretty much in the ratio of the atomic masses of these elements, which means no DCDMH is claimed in the product. The available bit is a funny trick which doubles the apparent mass of active ingredient, for chlorine you will get this amount by reacting with HCl since it supplies the missing half, but for Br this is nonsensical, since the consumer is not supplied with sodium bromide or hydrobromic acid. BUT If you were to react the BCDMH with HBr you would get about 650gms/kg bromine. Hence they are claiming almost pure BCDMH, certainly not at the 75% level. Further on their web site they claim 89-97% active ingredient. So the product is outside spec. Its a bit like the sultanas in the cornflakes.

[Edited on 9-11-2009 by len1]

1281371269 - 6-3-2010 at 07:35

I've been away for a while due to having a lot of schoolwork but it's easing up a bit so I've been back in my lab and have been doing some basic organic experiments.

Ozonelabs kindly sold me some cyclohexanol and 2-ethyl-hexanol, quantities of which I will be dehydrating to form the corresponding alkenes. To test if the dehydrations have been successful, some Br2 water will be useful.

As I only need about 1.7g max. of actual Br2 for 50ml bromine water solution I think the easiest way to go about it would be:

  1. Saturate 50ml water with KBr
  2. Make a chlorine generator using HCl and KMnO4
  3. Bubble chlorine through KBr sol. until no more will react (tested by holding damn litmus above reaction vessel)
  4. Distill the Br / KCl sol. until all the water and bromine has gone over


Can anyone see any problems with this method or suggest an easier way around?

Thanks for the help!

[Edited on 6-3-2010 by Mossydie]

woelen - 6-3-2010 at 08:06

I have tried that reaction and I must say it is a failure. I bubbled the chlorine through a fairly concentrated solution of NaBr, but the chlorine does not react with the bromide sufficiently fast. So, at least 90% of the chlorine bubbles through the liquid and the rest reacts with the bromide. This was not a success at all.

At a certain point, things became even worse. The air-mix above the liquid became bright yellow/orange, indicating that the chlorine, which bubbles through the bromide takes with it quite some already formed bromine, which gives its color to the air above the liquid.

I am considering making a web-page about this failed experiment (I made nice pictures of cool glasswork to show a nice synthesis, but the outcome was otherwise, but I still have the pics).

1281371269 - 6-3-2010 at 08:35

The reason I wanted to go by this route is that as toxic gases go, chlorine bothers me fairly little, and the reaction would of course be done outside on a windyish day. Then I could take the Br / KCl sol into the lab for distillation. This in my mind makes it one of the safest routes as there's no potential for large amounts of Br2 fumes. Poor efficiency isn't really a concern, I've got 500g of the KBr

I have tried the reaction and it seemed to go well, however I could smell Cl2 from the start and the solution anyway had under a gram of KBr in it. It could be improved though, by warming the KBr sol, by diffusing the chlorine via a funnel, or even by adding the chlorine to a sealed flask which was warmed / shaken, until it reacted, then adding more and repeating. With those improvements would it be satisfactory, or is there a better way around?

entropy51 - 6-3-2010 at 08:46

There is a prep somewhere on the internet for bromine water. It's basically dissolving KBr in diluted bleach and acidifying with a little HCl.
Quote:
I have tried that reaction and I must say it is a failure. I bubbled the chlorine through a fairly concentrated solution of NaBr, but the chlorine does not react with the bromide sufficiently fast. So, at least 90% of the chlorine bubbles through the liquid and the rest reacts with the bromide. This was not a success at all.
Woelen, I've made lots of Br2 that way. I can't imagine what went wrong, but it does work. Maybe it just doesn't work in the Netherlands.:D

1281371269 - 6-3-2010 at 08:54

Entropy - thanks, I've found the synthesis. There were two on the page, the first one went something: Get liquid bromine and add it to distilled water. Gee thanks.

But the second one was as you said NaBr + NaClO + HCl.

What is the actual reaction there?

entropy51 - 6-3-2010 at 09:00

Without balancing the equation, acidifying the hypochlorite makes Cl2, which just displaces the Br2 from NaBr. Same as bubbling Cl2 through NaBr as far as I know. I can't imagine why it didn't work for Woelen. He's forgotten more chemistry than I'll ever know.

I've used HCl + KMnO4 to make Cl2 many times, without incident. But some have said there is danger of explosion, for what it's worth.

[Edited on 6-3-2010 by entropy51]

woelen - 7-3-2010 at 12:27

Quote: Originally posted by entropy51  
Woelen, I've made lots of Br2 that way. I can't imagine what went wrong, but it does work. Maybe it just doesn't work in the Netherlands.:D

How fast did you bubble the chlorine through the solution and what concentration of bromide did you have in your experiment?

The experiment I did was just for educational purposes, I wanted to demonstrate the classical schoolbook reaction for making bromine from a bromide and chlorine gas. I made the chlorine gas from Ca(OCl)2 and 15% HCl.

Of course I did get bromine, but the reaction simply was not as good as I expected. Too much of the chlorine simply bubbled through the solution of NaBr while this solution was quite a concentrated one. I simply did not want to spoil so much chlorine in this reaction and after a few minutes I stopped. I used 10 grams of NaBr and all of this I have converted to Br2 by adding KBrO3 (which in the past I have made from electrolysis of KBr).

Just for the fun, I post the setup I used for this experiment.

The total setup is shown here:


At the right is the chlorine generator, 15% HCl is dripped slowly and regularly on solid Ca(OCl)2. The gas is lead through some glass tubing, through atest tube. The test tube contains a rubber stopper with two holes and a second glass tube leads the chlorine gas into a concentrated solution of NaBr in a second test tube. The final glass tube is a pasteur pipette with a very fine tip, allowing the bubbles of chlorine to be small.

A safety tube is inserted as well. If for whatever reason the system has a tendency to suck back, then I could turn open the little handle on the table, allowing air to be sucked in.


Some detail pictures of the chlorine generator, the safety setup and bubbler:







Finally I want to show how the bubbles of chlorine went through the NaBr. There was a nice constant stream of chlorine gas, with a few bubbles per second. But you can see in the picture that the bubbles are not absorbed by the liquid very much.



After a few minutes I had to quit the experiment, even with the window above the setup fully open, the stench of the chlorine became unbearable :o


[Edited on 7-3-10 by woelen]

entropy51 - 7-3-2010 at 13:02

One difference in my method was the use of KBr instead of NaBr. According to my notes I used 25 gm KBr dissolved in 100 mL H2O. The Cl2 was generated using KMnO4 + HCl, and the Cl2 was bubbled into the KBr using just a simple glass tube. My recollection is also that the Cl2 delivery tube was near the top of the KBr solution instead of the bottom as your pictures show. The flow of gas was relatively slow, but after a few minutes the KBr has turned dark red and soon Br2 begins to separate on the bottom of the flask. Some Br2 vapor definitely did escape the flask, but in the hood it was impossible to say how much. The KBr flask was an ugly red fuming mess contrasted to your picture.

I separated the Br2 in a sep funnel and then distilled quite a bit more from the aqueous solution. I did not keep an accurate record of the yields, but rather just repeated the prep a few times until I judged I had enough Br2 to last me for a while.

I wonder if you used enough Ca(OCl)2 to generate enough Cl2, but I'm sure you would have calculated it out to react with all the NaBr. There is a rather large volume of air to displace in your setup however, and that would require some additional Cl2. Because of the air being displaced I don't think the bubbles would have been absorbed initially. I don't think that the Cl2 is sufficiently absorbed in the NaBr solution to conduct this experiment without a hood or discharging the escaping Cl2 into caustic absorber. Perhaps if you had been able to add more Cl2 to the NaBr you would have seen more Br2?

[Edited on 7-3-2010 by entropy51]

woelen - 8-3-2010 at 01:18

I used a large excess of Ca(OCl)2 (maybe two times as much as needed). I have taken into account losses of Cl2 (I assumed loss of at least half a liter of gas) and the bubbles you see in the picture are at a time when the flask is really full of chlorine (you see the green color of the gas in the flask). Of course, the first bubbles of gas are not absorbed, but that is what I expected. But when the apparatus is full of Cl2 then still the absorption of gas is slow. I did not get any blob of Br2 with the Cl2. I had the chlorine in the bottom, because I thought it would be better because the bubbles of gas are in the liquid for a longer time. I really find it strange that the solution absorbs the chlorine so slowly.

But because you report success, I will retry this experiment soon. I have plenty of Ca(OCl)2 and HCl and I just ordered 1 kg of NaBr. So, I can repeat this experiment soon. I will make the production of Cl2 much slower than I did in the last experiment, I start thinking more and more that the high speed of chlorine production is the problem. Next time, I'll also use a delivery tube with a wide mouth instead of the thin needle-like tube I used in this experiment. In that way, the Cl2 might be absorbed by the liquid without many bubbles going through the liquid. This is the only thing I can think of what might have gone wrong. I had pure Cl2 gas, I had excess of Cl2 gas and the apparatus itself looks quite good for doing this kind of experiments. I have done this before in making KIO4 and Na3H2IO6 with Cl2-gas bubbled through alkaline iodide solution or iodate solution.

http://woelen.homescience.net/science/chem/exps/KIO4_synth/i...
http://woelen.homescience.net/science/chem/exps/Na2H3IO6/ind...

Both experiments are very similar, with chlorine bubbled through a liquid. In these experiments, this was very succesful and hardly any chlorine could be smelled.

len1 - 8-3-2010 at 07:38

A saturated solution of Br- will disolve about 4gms Br2/100ml before any is precipitated out. Also Cl2 bubbling through Br2 will volatalize it since the equilibrium favours BrCl - so the tube must not be at the bottom.

And finally this is an experiment for the fume hood - the chlorine bubbles do not dissolve fully that would mean you would see no gas reaching the liquid surface. Well over 50% of chlorine is lost as Cl2 and BrCl (this is the brown colour, Br3- is dark red). So while this method is demonstative it is inefficient.

woelen - 8-3-2010 at 12:27

Yes, I can confirm that it is not as efficient as I hoped it would be. But I did not expect it to perform so bad as I have experienced in my experiment. I'll retry it with the Cl2-delivery near the surface of the liquid and much slower generation of Cl2-gas.

My goal with this experiment indeed is not to have a perfect bromine generator. I have much better methods for that, based on making KBrO3 and acidifying a mix of KBrO3/NaBr. I just wanted to make a cool demonstration of a common schoolbook reaction which can be carried out by many people who do not have the equipment and the knowledge to make their own KBrO3.

len1 - 8-3-2010 at 14:06

You can have your reaction ttube, same as your safety - with a gas outlet. This way you can lead the chlorine outside (to your neighbours :) ) and the experiment can run longer (why not use silicone tubing instead of glass tubes). Reducing concentration of Br- will decrease time before Br2 appears as it has drastic effect on its solubility - i suspect it will not much affect reaction rate as that is determined be Cl2 mixing. I dont believe slower gas generation is what you want - halve the Cl2 rate, double the time you have to wait before Br2 appears. The amount of Cl2 released is about same in both cases

[Edited on 8-3-2010 by len1]

woelen - 11-3-2010 at 04:16

I do think that Cl2 flow rate must be reduced, because that allows the gas to be absorbed by the solution. Actually no bubbles should be formed at all, the gas simply must be absorbed. If the gas is produced too fast, then the gas forms bubbles and these may make it to the surface and get lost in the air (cough cough !!).

The reason why I use glass tubing simply is that I have a lot of it and I have bent them in all kinds of shapes. I like glass better than other materials, it is clean, is not corroded. The only disadvantage is that it easily breaks, but I work carefully and till now I never broke one of these tubes.

I expect my new stock of NaBr to arrive this week (I ordered 1 kg) and then I'll try again with a somewhat larger quantity.

entropy51 - 11-3-2010 at 13:18

In an attempt to see why the preparation of Br2 didn't seem to work for Woelen, I repeated this on a scale, and with conditions, intended to resemble those he reported.

This time I used NaBr instead of KBr. I dissolved 2.5 g of NaBr in 10 mL of water and poured it into a large test tube.

This time I made the Cl2 by dripping dilute HCl onto Ca(OCl)2 instead of the KMnO4 I used previously. I did not use a trap like Woelen did, since there is no suck back. The Cl2 delivery tube was immersed about half way into the NaBr solution, about 2 cm. The experiment was conducted in the hood so that I could let the Cl2 bubble thru the NaBr solution fairly rapidly. The Cl2 was generated from 5 g of Ca(OCl)2 and about 15 mL of 32% HCl diluted with an equal amount of water.

In less than a minute after starting the flow of Cl2, which bubbles up through the NaBr solution and doesn't appear to dissolve, the solution turned dark orange and orange fumes of Br2 filled the tube. After 3 or 4 minutes longer a large drop of liquid Br2 collected in the bottom of the test tube and the reaction appeared to be over.

I feel certain that Woelen will be able to make Br2 using this method when his NaBr arrives and he tries it again.

len1 - 11-3-2010 at 18:56

Quote: Originally posted by woelen  
I do think that Cl2 flow rate must be reduced, because that allows the gas to be absorbed by the solution. Actually no bubbles should be formed at all, the gas simply must be absorbed. If the gas is produced too fast, then the gas forms bubbles and these may make it to the surface and get lost in the air (cough cough !!).



Thats what I wrote - "the chlorine bubbles do not dissolve fully as that would mean you would see no gas reaching the liquid surface".

But the rate of bubbling has little to do with the efficiency of their dissolution provided they do not subsequently coalesce - thats because once a bubble is formed its rate of passage through the liquid is independent of the time interval with which the bubble behind it is formed. The latter is governed by the size of the initial bubble, its path length in solution, and reaction kinetics. The size of the bubbles is governed by the apparature, and to first order (provided the bubbling is not too fast) is independent of the rate of gas flow.

If you want to go without using bubbling - just a slow two-phase reaction, it will be more efficient but will take much longer. The rate is proprtional to the area of contact of the two phases, so with a 5mm tube its 100 times slower than if you used an inverted glass funnel 5cm diameter, whose tip is connected to a chlorine gasometer - if you wanted to do it in a beaker. Magnetic stirring underneath the funnel will also speed the rate up. But I prefer bubbling

[Edited on 12-3-2010 by len1]

woelen - 11-3-2010 at 23:53

Yes, I agree, once a bubble is moving upwards, its rate of absorption does not depend on how fast it was formed, hence my idea to prevent formation of bubbles at all. I also realize that I need to have a glass tube with a wide open end, a small inverted funnel might do the job, but maybe it also can be done (slowly) with a glass tube of 1 cm diameter (I also have quite a few of those). It's just a matter of experimenting and tweaking.


[Edited on 12-3-10 by woelen]

1281371269 - 13-3-2010 at 11:58

I did the experiment today and got fairly decent results, certainly made some bromine. I have about 200ml of more than saturated bromine solution which I will distill at a later date.



This was the best setup I could pull together with the kit that I have. It worked well, though of course most of the chlorine simply bubbled through the KBr. Thus the yeild was low, but on the other hand I only used about 25mL conc HCl and no more than 50g of KMnO4 so it was hardly wasteful.



A close up of the bromine formed.

As a side note: Could anyone tell me how to clean my glassware? The bromine stains I'll attack with a stong base, but what about the mess left in the RBF? How about pirahna acid?

entropy51 - 13-3-2010 at 12:48

Bromine does not stain glassware, at least not clean glassware. It might color traces of grease already on the glass.

Once you get of the mess left in the Cl2 generator flask, HCl should remove the Mn stains left on the glass. Just use a few mL and swirl it around. A little mild scouring powder and a brush should do the rest.

Do you have liquid Br2 in the bottom of the beaker? It looks like it. You should be able to distill a lot of Br2 out of the solution. Use magnetic stirring or boiling chips - it bumps badly. You can probably again gas the KBr solution with more Cl2 after you distill the Br2 out of it.

I wonder why the bromine water is so cloudy? I've seen it turn cloudy briefly but it clears up with continued gassing with Cl2.

For Cl2 generation I use about 6 mL of HCl per gram of KMnO4, which I believe is the ratio in Vogel. 1 gm of KMnO4 should produce about 1.1 gm of Cl2, so you can calculate how much you should use. It is much easier to clean up the generator if you use Ca(OCl)2 instead of KMnO4.

[Edited on 13-3-2010 by entropy51]

1281371269 - 13-3-2010 at 12:56

It looks a lot clearer in the contain I have it in at the moment. There is probably a fair amount of impurity, for example, I suspect that some KMnO4 managed to find it's way in as I poured it into the rbf.

woelen - 15-3-2010 at 01:36

Mossydie, this is exactly what I intend to do with my setup. My new order of NaBr arrived last weekend so I can do the experiment soon. You have a fairly large contact area between the Cl2 and the NaBr-solution. Did it stull bubble or is the gas simply absorbed by the liquid?

I won't use KMnO4 for making Cl2, but Ca(OCl)2 or TCCA because these are OTC and much cheaper. I'll use the apparatus I have shown further up in this thread, but I'll use a wide glass tube for delivery of the Cl2 in the NaBr-solution.

1281371269 - 15-3-2010 at 10:35

It still bubbled unfortunately, though I don't have a dropping funnel so had no control over the rate at which Cl2 was produced.

madprossor - 16-3-2010 at 01:20

hi, i have used the bromate method to produce Br2 without distillation.

80% molar yield is easy with this method-

dissolve 87.5g KBrO3 in 1650 ml h2o with heating. add 315g KBr and continue heating and stirring until it dissolves. remove heat and attach reflux condensors circulating ice water. continue stirring, turn on laminar flow exhaust, drip in 162ml strong h2so4 at a rate that maintains reflux.

after the acid is added, the reaction is cooled for about 30 minutes to the point where Br2 precipitates but not to the point where sulfate cyrstals fall from solution (this lowers yields). it must be KHSO4, right? Br2 is then tapped off. the highest yield obtained was 205.3g, more commonly 200g.

i put together a solubility chart in order to help understand the applicability of NaBr to a similar procedure. it looks like there is no downside to using NaBr because NaHSO4 is more soluble in cold water than the potassium salt. perhaps yields would be even higher.

grams soluble per 100g h2o



Solubility of Bromine is increased in the presence of its salts and in HBr

thanks to Klute and Woelen for first bringing KBrO3 to my attention. and len1, thank you also for your innovations.

[Edited on 16-3-2010 by madprossor]

[Edited on 16-3-2010 by madprossor]

[Edited on 16-3-2010 by madprossor]

woelen - 21-3-2010 at 08:27

I once again tried the method with Cl2 and NaBr. Now I used an inverted funnel-like construction on top of the surface of a solution of NaBr. Pictures will follow soon. But again, I am not happy with this method, it simply sucks. From a theoretical point of view it is very nice, and I can also imagine that from an industrial point of view it also is a good reaction, but for making Br2 as a home chemist it is only a last resort for me. Taking the effort of making KBrO3 and then reacting that with Kbr (or NaBr) and acid is better than using the Cl2-method.

I now did have a blob of Br2 at the bottom of the receiving flask, but it took a LONG time. I now was faced with a different problem. When I had the inverted full of Cl2-gas, then slowly the liquid was drawn into the funnel and Br2 was formed at the surface. Initially things went quite well, but when the concentration of Br2 increases, then red vapor escapes from the liquid. This vapor is more heavy than Cl2-gas and it slows down the absorption of Cl2 considerably. At a certain point the reaction almost comes to a halt and absorption is very slow.

In order to revive the reaction I had to increase the speed of chlorine generation, such that big bubbles escape from the inverted funnel, driving away the bromine vapor and replacing this by fresh chlorine. In this process of course a lot of Br2 and Cl2 is lost into the air. And that is what irks me, I simply have to accept too many losses.

entropy51 - 21-3-2010 at 08:53

I made about 100 mL of Br2 over the course of a few weeks, making a run when I had an hour or so to spare. I used a Cl2 delivery tube dipping into the KBr solution instead of the inverted funnel. If the Cl2 generates the Br2 at the surface of the NaBr of course the Br2 will be carried away with the Cl2 rather than being trapped in the liquid. I didn't see the problems you're having, but I admit that a bit of Br2 is lost as vapor. This is easily avoided by attaching a long reflux condenser to the absorbtion flask containing the NaBr (KBr in my case), but I didn't find it necessary since all the chemicals are readily available and cheap.

Looking back through this thread, I see that others have easily made Br2 by this method:
Quote: Originally posted by neutrino  
I decided to make some bromine by the Cl<sub<2</sub> + NaBr method. The chlorine was generated by adding bleach to NaHSO<sub>4</sub>. This was then led into a flask containing the dissolved bromide.


Quote: Originally posted by neutrino  

This process is interesting: while bubbling in the chlorine at a medium-fast pace with a crappy bubbler (end of a Pasteur pipette), large bubbles kept breaking the surface of the shallow solution, yet little bromine vapor seemed to escape, except when my chlorine source generated too much chlorine in one big blow.

The KBr + H2SO4 + H2O2 method posted by Bromic Acid, Magpie and others appears to be a superior method and I plan to use that method the next time I need to make Br2.

[Edited on 21-3-2010 by entropy51]

[Edited on 21-3-2010 by entropy51]

woelen - 21-3-2010 at 13:11

Oh yes, the Cl2-bubbling method can be used to make Br2, but it is cumbersome and I don't like the losses of bromine. Bromides are not really expensive, but still too expensive to accept large losses.

The method with H2O2 indeed is a nice method, especially if you have access to H2O2 of better than 3% concentration. I also have used that method in the past. I like this method because no chlorine is involved and no BrCl can be formed.

Another method which I also like quite well is the addition of solid Ca(ClO)2 to a solution of NaBr in 10% HCl. This method, however, requires precise weighing of the reagents. A slight excess of NaBr must be used. Adding too much Ca(OCl)2 will lead to formation of BrCl.

My favorite method, however is the use of KBrO3, added to a solution of NaBr or KBr in 10...20% H2SO4. With that, there is no involved chlorine and the production of bromine really is fast and quantitative, even better than with H2O2. The biggest disadvantage of course is that you first need to make KBrO3 by means of electrolysis of a solution of KBr.

Taoiseach - 22-3-2010 at 00:08

There's certainly no advantage over Painkilla's method of producing the chlorine in situ from HCl and TCCS. Except for the possible formation of interhalogen compounds this method seems to be very convenient. There's not a whole lot of strong oxidizers available OTC these days, and TCCS is one of them.

Btw TCCS is capable of oxidizing iodide to iodine (a highly exothermic reaction btw and a nice way to make I2). Why doesn't this work with bromine as well?

6KI => 6K(+) + 6I(-)
6I(-) => 3I2 + 6e(-)
C3N3O3Cl3 + 6e(-) => C3N3O3(3-) + 3Cl(-)
3(K+) + C3N3O3(3-) => K3C3N3O3 (tertiary potassium salt of cyanuric acid)
3(K+) + 3Cl(-) => 3KCl

In theory it should be as simple as mixing bromide with TCCS plus a little water and destilling.

[Edited on 22-3-2010 by Taoiseach]

Lambda-Eyde - 22-5-2010 at 15:08

Yesterday I prepared bromine according to Magpie's method. I used the same amounts of reagents, except for the H<sub>2</sub>O<sub>2</sub> which was only available as the 30 % solution.
Attached to the distilling apparatus was a scrubber containing a NaOH solution. It turned into a more brownish color during the synthesis. The smell of bromine could not be detected during the synthesis, only when I dismantled the apparatus. This was done with the fume hood running and the glass was promptly placed into a big bucket of NaOH solution.

Pictures:

The setup:




Quite an impressive sight...




The receiver, cooled in ice water and attached to the scrubber. I should have used more ice.




The product: About 25 mL of nasty bromine. :D What's left to do is to dry it with concentrated sulfuric acid and bottle it. I have already labeled a 100 mL Schott DURAN reagent bottle with a red, PTFE lined melamine cap specifially for this. :D



[Edited on 23-5-2010 by Lambda-Eyde]

Magpie - 23-5-2010 at 08:39

Thanks for sharing your pictures. It's nice to see that someone is actually making use of this procedure. Do you have any plans for the bromine?

Lambda-Eyde - 23-5-2010 at 10:02

Thanks for publishing the procedure. :) I wish I had better pictures, but the procedure was done at school and I don't usually bring a proper camera to school.

I have no plans for it yet other than to seal up two samples in ampules for my element collection, the other one I'll give away as a gift to an element collector I know.
I plan to explore organic synthesis in the following year and I think it will be put to use then.

Lambda-Eyde - 26-5-2010 at 14:44



Here's a picture of me separating the bromine from the sulfuric acid. Note that the stopper is still in. :D I noticed that a few seconds later... When will I learn?

Also, I got my first bromine burn when I tried to make the ampules. I didn't wear the clumsy gloves because they didn't allow me the dexterity I needed to handle the ampules. Also, I thought the ampule was sealed... Turned out it wasn't! :o Luckily I had a liter of thiosulfate solution a few meters away, allowing me to wash my hand very quickly. My finger is only slightly brown, and it didn't hurt at all. I guess I was lucky. The accident gave me quite some respect for bromine!

Magpie - 26-5-2010 at 14:56

You were lucky. As, I think woelen said, bromine is a little hellion. It attacks stainless steel and will attack your tissues with a vengance. I keep mine in the freezer so that it has no significant vapor pressure. Otherwise it would probably be leaking out past the Teflon lined cap.

It is fascinating to look at though. ;)

un0me2 - 15-7-2010 at 03:25

For those looking to store the Bromine they generate, pass it into an aqueous solution of SO2 (aka sulfurous acid - I've made it by dripping HCl onto metabisulfite, the SO2 gas dissolves in water VERY easily, endothermic too), which will reduce it to HBr (and oxidise the H2SO3 to H2SO4), pass the gas into a solution of alkali to get the KBr/NaBr salt(s).

For those in Oz, it is a nice way to get both pure H2SO4 & Bromide salts. Very few jobs require 'actual bromine' so instead of trying to store the mongrel stuff, turn it into a nice white powder that can be weighed out and stored easily.

slinky - 23-10-2010 at 23:20

This thread is awesome. Thanks everyone for all the cool pics and tutorials ^_____________^

verode - 14-11-2010 at 13:57

you must take care of bromine fumes
if you put into NaOH solution you may get NaBr and NaBrO3
then add ice and HCl
you shaking bromine with water and decanting

ElementalRebel - 18-1-2011 at 14:03

Ive been lurking about this forum for quite a while and decided it was time to finally sign up and say thanks to the people in this and a couple other threads on bromine synthesis and storage. Thanks to this site and those posting such useful info on it, I was able to produce two nice samples of bromine this past weekend using the sulfuric acid/potassium permanganate/Sodium Bromide method and my distillation apparatus.

There is one thing Im wondering about. I chose to keep my larger sample under concentrated sulfuric acid, as recommended on this site, but upon adding the acid the fume production increased and the acid became very cloudy and red. I let it settle and there is bromine under the acid, but the acid has stayed cloudy red and fumes continue to be produced.

The second sample I put under distilled water and all went well. Water turned slightly red, but fumes have just about ceased entirely and theres been no problems I could notice.

Do you think my acid had some sort of contaminant in it or wasnt concentrated enough? Heres a pic of the two flasks, just so you can see the difference. I can get a better pic need be.

woelen - 19-1-2011 at 02:29

It is strange that your acid becomes cloudy. Somewhere there must be a strong impurity. Normally, H2SO4 will float on bromine and any water, still present in the bromine is absorbed by the acid. After some time there will be a clear red layer of acid above an almost black layer of bromine and the amount of vapor, released from the acid is noticeable, but not excessive. It is a nice method of storing bromine. You keep it dry and when you open the bottle you do not have copious amounts of bromine vapor escaping from the bottle. But... this assumes clean H2SO4 of 96% concentration.

ElementalRebel - 20-1-2011 at 10:54

Ill have to work a little on my acid to check what may have gone wrong. There was something else I noticed that I found odd. It appears almost like theres some solid particles under the acid with the bromine.

As I tilt the flask and the bromine moves about, there also seems to be small, jagged shaped "chunks" that prob avg 2-3mm in width as well. They dont move like the droplets of bromine and also appear to hold their shape. The bromine was very clean before adding the acid, so now im trying to think of what those "chunks" could be.

Ill try to get a decent picture after work. Thanks for the reply and any future ones.

plante1999 - 18-7-2011 at 10:38

Today , I fond something very interressing , I do not know the balanced equation.

I was attemping to make bromine with NaHSO4 + Ca(OCl)2 + NaBr.

I take 5ml of 30% NaBr (from brominating solution for pool) and i ad 10g of NaHSO4 ( from pool Ph minus (an exess)

but my father was needing help so i go help him , 2 hour later i go continu my synthesis , but when i look at my capped vial , there was a layer of bromine , and pressur was formed in the vial. Take note that almost no heat is produce.

I am sure there is no too mush impurity.


[Edited on 18-7-2011 by plante1999]

plante1999 - 22-7-2011 at 06:37

I have somes new for my prosses , It seem only work at very low ph close to 2. also the gas formed is acidic in presence of water , SO2? , i tested purity of my reagent and there was a goo purity.

Bringing back an old topic

ScienceHideout - 2-6-2015 at 19:03

Quote: Originally posted by Engager  
To produce large ammounts of bromine with ease i use following reaction:

5KBr + 3H2SO4(aq.) + KBrO3 => 3Br2 + 3K2SO4 + 3H2O

Procedure is straight forward: 63g KBr is dissolved in 300 ml of water, 18 ml of 95% H2SO4 is added with stirring (car battery acid can be used to disslolve KBr, taken in such ammount that resulting H2SO4 concentration is about 10%). Solution is transfered to 500 ml flask and 17.5g of potassium bromate is added by small portions with stirring. Stirring is continued until large drop of liquid bromine is formed in the bottom. Bromine is separated on separating funnel and dried with concentrated H2SO4. Yield is almost quantative, but some ammount of bromine remains dissolved in water (it is not large and depends from temperature).
[Edited on 12-10-2007 by Engager]


I know that the last post in this topic was in 2011, but I would like to further discuss one of my finds that you may find interesting, but I didn't wish to open a new topic.

I was fascinated with this procedure above. I was inclined to try it myself- this sort of 'distillation-less' bromine production method. It is brilliant.

So, first thing I did was check the stoiciometry. See what was in excess, and what wasn't. I was really just planning a quick and messy isolation, so I wasn't interested in sigfigs, I did rough calculations and I rounded some numbers. Come to find out, the potassium bromide amount was .5 mol, sulfuric acid I converted to grams using density from wikipedia and found it was .3 mol, potassium bromate was .1 mol. It was about a tenth of the stoiciometric quantities, meaning a little less than 7.5 mL of bromine was expected. Not bad, so I decided to try it.

In lieu of KBr, I used NaBr, adjusting the mass to that which is respective to its molar mass.

Reaction was fun and surprisingly simple. I made an ice bath with water to get the maximum yield, stuck a flask in with DI water, and weighed my compounds. I put in 18.2 mL sulfuric acid, and 17.665 g KBrO3. These are VERY close to what the procedure calls for according to the procedure and stoiciometry, and with later calculations I found that the variations in the number of moles of each of these compounds would be too small to throw off the reaction.

Here is the crazy part. I added EXCESS NaBr. The sodium bromide I buy is from the pool store and put in little pouches, so I decided to just use a whole pouch instead of what it calls for, that way I don't need to seal up the extra bromide. The total mass of NaBr I put in was 66.352 grams, whereas the correct amount should've been 51.445 g. Again, though, just a quick and messy procedure, I wasn't too concerned with any of this.

So clearly, the sulfuric acid and bromate should've been the limiting reagents.

So here I go, I dissolve the bromide, add the sulfuric acid, mix, add the bromate bit by bit, mix, and put it in a sep funnel. My first thought: "holy crap, that is a lot of bromine!"

In the end, I found I had 10.4 mL of bromine- a 140% yield! How crazy! Not to mention, the top layer from the sep funnel was orange, and out of curiosity, I spun a small amount in a centrifuge and found that liquid bromine could come from that, too, meaning there is A LOT more bromine produced than there should've been according to this reaction. Perhaps there is water dissolved in it which raises the volume? I doubt it because of another calculation that I made.

Interestingly enough, if I calculate the yield of bromine in respect to the NaBr, and assume the other reagents were in excess, I end up with a quantity that is very close to what I got. The problem is, the other reagents WEREN'T in excess.

So, is it just me, or is something else going on in this reaction, or is it all just an odd coincidence? :o

gdflp - 3-6-2015 at 07:35

You're calculations are off. Based on your quantities, potassium bromate is the limiting reagent. 17.665g KBrO<sub>3</sub> corresponds to 105.8mmol. Based on the stoichiometry of the reaction, 3 moles of Br<sub>2</sub> are produced for every mole of KBrO<sub>3</sub>, thus 317.4mmol of Br<sub>2</sub> is the theoretical yield, which corresponds to 50.72g of bromine, or about 16.3ml.

byko3y - 3-6-2015 at 10:29

KBrO3 is relatively expensive for bromine synthesis. You can use KClO3 for the same purpose.

woelen - 3-6-2015 at 11:55

The use of KClO3 is not recommended. Your bromine will become impure, it will contain BrCl and maybe other contaminants like ClO2. KClO3 is not a clean oxidizer. Many people think that it works like

ClO3(-) + 6H(+) + 6e --> Cl(-) + 3H2O,

but this is certainly not its only mode of operation. It also forms Cl2 and ClO2.

If you use excess bromide, then BrCl can be destroyed, but ClO2 will linger.

A very nice method of making Br2 is electrolysing a solution of NaBr or KBr, such that appr. 1/6 of all Br(-) is converted to BrO3(-). No need to isolate the bromate, just electrolyse and then add acid. I wrote a webpage about that. Using this method you do not need to buy the relatively expensive bromate.

http://woelen.homescience.net/science/chem/exps/OTC_bromine/...

Making the required chromate or dichromate is even easier than I write in the webpage. Take 250 mg or so of any chromium(III) salt (e.g. chrome alum, which is easy to obtain) and dissolve in a few ml of water. Heat the liquid gently, so that it becomes quite warm but not boiling hot (50 to 60 C is OK). Add bleach to the warm solution dropwise while swirling. Add just enough bleach to make the liquid clear and yellow. This liquid can be used as such, no need to purify or concentrate. Add it to the solution of NaBr or KBr and use the resulting solution for electrolysis.

Another reason for not using chlorate is that this compound will soon be prohibited in the EU. Possession of chlorates is illegal, so using this will become a problem next year.

ScienceHideout - 3-6-2015 at 13:01

Quote: Originally posted by gdflp  
You're calculations are off. Based on your quantities, potassium bromate is the limiting reagent. 17.665g KBrO<sub>3</sub> corresponds to 105.8mmol. Based on the stoichiometry of the reaction, 3 moles of Br<sub>2</sub> are produced for every mole of KBrO<sub>3</sub>, thus 317.4mmol of Br<sub>2</sub> is the theoretical yield, which corresponds to 50.72g of bromine, or about 16.3ml.


I am such an idiot, I used the mass for Br instead of Br2. Thus, the amount of bromine, in grams, is supposed to be twice as much as what I got since it is 160x0.3 rather than 80x0.3.

That is where my calculation was thrown off!

Thanks for making me double check this! Next time I do some stoiciometry I will be double sure to check the formulas.

It is sort of funny- I always tutor chemistry students and I can't tell you how many times I've said "Mass of oxygen is 32, not 16!" I guess I better practice what I preach! :P

mr.crow - 3-6-2015 at 14:21

Great post woelen

Its worth noting again that Br2 will react with bromide ions to make tribromide which is soluble. Any excess bromide will reduce your recovery!

byko3y - 3-6-2015 at 17:19

woelen, to produce ClO2 in a reducing medium you need to draw it from the mixture as fast as it is forming:
HClO3 + HCl => HClO2 + HOCl
ClO2 + H2O => HClO3 + HClO2
And chlorine reacts with bromide, leading to Br2, unless you draw it from the reaction.
Anyway, I think ClO3- is overkill, and hypochlorite can do the job. Then you just leave your bromine to evaporate with distillation column, so BrCl just flies away.
mr.crow, AFAIK there's not NaBr3/KBr3 compound.

blargish - 3-6-2015 at 18:50

Quote: Originally posted by byko3y  
AFAIK there's not NaBr3/KBr3 compound.


I'm pretty sure the tribromide ion does exist in a manner similar to the triiodide ion, where iodine can dissolve into NaI/KI solutions due to the formation of [I3]<sup>-</sup>. (you cannot necessarily isolate NaI3/KI3)

byko3y - 3-6-2015 at 21:44

Even if this compound exists - I don't think it will somehow prevent bromine distillation.

j_sum1 - 3-6-2015 at 22:47

Thanks for resurrecting this thread. I can see that I will need to take the time to read the whole thing. A couple of comments before I do...

1. The thread emerged just moments after I had completed an extraction of Br2 from NaBr using acidified permanganate as an oxidant. I wanted some bromine dissolved in a non-polar solvent for my students to investigate saturation in lipids. I know I could have used the bromine water available, but, you know oil and water don't mix that well.

2. Making a decent amount of Br2 for my element collection has been on my short list for a while. I have put it off until after my next glassware purchase. I may as well have the stuff contained. I have almost decided on a displacement of NaBr using Cl2 gas. I know that there is some chance of getting the interhalide compound BrCl but I understand that that is fairly minimal. (source -- youtube clip entitled something along the lines of "my favourite way to produce bromine".) Based on woelen's post I might however revisit the electrolysis idea. Mrhomescientist has a nice clip on that and I have read woelen's linked page in the past.

3. I thought this thread was going to be about sources of bromine compounds for use as a starting reagent. I have some reagent grade NaBr but wondered about other sources. In my section of the world it is not possible to find bromine salts. I understand that they are quite readily available in some locations as pool chemicals. The most common bromine chemical here is 1-Bromo-3-chloro-5,5-dimethylhydantoin which has the problem of chlorine being present also. I am on the lookout for suitable otc sources if I can get it. I did come across a mixture of 85% NaDCCA and 15% NaBr. I wondered about the logistics of adding HCl directly to this mix so that the Cl2 gas is produced in situ with the NaBr with the probable evolution of a mixture of Cl2, ClBr and Br2. I need to do some more research before I try this.

[end of ramble]

byko3y - 3-6-2015 at 23:37

Quote: Originally posted by j_sum1  
I have some reagent grade NaBr but wondered about other sources. In my section of the world it is not possible to find bromine salts.

I'm sure you can purchase few kilos of it on ebay without any problems.
Looking at this I realize the day will come when some people will not be able to obtain NaCl.

woelen - 4-6-2015 at 00:16

@byko3y: Tribromide certainly exists, and its formation is responsible for reducing the yield dramatically, if excess bromide is present. Bromine dissolves very well in a concentrated solution of a bromide salt. By means of distillation, the tribromide can be broken apart and bromine escapes from the liquid, but if you want a distillationless synthesis (just separating the bromine from the aqueous layer), then the formation of tribromide is highly undesirable.
Even when distilling, the formation of tribromide is undesirable. The bromine is more tightly bound to the solution and you have to make the solution hotter. This makes separation from the water less efficient. I know this from personal experience. The distilled bromine then is very wet, a lot of water also distills over together with the bromine, due to the higher temperature needed to drive off all bromine.

ClO2 also is a nasty beast. On the one hand, it is very reactive when present as gas and it can explode. On the other hand, it reacts very sluggishly in aqueous solution. It is remarkable how slowly it reacts with many reductors. In the production of bromine (where excess oxidizer is needed to avoid formation of tribromide ion) the ClO2 can linger for a long time and it can dissolve in the Br2. You do not have to fear explosion of ClO2 in such cases, but it can be a highly undesirable contaminant of the bromine. For this reason I do not like chlorate as oxidizer in aqueous solution. If you want to see the effect yourself, just try it. Take some chlorate salt and add a reducing acid mix (e.g. HCl which has chloride as reductor, or dilute H2SO4 with a little sulfite added which has SO2 as reductor) to the chlorate. You'll see formation of bright yellow ClO2 and you'll also see that it lingers for hours at least without any reaction.

Hypochlorite can be used as oxidizer for formation of bromine, but it has the disadvantage, that using excess of this oxidizer leads to considerable losses in aqueous solutions due to easy hydrolysis. Just try this to see it yourself: Take some NaBr, dissolve in some dilute H2SO4 and then add conc. bleach. You get a point at which there is a lot of bromine. If you go on adding bleach, then quickly the Br2 redissolves again and you get a piss-yellow solution, which contains hydrolysis products of BrCl (HBrO3, HCl). This issue exists for chlorate used as oxidizer as well. It does not exist for use of bromate as oxidizer.

NaBr3 and KBr3 cannot be isolated from solution. The cesium salt, however, can be isolated. I made some CsBr3 myself and keep a small sample of it in my lab, just as a curiosity. It is a red/orange compound.

byko3y - 4-6-2015 at 06:26

Everybody knows iodine forms triiodide, and bromine probaby forms it too, and the problem is that triiodide is just as well water soluble as the iodide salt. But... what is the atual equilibrium of the trihalide formation? Nobody seems to know.
http://pubs.acs.org/doi/abs/10.1021/ja02202a004
"The solubility of bromine in aqueous solutions of sodium bromide" - James M. Bell , Melville L. Buckley
J. Am. Chem. Soc., 1912, 34 (1), pp 14–15; DOI: 10.1021/ja02202a004
Solubility of bromine in a distilled water is 34 g/L, which is 0.21 mole/L. At 96 g/L concentration of NaBr (0.932 mole/L) solubility of bromine in water becomes 1.24 mole/L (397 g/L) (original work mentiones 2.48 gram-ATOM/liter).
As you can see, basically one mole of bromide converts into one mole of tribromide at low concentrations, while at higher concentration, something happens that makes possible for 1.5 more moles of bromine to be soluble in the same solution.
But anyway, the most important thing for us is to know that if we have not high enough concentration of bromide, then it stochiometrically turns all the bromide into tribromide. But small amount of bromide won't prevent extractive separation.
Basically, you can use excess of oxidizer, which will result in BrCl and Cl2 formation, but I want to remind you that we have 5:1 proportion of Br:Cl.
The main source of ClO2 is from chlorite (NaClO2) in case of HCl oxidation, ph of the solubtion is another question, because at different ph the products may be different.
I'd say the major pathways that give stochiometric conversion of chlorite are
2 NaClO2 + 2 HCl + NaOCl → 2 ClO2 + 3 NaCl + H2O
and gas-solid reaction
2 NaClO2 + Cl2 → 2 ClO2 + 2 NaCl
Actually, there's a crazy intermediate http://www.sciencedirect.com/science/article/pii/13811169950... having formula BrClO2.
And I want to emphasize that bromine is more readily oxidized than chlorine.
Chemistry of halides high oxidation states is very complex, so I'd say it's all about experimental data. The question is: will the ClO2 be formed at all in this reaction?
I hope I will try soon to experiment with NaOCl+NaBr, but for now I'm pretty much sure the problem is about ph of the solution.

morganbw - 4-6-2015 at 07:11

Quote: Originally posted by byko3y  
Everybody knows iodine forms triiodide, and bromine probaby forms it too, and the problem is that triiodide is just as well water soluble as the iodide salt. But... what is the atual equilibrium of the trihalide formation? Nobody seems to know.
http://pubs.acs.org/doi/abs/10.1021/ja02202a004
"The solubility of bromine in aqueous solutions of sodium bromide" - James M. Bell , Melville L. Buckley
J. Am. Chem. Soc., 1912, 34 (1), pp 14–15; DOI: 10.1021/ja02202a004
Solubility of bromine in a distilled water is 34 g/L, which is 0.21 mole/L. At 96 g/L concentration of NaBr (0.932 mole/L) solubility of bromine in water becomes 1.24 mole/L (397 g/L) (original work mentiones 2.48 gram-ATOM/liter).
As you can see, basically one mole of bromide converts into one mole of tribromide at low concentrations, while at higher concentration, something happens that makes possible for 1.5 more moles of bromine to be soluble in the same solution.
But anyway, the most important thing for us is to know that if we have not high enough concentration of bromide, then it stochiometrically turns all the bromide into tribromide. But small amount of bromide won't prevent extractive separation.
Basically, you can use excess of oxidizer, which will result in BrCl and Cl2 formation, but I want to remind you that we have 5:1 proportion of Br:Cl.
The main source of ClO2 is from chlorite (NaClO2) in case of HCl oxidation, ph of the solubtion is another question, because at different ph the products may be different.
I'd say the major pathways that give stochiometric conversion of chlorite are
2 NaClO2 + 2 HCl + NaOCl → 2 ClO2 + 3 NaCl + H2O
and gas-solid reaction
2 NaClO2 + Cl2 → 2 ClO2 + 2 NaCl
Actually, there's a crazy intermediate http://www.sciencedirect.com/science/article/pii/13811169950... having formula BrClO2.
And I want to emphasize that bromine is more readily oxidized than chlorine.
Chemistry of halides high oxidation states is very complex, so I'd say it's all about experimental data. The question is: will the ClO2 be formed at all in this reaction?
I hope I will try soon to experiment with NaOCl+NaBr, but for now I'm pretty much sure the problem is about ph of the solution.


You seem to have a better understanding of the chemistry than I do, however, unless you have more need to understand the chemistry, versus actually getting the product????

There are some tried/proven methods here???

If you life, is as a research chemist then, by all means, do your experiments, I bow to you.

If you just need the element sir, then you do not need to reinvent the wheel.

For me, the years are passing and time is becoming a factor, I lean on others experience.

[Edited on 4-6-2015 by morganbw]

ScienceHideout - 4-6-2015 at 07:23

Quote: Originally posted by morganbw  


You seem to have a better understanding of the chemistry than I do, however, unless you have more need to understand the chemistry, versus actually getting the product????

There are some tried/proven methods here???

If you life, is as a research chemist then, by all means, do your experiments, I bow to you.

If you just need the element sir, then you do not need to reinvent the wheel.

For me, the years are passing and time is becoming a factor, I lean on others experience.

[Edited on 4-6-2015 by morganbw]


morganbw,

If you just do chemistry for the hell of it, perhaps you should try a different hobby. Byko3y is doing chemistry the RIGHT way. Why would anyone want to do something that they don't understand?

Everyone should aspire to know what they do and do what they know, and some are better at it than others. If you are jealous at the incredible understanding byko3y has with this topic, good. I am too, and it comes with years and years of discipline and study. However, I don't think that snotty remarks will better YOUR understanding. Only you can do that.

byko3y - 4-6-2015 at 10:15

NaOCl + 2NaBr + H2SO4 => Br2 + NaCl + Na2SO4 + H2O
150 g of 3.7% of NaOCl = 0.075 moles
35 ml of 35% H2SO4 (0.125 moles)
12 g of KBr (has the same properties as NaBr) = 0.1 mole
Extracted first with hexane 30 g + 10 g + 12 g + 12 g
Then I added 50 more grams of NaOCl, keeping the solution acidic. The result you can see on the last picture. I don't think that addition changed anything in terms of bromine yield, just some more chlorine and bromine chloride was created. You can compare the color https://en.wikipedia.org/wiki/Bromine_monochloride
First extraction (30 g hexane added) :
IMG_0793.JPG - 98kB

Second etraction (10 g more added) :
IMG_0797.JPG - 97kB

Fourth extraction (+2x10g) :
IMG_0798.JPG - 91kB

After addition of more 50 g of NaOCl solution and 12 g of hexane (the hexane layer had transparent red color, not shown):
IMG_0802.JPG - 96kB

I'd say that extraction leads to even less byproducts than distillation, because the BrCl has a chance to react with bromides, yielding more Br2.
Probably my final product contains a decent amount of BrCl, but it's at least 80% bromine. For some reason it smokes in air, probably it's a mixture of Br2+BrCl.
If you use slightly less than stochiometric amount of hypochlorite and use extraction for bromine separation, than you will leave something like 10% of bromine in the solution, but the extracted bromine will be 95%+ pure.
Bromine attacks polypropylene, so use glass or PTFE stoppers. I've used glass stopper tappered with PTFE and sealed with paraffin.

[Edited on 4-6-2015 by byko3y]

Oscilllator - 4-6-2015 at 19:01

ByKo3y, if you conducted your experiment in a brightly lit room you could have contaminated your hexane with hexane chlorides, since the UV light present would cause the bromine to react with the hexane.
Also, it is highly probably that your hexane is not very pure and contains some alkanes, especially if you distilled it from petrol and the like. If you want to test for the presence of alkenes you can add a very small amount of your bromine to some hexane, and then observe to see if the red colour disappears. If it does, your hexane is contaminated.

byko3y - 4-6-2015 at 20:35

Oscilllator, you are so right. Today at the morning I found the glass stopper popped high in the air and strong stream of something was coming from the bottle. Most likely it is indeed HBr.
Also, bromine has poor solubility in hexane - I was too stupid to use it, but it was too late, because when I realized I'm using the wrong non-polar solvent, the experiment was almost over.
My hexane is 99% pure reagent grade. No, it does not contain any olefins.
So there will be no report on yields, sorry.
I'd say that bromine is a really nice reagent for alkanes halogenation ^_^. Much more usefull than chlorine that boils at -34°C. You can halogenate hydrocarbons starting from butane/isobutane/butylene using a home freezer. Nice way to 2-butanol/tert-butanol.
PS: I've just understood I can actually measure the amount of hexane that was halogenated. However, there will be no way to distingiush between chlorinated and brominated.

[Edited on 5-6-2015 by byko3y]

morganbw - 5-6-2015 at 07:10

Quote: Originally posted by ScienceHideout  
Quote: Originally posted by morganbw  


You seem to have a better understanding of the chemistry than I do, however, unless you have more need to understand the chemistry, versus actually getting the product????

There are some tried/proven methods here???

If you life, is as a research chemist then, by all means, do your experiments, I bow to you.

If you just need the element sir, then you do not need to reinvent the wheel.

For me, the years are passing and time is becoming a factor, I lean on others experience.

[Edited on 4-6-2015 by morganbw]


morganbw,

If you just do chemistry for the hell of it, perhaps you should try a different hobby. Byko3y is doing chemistry the RIGHT way. Why would anyone want to do something that they don't understand?

Everyone should aspire to know what they do and do what they know, and some are better at it than others. If you are jealous at the incredible understanding byko3y has with this topic, good. I am too, and it comes with years and years of discipline and study. However, I don't think that snotty remarks will better YOUR understanding. Only you can do that.


We should perhaps share a box of tissues:)

I actually admire (byko3y ) he seems to be well grounded and a thinker (that is admirable).

Using known reactions does not imply a lack of understanding, and it for sure does not mean (just for the hell of it).

I have seen some who even consider this as a part of research.

I am not able to type very well this day, broke a shoulder blade as well as a leg yesterday. :(

byko3y - 7-6-2015 at 13:50

For some reason nobody says about bromine cumulative toxicity, and it actually exists. I've got slightly poisoned by it, because I was sure it is completely safe until you can breath:
"Potassium bromide (KBr) is a salt, widely used as an anticonvulsant and a sedative in the late 19th and early 20th centuries, with over-the-counter use extending to 1975 in the US...
Bromide ion is a cumulative toxin with a relatively long half life (in excess of a week in humans)".
Symptoms: reduced cardiac activity, low arterial pressure, general weakness, headache, memory loss, skin rash, decreased skin sensitivity, general depression, bronchitis, conjunctivitis, damage to the kidneys, delirium.
Antidote - sodium chloride and a lot of water.
This shit needs a lot of time to be removed from the body :( I have to admit that breathing chlorine on a regular basis is much safer.

Bromine source and synthesis

Brom - 19-7-2015 at 10:13

I use the the method from my old high school chemistry book. NaBr+MnO2+dil. H2SO4. If you have an all glass distillation setup this is definitely the way to go. Very fast and simple.

image.jpg - 173kB

Bromine source and synthesis

Brom - 19-7-2015 at 11:04

I was able to obtain approx. 10 ml or so in about 15 minutes using very small quantities of reagents. As soon as I get a suitable storage container I will scale it up a bit and make a larger amount.

image.jpg - 155kB

diggafromdover - 20-7-2015 at 07:50

Very nice using heat gun as heat source. Where I come from their main use is thawing pipes and occasionally setting fire to the buildings in which the pipes are being thawed.

Brom - 20-7-2015 at 10:27

I had a feeling someone would notice my heating method. I need to get a proper heating mantle but they cost as much as the entire distillation apparatus! But in the meantime the heating gun works pretty well.

diggafromdover - 20-7-2015 at 11:49

I just bought a used Fisher Thermix for $15 plus shipping (more) but I will give your method a shot.

ChemPlayer_ - 21-7-2015 at 00:17

Awesome set up Brom - congratulations!

I see ice in the waterbath, but how cold do you have the water in the condenser? My problem has always been that the bromine is so volatile that some of it escapes to the atmosphere and it's very difficult to condense.

agent_entropy - 21-7-2015 at 08:33

@ChemPlayer_ I find that circulating a bucket of icewater through the condenser using a pump for a small garden fountain works well and condenses the vast majority of the bromine.

Brom - 21-7-2015 at 08:37

I just filled a 2 gallon jug half way with a ice and the rest water. Then I use a small fountain pump for circulation. With a 300 mm graham condenser I get 100% condensation.

Brom - 21-7-2015 at 08:43

Also what I like about this method is that you simply charge your flask with the required materials and the reaction does not start till you apply heat. No need for an addition funnel as with the H2O2 method.

Brom - 21-7-2015 at 08:46

Nice agent entropy. Great minds Think alike. I think we posted the same thing at the same time

Brom - 21-7-2015 at 08:49

Never mind I looked at the post times and you beat me to it by 4 minutes. But seriously bucket of ice with $20 fountain pump from lowes works wonderfully
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