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Author: Subject: Bromine Source and Synthesis
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[*] posted on 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.
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[*] posted on 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]
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[*] posted on 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.




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[*] posted on 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]
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[*] posted on 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.




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[*] posted on 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]
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[*] posted on 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]
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[*] posted on 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?



The single most important condition for a successful synthesis is good mixing - Nicodem
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[*] posted on 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.
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[*] posted on 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!
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[*] posted on 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. ;)




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[*] posted on 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.
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[*] posted on 23-10-2010 at 23:20


This thread is awesome. Thanks everyone for all the cool pics and tutorials ^_____________^
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[*] posted on 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
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[*] posted on 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.

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[*] posted on 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.



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[*] posted on 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.
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[*] posted on 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]




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[*] posted on 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.



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[*] posted on 2-6-2015 at 19:03
Bringing back an old topic


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




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[*] posted on 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.



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[*] posted on 3-6-2015 at 10:29


KBrO3 is relatively expensive for bromine synthesis. You can use KClO3 for the same purpose.
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[*] posted on 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.




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[*] posted on 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





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[*] posted on 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!




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