Sciencemadness Discussion Board - Bromine Source and Synthesis

Bromine Source and Synthesis

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Blind Angel - 10-4-2003 at 12:57

I was wondering, is there any OTC source of bromine that can give good yield if we try to isolate it?
I know that most cough syrup contain DMX.HBr but i don't think that the bromine quantitie is enough to be useful...
any idea?

Edit: Merged with bromine thread as this contains plenty synthesis info

[Edited on 18-8-2004 by chemoleo]

DDTea - 10-4-2003 at 16:06

Sodium Bromide is sold at Pool and Patio stores under that very name: "Sodium Bromide." I don't remember specifically what it is used for though. However, some Pool and Patio guy was selling some of the stuff on eBay. It comes in a cylindrical container, and costs $9.99.

From there, the way I would go about isolating Bromine would be the same way as isolating Iodine from Iodide salts...

First react the NaBr with Sulfuric Acid to make Hydrogen Bromide, and then react the Hydrogen Bromide with a strong oxidizer to separate H2 and Br2.

Blind Angel - 10-4-2003 at 19:27

cool, didn't thought it was that easy lol

PHILOU Zrealone - 11-4-2003 at 05:01

NaBr + H2SO4 --> NaHSO4 + HBr(g)
8HBr + KMnO4 --> KBr + MnBr2 + 5/2 Br2(l/g) + 4H2O
KBr + H2SO4 --> KHSO4 + HBr(g)
MnBr2 + 2H2SO4 --> Mn(HSO4)2 + 2HBr(g)

So:
5NaBr + 5H2SO4 --> 5NaHSO4 + 5HBr(g)
5HBr +3H2SO4 + KMnO4 --> KHSO4 + Mg(HSO4)2 + 5/2Br2 + 4H2O

And in a one pot reaction:
5NaBr + 8H2SO4 + KMnO4 --> 5NaHSO4 + KHSO4 + Mg(HSO4)2 + 5/2 Br2 + 4H2O

Br2 can be extracted with CH2Cl2, CCl4, CHCl3, CCl3-CH3, ...

Beware of its low boiling point and heavy gaseous fumes that makes it very concentrate much more than Cl2; so even if a little less reactive the concentration can make it very dangerous to handle, to inhale or to get on the skin!!!!

:cool:

Haggis - 11-4-2003 at 07:54

How about bubbling chlorine gas through a sodium bromide solution. The chlorine should knock out the bromine and for sodium chloride with the bromine free. There are many methods for generating chlorine gas, but common bleach and sodium bisulfate does the trick.

trinitrotoluene - 11-4-2003 at 21:37

Yes that method does work 2 NaBr(aq)+Cl2 = 2NaCl(aq)+Br2. Watch out for the undissolved chlorine.

Home Depot

Boob Raider - 13-4-2003 at 21:34

sells NaBr soln ~35%. It comes as a part A and part B set meant for brominating the pool IIRC. The other part contains K2SO5 soln. ~ 33%. There .... mix the 2 and you get bromine water .

You might also find brominating tabs .... they, IIRC are cyanuric tribromide. I suppose you bubble Cl2 through a soln. of that to get Br2.

DDTea - 14-4-2003 at 20:10

Bubbling Cl2 through an NaBr solution does seem like a much cleaner way to go. However, with me, I avoid working with straight Chlorine gas whenever possible; it's really not something I enjoy being around, partly because of its irritating nature, and partly because i just don't like the smell

.

However, from what I've read, straight Br2 doesn't seem much nicer...

Haggis - 15-4-2003 at 07:59

Yea, bromine gas smells quite like chlorine. They do sell respirators at the hardware store for 29.99 that will filter chlorine along with acid gases, vapors and mists. How about running the residual gas through a drying tube full of activated carbon or into a container with activated carbon in it. The chlorine should be absorbed and trapped in the pores of the charcoal.

Is it not possible

Boob Raider - 15-4-2003 at 12:42

for the free Br2 to react with the bubbling Cl2 in the solution to form other weird Halo-halides etc ???
Once you get Br2*H2O is it possible to distill it to get pure bromine ??
What do you need the Br2 for anyways ... Blind Angel ??
Oh and a Bromine burn ...... Hurts like a bitch and for a long time.

Blind Angel - 15-4-2003 at 13:35

Oh, i dont really need it, i just know source for Chlorine, (theoricly) for Fluorine and for Iodine but i didn't knew for bromine

trinitrotoluene - 12-7-2003 at 22:51

Coulden't you use HCl instead of pure Cl2 gas? I had used HCl and Ca(OCl)2 as a source of chlorine gas before. So HCl seems like a good possability maybe the Cl will displace the Br as its more electro negeative. HCl is OTC and very cheap its $3.79 a gallon (3785ml) 31% soloution.

BromicAcid - 13-7-2003 at 17:43

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.

Bromine

procrastinator - 18-8-2004 at 09:23

Recently, I tried to make some bromine. I added NaOCl (bleach w/6% NaOCl), NaBr, and NaHSO4 to a flask and saw a nice reddish-tan color in the solution. Upon allowing it to sit for a few minutes, I noticed reddish-tan fumes above the solution and a black drop on the bottom, which I took to be bromine. I then put the flask in an ice bath in an attempt to condense some of the fumes and keep more of the stuff from evaporating. Wanting more bromine, I added more bleach to the solution. Nothing happened. Then, I added some more bisulfate. Still nothing. Thinking the solution had run out of Br- ions, I added some more NaBr. Now, this is the weird part: the salt dissolved and the bromine on the bottom disappeared! Does anyone know what happened? I know that I- forms the I3- ion, does bromine do anything similar?

Edit: Merged with existing thread.

[Edited on 18-8-2004 by chemoleo]

JohnWW - 18-8-2004 at 10:14

It looks that way. Such a Br3- anion, like I3-, would be linear, with the central Br being trivalent and bearing the negative charge.

John W.

procrastinator - 18-8-2004 at 11:51

Has anyone seen proof that such an ion definitely exists?

Polverone - 18-8-2004 at 12:21

A quick Google search for "tribromide ion" reveals a number of references to it.

oops

procrastinator - 18-8-2004 at 13:12

I was just searching "Br3-". Curse my lack of googling (sp?) skills.

:edit:
Okay, I just added more hypochlorite/bisulfate to the solution. instead of giving me liquid bromine, I got a piss yellow solution! :mad:

So, I've concluded that a slight excess of Br- in solution will destroy some Br2 and that an excess of Cl2 will also destroy it! :mad:

Oh, well

[Edited on 18-8-2004 by procrastinator]

vulture - 18-8-2004 at 13:38

I'd suggest you'd carefully boil down (not to dryness ofcourse)the solution outside to destroy all the hypochlorite that might still be in there.

Then put it in the freezer for a night and see what happens. If that doesn't work, KMnO4 + H2SO4 time baby!

procrastinator - 18-8-2004 at 14:55

Quote:

Originally posted by vulture
I'd suggest you'd carefully boil down (not to dryness ofcourse)the solution outside to destroy all the hypochlorite that might still be in there.

I'm pretty sure that'll also make most of the bromine evaporate as well. Remember that I started out with only 6% NaOCl. Also I unfortunately don't have an expendable freezer.

Although I'm putting a good deal of ice into it. Anyway, if I finally do get anything from this batch, I'll store it in a glass vial. Unfortunately, the underside of the cap is made of some junk which will no doubt liquefy on contact with the bromine just like it did with iodine. Does anyone have any ideas about insulating it with something inert? PE & PP probably won’t work and my PTFE tape is completely permeable to corrosive vapors.

BromicAcid - 18-8-2004 at 18:07

Bromine preparations:

Bromine presents a challenge different from producing chlorine as hydrobromic acid is not readily available.

As mentioned, permanganate, bromide, and acid produce bromine.

Hypochlorite, bromide, and acid produce bromine. (Either by the insitu production of chlorine and it's action on the bromide, or by the liberation of free bromine by hypochlorite oxidation and the bromine, or by the liberation of free bromine by one of the two aforementioned methods, followed by the replacement of the chloride in the hypochlorite with bromine, and the decomposition of the hypobromite anion.)

As shown in the picture that is attached to this, H2O2, H2SO4, and bromide produce bromine.

I've heard that adding concentrated H2SO4 to a bromide also produces an appreciable amount of free bromine by oxidizing the bromide anion in addition to the HBr produced.

Another source of bromine that is over the counter is 1-Bromo-3-chloro-5, 5-dimethylhydantoin which is one of the bromine carriers that is available for treating spas. But I'm not sure how one would recover the bromine from this.

Bromine is my favorite halogen. My method with H2O2 and H2SO4 only gives a 50% or so yeild with respect to sodium bromide, some of this is due to inefficent condensing and evaporations from my receiving flask, but these really are not acceptable looses.

Br2setup.jpg - 49kB

H2O2 method

procrastinator - 18-8-2004 at 18:28

So, adding caro's acid to a bromide should give free Br2. Couldn't it just be decanted instead of having to be distilled? NaHSO4, HBr, H2O, H2SO4, H2SO5 should form a separate layer from the nonpolar Br2, right?

OT, I'm slowly realizing that my name may make me seem kind of...dumb. Is there any way to change it or do I have to re-register?

BromicAcid - 18-8-2004 at 18:41

The bromine could probably be siphoned off, but my main reason for distilling is that the reaction is exothermic so distilling actually requires no heat. I dissolve my NaBr in 35% H2O2 and add the H2SO4 (conc.) slowly with stirring, the Br2 comes off.

So if you keep the reaction mixture very cool it should keep the Br2 in the main mix though.

And yes procrastinator, I believe you have to reregister.

One other thing that I mentioned before, the activity series of the halogens when it comes to oxocompounds is reversed. Therefore free bromine will replace chlorine in hypochlorite.

Addition of hydrobromic acid to a solution of bromate results in very high, almost quantitative bromine yields.

bromine

mick - 19-8-2004 at 11:52

When I was young, I used to love distilling bromine, I think it came from potassium bromide and I collected it under water cooled with an ice bath. A lower layer of bromine formed once the water was saturated with bromine. Makes you cough.
Mick

patu - 20-8-2004 at 11:00

I've got a good way of getting about a ml of bromine. First you get a big volumetric flask or a clear plastic container with a lid and fill it with chlorine from a bleach and hydrochloric acid reaction. You can get sodium bromide packets at walmart or any pool store. dump a small packet into the container and shut the lid. Within about two minutes the entire container turns from green to a very deep red from the bromine liberation. the chlorine is not dry, the excess water collects at the bottom and hardens the newly made sodium chloride which prevents it from being loose. This is the fun part. Take one of those liquid air cans to clean computers, turn it upside down and spray the container. the liquid instantly freezes on the urface and the bromine condenses. flip the container upside down and take off the lid the liquid bromine will happily drip into a small jar surrounded by ice. if this is done a couple of times you can get a ml of bromine. As for containing it for any length of time, all bets are off.

thefips - 24-8-2004 at 12:24

I made bromine from dilute sulfuric acid,potassium bromide and MnO2.This mixture was distilled.On the bromine was a water layer,which can be removed by separation,or the bromine with the water can be stored in a small bottle.To get the bromine use a syringe to get through the water layer.

garage chemist - 24-8-2004 at 12:42

I made bromine by mixing H2SO4, NaBr and KClO3.
The KClO3 is actually a powerfull enough oxidiser to oxidise HBr to Br2 in aqueous solution at room temp. It works great, I've tried it. NaClO3 can also be used, of course. Even the 53% NaClO3 weedkiller (or the famous "Unkraut EX"

works if you add the correct amount. The flame retardant doesn't disturb the reaction.
You have to use 6 moles of H2SO4, 6 moles of a bromide (NaBr) and 1 mol of a chlorate.
The Bromine is then separated from the reaction mix by distillation.
Unfortunately, NaBr can only be obtained from chemical suppliers where I live (Germany), and it's quite expensive.

H2O2 can also oxidise HBr to Br2 at room temp.

Caution

hodges - 4-10-2004 at 15:11

I ended up with about 3g of bromine vapor in my kitchen! I did the following:

2NaBr + 2H2O2 + 2HCl --> 2NaCl + 2H2O + Br2

I decided to make around 1ml of liquid bromine, which would be about 3 grams given the density of bromine. So I dissolved 4g of NaBr in about 5ml H2O. I added 4ml of muriatic acid and 5ml 35% H2O2, based on my mass calculations from the above formula. I did this is a tiny plastic cup.

At first not much happened, and the solution just turned orange. But then I noticed some gas bubbling out. Apparently the muriatic acid contains a bit of impurity (possibly iron?) and it was causing the H2O2 to decompose. The reaction got faster and faster until apparently the boiling point of the bromine was reached. I ended up with a kitchen sink half full of red vapor. I quickly opened all the windows and put an exhaust fan in the window. The bromine gradually migrated out of the sink and into the surrounding air. I could see a slight brown color to the entire side of the kitchen nearest the sink. I stayed out of the kitchen. In the other rooms, the bromine was strong enough to burn my eyes slightly but not enough to cause coughing. Even my pet cat, who usually pays little attention to my chemistry experiments, started complaining to me about the smell.

It took about 45 minutes to exhaust enough of the bromine out of the rooms so that it was no longer too noticable. What I had left behind was a weak solution of bromine water, but no liquid bromine.

BromicAcid - 4-10-2004 at 15:33

The reaction itself, oxidation of the bromine anion, is exothermic. Although impurities may have been a contriuting factor, when I do this on the large scale it boils due to the heat of reaction and the heat of solution of sulfuric acid (okay, mostly the latter, but with HCl alone it will heat rapidly after initiation). I have pictures of this up around here somewhere, it works well if contined in a distillation appratus, I guess I should have said not to do this out in the open.

fritz - 5-10-2004 at 09:06

The way I made Br2 were the mentioned methods with H2SO4; KBr & KMnO4/MnO2. My aim was not the production of pure Br2 but Br2-water, so the bromine distilling over was solved in water. I read somewhere that this should be a good method for preventing the bromine to evaporate from the receiver flask (if it is collected under a layer of water). As mentioned it could be separated by means of a syringe.

To Bromic acid:
your set-up for Br2-Distillation looks very good! I would suggest to fill the bucket, where the cooling-water for the Liebig-condenser comes from, with ice/water, so your apparatus perhaps won?t be filled so much with the Br2-vapours. Also you may put the receiver-flask in a real cold cooling-mixture (<0°C!), Perhaps you may use a column between reaction-flask and condenser so you could separate the Br2 from water-vapours which come from the reaction mixture (also perhaps the method jubrail mentioned with conc. H2SO4 instead of diluted may give Br2 without much water)

Another more common suggestion is to blow the bromine from the various reaction mixtures with air or some inert gas out of the reaction flask. Maybe it?s possible this way to carry out the reaction at a lower temperature so the Br2 may be condensed faster.

About the reactions with Cl2/Br-:
May be as mentioned above you may blow the bromine out of the flask (with hot Br?solution) with a Cl2/air mixture. In this case I would also suggest a long column to condense as much water as possible. With the out-blowing of the bromine the oxidation to BrCl, BrCl3, or HBrO3 may be avoided. If someone needs some pure Br2 he could try to distil the bromine from KBr (reacts with the chlorides as well as with bromic acid to give bromine). Drying could be carried out by shaking with H2SO4 (and separating) or distilling from P4O10

hodges - 5-10-2004 at 14:26

It was my reaction (with the H2O2), as well as my decision to do it in the open. I certainly don't blame BromicAcid or anyone else - I just wanted to post the results of what I tried as a caution in case anyone else had an idea similar to mine.

Theoretic - 6-10-2004 at 09:03

H2SO4 can be tried instead of HCl for the H2O2 method. Or SO3 can be used on its own to oxidize NaBr:

2NaBr + 2SO3 => Na2SO4 + SO2 + Br2.

SO2 and Br2 don't react, so bromine can just be condensed out.

hodges - 7-10-2004 at 15:26

I re-tried my experiment using the same quantities of reagents (4g of NaBr, 5ml of muriatic acid, and 5ml of 35% H2O2). This time I chilled the NaBr/HCl mixture in an ice bath and added the H2O2 SLOWLY. The reaction temperature stayed under 10C. This time I did not get a runaway. The solution slowly turned from yellow to orange to red.

I then covered the glass with a piece of plastic wrap and placed it in the refrigerator. Unfortunately, it doesn't look like any free bromine was formed (only bromine water). The air above the solution is yellow due to the bromine, but I can't see any little drops of bromine in the solution.

I used little enough liquid that only a fraction of the bromine should have dissolved, based on bromine solubility. I tested the pH and it is still <1, and I see bubbles of oxygen being slowly evolved still so I'm sure I still have H2O2 and H+ present. I'm thinking that the bromine is dissolving due to the presence of NaBr, much like iodine dissolves in NaI solution.

Since the bromine is not much use to me in the water, I had thought of trying to make bromoform, given that a way to make iodoform starts with iodine dissolved in a water solution of sodium iodide, and what I have is bromine in a water solution of sodium bromide. However, I'm sure I still have H2O2 present so I'm not too anxious to add acetone to my acidic solution. Perhaps if I added a small amount of MnO2 it would decompose any remaining H2O2 and then I could add the acetone once all the H2O2 is gone?

BromicAcid - 7-10-2004 at 15:42

Besides, the haloform reactoin is run under basic condtions so you at least have to basify it. And that eliminates the worry of acid+oxidizer+acetone, although you might have a basified peroxide solution that does have a lower oxidation potential of an equivilent volume of H2O2 with acid.

Theoretic - 8-10-2004 at 12:07

You could use Na2SO4 to absorb all the water, since 15.4 grams of it will absorb 18 grams of water, so you don't need very much (alternatively CaCl2, 11 grams for 10.8 grams of water absorbed). All your bromine will be freed, the remaining H2O2 will be catalytically decomposed to H2O and O2:
H2O2 + Br2 => 2HBr + O2
2HBr + H2O2 => Br2 + 2H2O.
The sodium bromide will stay as is, you can then separate it by distilling the dry residue with NaHSO4 (in contrast to H2SO4, it will not cause oxidation of the bromide ion).

hodges - 8-10-2004 at 14:47

I now get the feeling that, for whatever reason, I never had very much bromine to start with. For one thing, the water was red, but not extremely dark like I would expect if there was a high concentration of bromine or even tribromide. Look at how dark a 3% or less solution of iodine is, for example. I should have had more than 3% concentration of bromine and it was not nearly that dark.

Today I tried adding some NaOH. The solution, of course, became colorless, as the bromine was presumably changed to hypobromite. I added a small amount of acetone and - nothing happened. No cloudiness to the solution. After several hours the solution was clear, except for a few tiny pieces of acetone peroxide that had formed despite the basic solution (I measured pH > 12).

Does H2O2 not oxidise bromide very well or something? I know it works fine for oxidizing iodide, but of course it doesn't work for chloride. So I'm thinking maybe it doesn't work too well for bromide, and that I need something stronger next time.

neutrino - 8-10-2004 at 18:44

When I successfully made the stuff, the water wasn't too red. The odd thing is that I can't reproduce my results--this is the only successful trial I had. If I remember correctly,
there is about 1 cm of water in this flask and a good deal of Br2 vapor.

P.S. I am the procrastinator earlier in this thread.

bromine.JPG - 41kB

More experimentation with bromine.

BromicAcid - 20-10-2004 at 12:07

I tried to prefect the method of preparing bromine by the following method:

2KBrO3 + 10NaBr + 6H2SO4 ---> 3Br2 + 5Na2SO4 + K2SO4

I've tried two different bromate sources, potassium bromate and barium bromate, the barium bromate gave the predictable result of a solid in the bromine the collects at the bottom, of roughly the same specific gravity and is hard to separate. The potassium bromate gives good results, actually, great results with HBr(aq) I would go so far as to say quantitative. However all my HBr has to be made produced by my self and it is therefore a scarce commodity.

So I've been working on the use of sodium bromide, which I have great access to. The problems with this method being that the sodium bromide hardens when water contacts it, the whole reaction mixture has a habit of forming a mass at the bottom of the mixture that does no react further, and in order to get them to react completely the mixture has to be diluted again and again to dissolve whatever film is preventing the two reactants from reacting with the acid further therefor resulting in significantly lessened yields.

I was doing good today, enough reactants had been added to the graduated cylinder to form approximately .125 mol of Br2 and ice cubes where added to the powder, then dilute (36%) H2SO4 was added slowly. Immediately everything turned red then orange at the top where the iced cubes floated to. Red liquid was collecting rapidly at the bottom mixed with unreacted KBrO3/NaBr and the whole volume at the bottom was down to about 8 ml. I stirred the mix and came back, and the bottom of the cylinder was full of crystals, lots of them. At first I thought sulfates had crystallized out, I removed some with a pipette and put them on a watch glass, instantly they melted and released a lot of bromine. "Hummm..." I thought to myself, "Bromine hydrate..." It was somewhat cold outside but I believe my excess ice was really at fault, upon allowing the solution to stand and heat up a little I watched the crystals disappear, leaving behind a small quantity of bromine.

Pretty.....

hodges - 20-10-2004 at 15:10

I wonder if KBrO3 can be make via electrolysis of NaBr like KClO3 can be made via electrolysis of NaCl (adding KCl, of course)?

neutrino - 20-10-2004 at 19:19

Why do you need to make the reaction anhydrous?

neutrino - 4-12-2004 at 13:08

I decided to make some bromine by the Cl<sub<2 + NaBr method. The chlorine was generated by adding bleach to NaHSO4. This was then led into a flask containing the dissolved bromide. I used granulated bisulfate, which turned out to be a big mistake, as it took to long to dissolve in the bleach. I eventually gave up on this and used a solution of bisulfate, which worked much better, but still retained a good deal of dissolved chlorine.

The flask containing bromide first turned reddish-brown as the Br- was oxidized into Br3- and then turned cloudy and black as this formed elemental bromine. At this point, I let it settle into a black pool at the bottom. After adding more chlorine, the air above the solution started turning clearer and no more bromine was being generated. I then removed the bromine and put into a vial, the cap of which is slowly being destroyed because the PE saran wrap protecting it from the bromine vapors is permeable. It turns out that the extra chlorine didn’t affect the bromine, as it was being added slightly above the pool. Here’s a picture of the bromine sitting in a (25mL?) vial.

Br2.jpeg - 14kB

Eclectic - 4-12-2004 at 15:54

H2O2 works very well for liberating I2 from acidic solutions of iodides, but for Br2, not so much. (K2MnO4 or CaOCl) and (H2SO4 or HCl) would be much more cost effective

psychokitty - 5-12-2004 at 13:06

If the chlorine method of synthesizing bromine works then I see no reason why the method of adding sodium bromide to sodium hypochlorite at low pH would not work equally well.

I have an idea based on something that I once read.

Apparantly, residual traces of chlorine can be removed from water by boiling it. The same technique could be applied to bromine water. One could oxidize, in a reflux set-up, slowly or all at once, sodium bromide in an aqueous solution of oxone. At the top of the reflux condenser could be placed a polyethylene tube the end of which could lead to another flask and be immersed in a suitable amount of solvent which could serve to absorb the bromine vapor. The water running through the reflux condenser would have to be somewhere above the boiling point of bromine. In this way, the bromine could be collected for later or immediate use, and the yield of the oxidation could be calculated by weighing the volume of solvent after the absorption of the bromine.

A stream of air could be used in place of heat, possibly to good effect. In any case, an air bubbler easily obtained from an aquarium supply outlet in conjunction with ice-maker polyethylene tubing should be used to blow the last remaining bromine vapor out of the reflux system and into the solvent. Doing so should ensure a bromine-free atmosphere and work environment, improving the overall safety of the experiment.

[Edited on 5-12-2004 by psychokitty]

neutrino - 5-12-2004 at 16:01

There isn’t really a need for distillation during synthesis unless you want pure bromine. PE isn’t that resistant to bromine, btw. In any case, most commercial ozone-producing machines only produce tiny amounts of the stuff, while your bromine would be washed away with the unconverted gas.

As for the bleach method, I’ve had nothing but trouble with that. If you add a little extra bleach, you end up destroying your bromine. A little extra bromide and you’re stuck with the tribromide ion, which I would rather not go through the trouble of distilling the bromine out of. I’ve only had one successful run with this method, where I produced about a drop of the stuff.

Saerynide - 6-12-2004 at 01:40

A long time ago, I mentioned a drain cleaner I had which was a strong acid, but couldnt figure out what it was. It was light yellow brown. At first, it seemed like sulfuric + HCl (cause it smelled like something sulfurous when it reacted with metals and it attacks Al and floor tiles like mad). I tested to see if it was HI with H2O2. Zippo.

Finally, I thought it could be HBr, and added a crystal of KMnO4 to it. The solution started to turn purplish then turned brownish red. And then, it started to stink. God, now I know why bromine is called bromine

Esplosivo - 6-12-2004 at 21:39

Did the reddish brown substance precipitate or something, or was the reddish brown colour simply in solution? It seems strange that a company would use HBr instead of cheaply available HCl for drain cleaner right? KMnO4 also reacts with HCl to liberate light green fumes, and if not acidified a suspension of brownish Mn Oxide forms IIRC, which takes a long time to ppt - you see now why I asked about the colour of the solution Could you post a couple of pics of the stated rxn please.

[Edited on 7-12-2004 by Esplosivo]

[Edited on 7-12-2004 by Esplosivo]

Saerynide - 7-12-2004 at 01:51

It was Br2. Im positive it was not Cl2 and it was not I2. Cl2 and I2 are nice smelling in comparison to Br2

I'd like to post pics, but I am NOT doing that again

I used about 0.5ml of drain cleaner, and 1 KMnO4 crystal, and that made the whole bathroom reek for over an hour - the window was open too. As soon as the crystal stopped reacting and I had a good look it at from all angles (and smelled enough bromine

), I immediately dumped the few drops down the drain and turned on the tap to wash it down good, but god, it still stunk to high hell

After my unfriendly encounter with Cl2/HCl, Im quite scared of halogens indoors... Maybe when my parents arent home, then I'll try to get some pics.

For the time being, I'll try to describe it better. The drain cleaner was light yellow/tan at first, kinda like this #E1,F0,9D. Upon adding the KMnO4 crystal, it started to bubble, dissolve, and make wisps of this color, #AB,03,92, the regular KMnO4 color. Parts of it was yellow and other parts purple. These started to mix and it started to change from purplish (#7E,14,47)to dirty brown purple (#63,01,34) (it also started to smell like halogen now and was starting to stink), to brown (#99,33,00), and eventually to really dark red brown (#46,16,00)<- not very accurate - cant really replicate the brownness. In transmitted light, it was clear, dark red brown (more red than brown, like #79,00,00), but in reflected light, it was brown black (#2F,0F,00) and kinda opaque looking.

Hope this helps you guys a bit. I am curious and would like to try it again, but yes, I am quite afraid to

Boochit

psychokitty - 7-12-2004 at 23:23

Neutrino: Cl2 added to water forms hypochlorous acid. Addition of NaOH to this forms sodium hypochlorite. Addition of NaBr formes NaCl and sodium hypobromite. Addition of enough HCl to neutralize the sodium hypobromite forms hypobromous acid; additional HCl forms Br2.

I don't understand why your experiments failed but it's not because there's any difference between gassing a water solution of NaBr with Cl2 and acidifying a sodium hypochlorite/NaBr mixture with HCl.

Esplosivo - 8-12-2004 at 00:22

Quote:

Originally posted by psychokitty
Neutrino: Cl2 added to water forms hypochlorous acid. Addition of NaOH to this forms sodium hypochlorite. Addition of NaBr formes NaCl and sodium hypobromite. Addition of enough HCl to neutralize the sodium hypobromite forms hypobromous acid; additional HCl forms Br2.

This much depends on the conditions the reaction is subjected to. Bubbling Chlorine through water does not produce HOCl solely. HCl is also produced as is evident from a simple reaction:
Cl2 + H2O --UV--> HOCl + HCl
( If no UV is present chlorine simply dissolves in water to form what is known as 'chlorine water' )

Addition of NaOH therefore produces both NaOCl and NaCl. Bubbling Chlorine through cold and dilute NaOH would produce the chloride and hypochlorite (Chlorate (I)) while bubbling chlorine through hot and conc NaOH woud give the chloride and the chlorate (V). All these are dispoportionation reactions.

[Edited on 8-12-2004 by Esplosivo]

neutrino - 8-12-2004 at 03:18

The reason my experiment failed was that I didn't know the exact molatity of hypochlorite in the bleach and therefore couldn't get the quantity of other reagents just right. It is true chemically that the redox reactions are similar, but in practice the bubbling method works much better.

[Edited on 10-12-2004 by neutrino]

And?

psychokitty - 9-12-2004 at 18:32

You likely didn't know the exact molality of hypochlorite formed by the addition of Cl2 to water but that didn't seem interfere with the progress of your oxidation of NaBr to Br2.

There must be some other explanation as to why your first few hypochlorite experiments failed.

neutrino - 9-12-2004 at 19:13

What Cl2? The whole point of the one-pot reaction was to avoid having to use chlorine gas. What exactly are you getting at?

BromicAcid - 9-12-2004 at 19:18

Too much hypochlorite and the experiment may fail, too little and it will fail, just right and added too fast or too slow and it will fail. One main side reaction eats up yields:

Br2 + 2OCl- ---> 2OBr- + Cl2

Solution has to be acidified afterward to recover the bulk of the bromine, the activity series is reversed for oxo halogen compounds. With stright chlorine this is adverted however inner halogen compounds are more likely to form, though they usually have short half lifes in aqueous mediums.

The Cl2 you bubbled in . . .

psychokitty - 9-12-2004 at 21:38

Cl2 bubbled into an aqueous solution of NaBr = Br2 and NaCl in water, assuming, of course, exact stoichiometric amounts; however, too much Cl2 in solution will form hypochlorous acid. Thus, you're back where you started as far as the problems you had using sodium hypochlorite as an oxidizer in this reaction.

The only difference between sodium hypochlorite and Cl2 in water is that one is basic and one is not. Got it?

Damn . . .

BromicAcid - 9-12-2004 at 22:07

Yeah, hypochlorous acid is formed if an excess of chlorine is added, however as long as the solution is not basic (which it wouldn't be) then less then 2 g of chlorine will be solvated in the equilibrium:

Cl2 + H2O <---> HCl + HOCl

At any time, therefore the hypochlorite problem is largely eliminated.

neutrino - 10-12-2004 at 03:26

Ah, you’re referring to my recent experiment. The hypochlorite reaction isn’t really that much of a problem. For the most part, the hypochlorous acid would form near the surface, as the chlorine is bubbled about halfway down into the solution, while the bromine collects in a pool at the bottom. Aside from that, I discontinued the chlorine when the color of the atmosphere began to change, indicating completion of the reaction.

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.

enima - 21-12-2004 at 11:36

Oxone and a bromine salt will yield elemental bromine, I'm not sure your purpose for its usage, if you are using it for aromatic halogenation I'd recommend using ammonium bromide in an acetic acid solvent, dripping 30% peroxide onto the solution (containing your aromatic of course) stir for 5-6 hours and you should have a nice brominated product..

If it is another use, you can try forming the HBr with sulfuric acid (So2 is produced too) and then you can drip peroxide onto that, (keep it cool at 0C when doing the h2o2 dripping) this should yield bromine.

You can purify it via distillation, although IMO its best to keep away from elemental bromine whenever possible.

Nick F - 21-12-2004 at 12:13

Ah, I can't wait to get my NaBr! Lots of things I want to try...
Apparently BrNO2 can be made with Br2 and NaNO2 (damn, or was it NO2 and NaBr? Some red gas and some white powder, I forget which way round it was! :p), that could be quite fun stuff

.
I also want to use it for making NaOBr solutions to try to make hydrazine with, just because it's easier for me to make lots of bromine than it is to make lots of chlorine. Concentration of the hypohalite solution seems to play a big role in determining how much hydrazine you get out. I think it should be easy to make a solution equivalent to >15% NaOCl solution. And sodium bromide leftovers could be easily recovered.... One concern is adding the sulphuric to ppte the hydrazine - does dil. sulphuric still oxidise HBr? That could destroy your hydrazine real quick....

aludel - 22-12-2004 at 14:52

First post here after some reading - great forum! My chemistry passion is awakened after 25 years of slumbering... unfortunately I can't set up a proper lab in my apartment at the moment but I still have some chemicals.

One of my favourites is bromine and I tried to make it a few times using KBr, MnO2 and concentrated H2SO4, distilling it. Bromine was developed immediately and pretty violently when I poured the H2SO4 in. Being an ignorant fool, I used a lot of rubber tubing the first time, which happily reacted with the bromine and was destroyed (and consumed most of the Br2). But I got my first glimpse (and smell!) of this beautiful element.
It's difficult to cool it enough, it's very volatile. Another reason for my poor yields.
But I ended up with a few ml of bromine. Hard to store, it evaporates right through plastic screw caps.
I'll never forget the smell, it's getting into your clothes, hair, the whole house! My poor mom

HNO3 - 23-12-2004 at 21:26

I made ~2mL bromine. I started with some bleach (left over from a chloroform reaction

). I then added some NaBr. This formed a dark yellow solution. I then added 3mL 31% HCl. there was a barely noticable reaction, but no bromine was produced. More HCl. More HCl. Still no reaction. So I added some Ca(OCl)2, then some more HCl. No reaction. So, I added some leftover KMnO4 solution. No reaction. More HCl. No reaction. Aha! I get my sulfuric acid out and add 3mL. That did it. Bromine starts being formed, and some starts boiling off. I hold my breath and bolt for the door (I'm not too fond of the halogens

). I take a couple breaths and go get my gas mask and put it on. I go back in. I soon discovered that its filters needed replaced, as it was leaking bromine fumes. Oh well, it wasn't that bad. I dumped the top layer off and poured the bromine and a little bit of solution above it into a graduated cylinder. I then pipetted the bromine into a home-made glass ampoule. Then I made a neck on the ampoule and sealed it shut. There is one picture at http://www.geocities.com/stwrt_kck/mypage.html
Another one

Br in Grad cyl 6.jpg - 54kB

Theoretic - 24-12-2004 at 09:21

"Apparently BrNO2 can be made with Br2 and NaNO2 (damn, or was it NO2 and NaBr? Some red gas and some white powder, I forget which way round it was! :p), that could be quite fun stuff ."

Both would work I think.

No, really, both of these reactions would work:
Br2 + NaNO2 => NaBr + BrNO2
NaBr + 2NO2 => NaNO2 + BrNO2.
So NaBr (or NaNO2) can serve as a catalyst for combining bramine and NO2 into nitryl bromide.

Bromine Source and Synthesis

Protium - 25-12-2004 at 20:39

I have tried making bromine a few times.

I've tried passing Cl2 through NaBr in H2O solution, then distilling...
didn't work well...

I tried dripping conc. H2SO4 directly onto dry NaBr salt in large RB flask. What ended up happening was that the NaBr formed a solid chunk that was encapsulated within a shell of NaSO4 so that no further reaction occurred. Some (few ml's) actually distilled over being caught in an acetone/dry ice cooled falsk. Magnetic stirring did little to break up this chunk of NaBr trapped in NaSO4.

My most recent idea was this...

In a large test tube, one would add a portion of xylene, then adding conc. H2SO4 to the tube while it is tilted so that the H2SO4 would flow nicely to the bottom of the tube. Then addition of NaBr in small amounts would cause bromine formation.

I tried this because I thought that the bromine would be soluble in the xylene and I could just remove the xylene layer and distill the bromine directly because of the large difference in bp's of bromine and xylene.

In a trial run, a small amount of NaBr was added to a large test tube as described above, though reactants were previously cooled to about 5 degrees C. Immediately small amounts of gases started to form. A tinge of orange color formed as well. A pH test paper placed across the top of the tube indicated that some HBr escaped the system without being oxidized by the H2SO4. I added a few more spatula tip-fulls of NaBr and let tube sit for a while as small amounts of gasses continued to escape into the environment....no fume hood

.... It turned out that the bromine had all dissolved in the H2SO4 rather than the xylene that I had thought it would dissolve in.

I have never been able to get a layer of bromine to form on the bottom of any experiment to where I would merely be able to pipette it out. Should I dilute my H2SO4 with H2O?

I would like to collect around 50mL or so of Br2.

What method would be best for collecting larger amounts?

neutrino - 26-12-2004 at 12:39

Bromine is insoluble in sulfuric acid, that’s why it’s usually dried in sulfuric acid. How long did you let it sit before giving up on it?

I don’t know why your bubbling experiment failed, mine worked very well. My chlorine was being bubbled about 1cm below the surface of the liquid (~100mL of NaBr solution in a 500mL round bottom flask) through the end of a Pasteur pipette. The pipette was stuck through a thermometer adapter (without the rubber part) and loosely plugged with plastic wrap. Even with my archaic bubbler, I still got a good yield.

HNO3 - 27-12-2004 at 09:45

I tried passing Cl2 over NaBr before. I was using trashy equipment (plastic bottles, anyone?). The bromine formed colored the plastic orange/brown, but I didn't get any bromine.
The problem with your NaBr/H2SO4/xylene experiment is that the halogens do a substitution reaction with organics. Br2 + C6H4(CH3)2 --> C6H3Br(CH3)2 + HBr.

p-methylbenzylbromide

psychokitty - 27-12-2004 at 18:53

The bromine will add to one of the methyl side chains of the xylene. I believe I read somewhere that the product of this reaction, p-methylbenzylbromide, is a powerful lacrymator. Use another solvent, one that is inert to bromination.

HNO3 - 27-12-2004 at 18:59

I wondered if it wouldn't react with a methyl group, but I wasn't sure, so I went with what I knew.

kryss - 2-1-2005 at 16:33

The easiest way of making Br2/Cl2/I2 from the halides is by mixing a solution with MnO2, then adding dil H2SO4 and gently warming or distilling. With KI/NaI you'll get loads of I2 produced quickly and exothermically ( easy to lose control), KB2 is a little less violent and NaCl will give you loads of Cl2 on warming.

phanchem - 12-1-2005 at 14:05

About handling liquid bromine: I once had to get 10 ml of liquid bromine out of the lab refrigerator. So here I was, pouring out of a half-liter glass bottle...and the mass of frozen bromine lurking on the bottom of the bottle sloshed about an ounce of the liquid all over my hand once I got past a certain angle in my pouring. Fortunately, the sodium thiosulfate was nearby, and I could neutralize it in place, the only lasting evidence of the spill being a brown crust of brominated skin on the back of my left hand that wore off in a week or so. Refrigerator temp was turned too low--profit from my mistake and don't repeat it.

HRH_Prince_Charles - 13-1-2005 at 04:35

Would it be impudent of me to suggest wearing gloves?

Writeup of today

BromicAcid - 16-1-2005 at 14:29

Premise:
Upon mixing HBr(aq) with KBrO3 a top water layer and a bottom bromine layer was observed, the bromine appeared to be clean, no precipitate, and appeared in large quantity indicating a good yield. When the reaction is done with NaBr and H2SO4 with KBrO3 dirty bromine (almost a suspension of bromine with precipitate) appears in the bottom. My attempt today was to make a solution of HBr in situ with NaBr and H2SO4, cool to allow the sulfate to precipitate, then add the bromate in the cold weather.

Equation:
10NaBr + 2KBrO3 +6H2SO4 ---> 6Br2 + 5Na2SO4 + K2SO4 + 6H2O

Procedure:
To 200 ml of room temperature distilled water in a 1 L Erlenmeyer was added 120 g (~1 mol) NaBr*xH2O, with slight agitation the sodium bromide was dissolved in 7 minutes. 92 ml (a slight excess) of concentrated H2SO4 was measured out in a graduated cylinder and was added in several portions with significant swirling to the sodium bromide solution. A slight red color from bromine evolution was observed and the reaction mixture heated to an estimated 40 C once all the H2SO4 had been added. The mixture was allowed to cool for two hours outside at -10C. Upon inspection of the mixture it was noted that it was now a slush, the expected precipitate of Na2SO4 had either been over shot or occluded by the mixture freezing entirely, it was wrong to assume that the Na2SO4 would precipitate out nicely and the rest of the solution would remain liquid due to the freezing point depressive abilities of the HBr contained therein.

The slush and chunks were broken up to a homogenous consistency and 33.5 g of KBrO3 was weighed out, roughly .2 mol and half of this was added to the slush mixture. The addition produced a hissing sound and the mixture quickly started to turn red from bromine evolution. After the fist half had been added a watch glass was placed on top, lifting the mixture drops of bromine were evident on the bottom, at these temperatures the bromine was significantly less volatile and only small amounts wafted off the top.

The mixture was allowed to cool for twenty minutes further outside. The addition of the second half of the bromate was smoother then the first addition, no sound or other indicative factors of a fast reaction. The color slightly deepened and bromine was more pronounced at the opening of the container and little droplets of it clung to the sides at the water line. The top was covered yet again and the mixture was allowed to react to completion over the course of another 20 minutes.

The bromine at the bottom was heavily contaminated with precipitate, separation may have been accomplished by straining through glass wool however due to the temperature outside I was ready to call it quits so I had to find something to store it in. The entire mixture was poured into a glass reagent bottle which was accompanied by significant fuming. After the bottom precipitate layer was poured in there was sediment left in the container, shaking the flask the sediment sounded distinctly like stones in the flask, removing one and prodding it with a glass stirring rod they were found to be very hard, their composition is unknown.

In the smaller reagent bottle the bromine at the bottom seems to be divided into layers, a top green/red layer, a bottom bromine/sediment layer, if this separation persists through the winter actions will be taken to ascertain the identity of the layer constituents. Theoretically roughly 80 grams of bromine (26 ml), although not measured directly the container appears to contain at least 20 ml of bromine, which would be quite good considering the bromine lost to evaporation and solvated in the water above.

Conclusion:
HBr should be previously purified to give a somewhat pure bromine product in one shot. Filtration through glass frit/wool should be considered if taking this on following this procedure to remove suspended particulate, distillation is always an option. Unlike a previous attempt, bromine hydrate was not a problem. Overall this reaction quickly gave quantities of bromine in a high yield without the need for distillation, if a pipette were to have glass wool inserted into the intake then the filtration may be accomplished by simply withdrawing the bromine from the container.

After sitting for 15 hours or so the bromine has separated into a distinct phase, the particulate has frozen into a layer above it and does not travel with the bromine. It is now believed the solid rock like pieces left in the reaction flask are pieces of bromate that went into the mixture too quickly and formed an inert crust. I plan on just suctioning off the bromine only a thin sheet of ice stands in my way, the mixture is for the most part slushy. The above container has a volume of 450 ml.

[Edited on 1/17/2005 by BromicAcid]

garage chemist - 16-1-2005 at 14:34

Where do you get so much KBrO3?
You seem to use lots of it...
Did you buy it or make it?
It is rather carcinogenic after all.

Maybe we should open a thread on the production of bromates.

Nice bromine synthesis by the way.

BromicAcid - 16-1-2005 at 14:38

KBrO3 is used in baking as a levening agent or something of that extent, I was in a bakery store and they sold KBrO3 by the pound for $2.75 so I bought a few pounds, although my frequent use of it has reduced my stock so I must once again go to a bakery store. I wonder what they think I need that much KBrO3 for? I don't think bakery stores get warned that one of their chemicals can function as an oxidizer and they should watch the people that buy it.

Edit: Ahhh.. here it is; "Maturing agent in flour, dough conditioner, food additive."

[Edited on 1/16/2005 by BromicAcid]

garage chemist - 16-1-2005 at 14:49

Wow, very interesting!
But I never saw KBrO3 on the list of the allowed food additives (the "E-numbers"

.
I'm quite sure it isn't allowed in Europe.

Anyway, what kind of bakery did you visit?
Some professional, large scale industrial bakery?

BTW, E252 is KNO3, so you could ask for this ,if you needed it ( only in Europe of course).

"I'd like to have some E252" sounds much less suspect than "I'd like to have some potassium nitrate".

[Edited on 16-1-2005 by garage chemist]

The_Davster - 16-1-2005 at 15:28

"Causes irritation to the gastrointestinal tract. Symptoms may include nausea, vomiting and diarrhea. May cause abdominal pain, reduced urinary output, low blood pressure, methemoglobinemia, convulsions, liver and kidney damage, and coma. Cyanosis may occur as a later symptom. Death may occur from renal failure, within 1 to 2 weeks. Estimated lethal dose is 4 grams."

Sounds like something really great to have in food products.

The MSDS must be exagerating...
Source:
http://www.jtbaker.com/msds/englishhtml/p5576.htm

verode - 22-3-2005 at 06:29

the NaBr with ice + Cl2+ some HCl
You get Br2 Its should be acid becose it could happen Br->HBrO3
The Br2(l) is heavy than water

HNO3 - 18-5-2005 at 20:26

I'm working on a method that allows absolutely no formation of halohalides. I also don't want to use H2SO4 because the heat spike volitalizes so much Bromine. So, the current results.

I was not in a mood to measure very carefully and I didn't have my scale there, so its pinch of this, dash of that chemistry. I used Sodium Bromide, then added 6% bleach and swirled the mixture until it dissolved all the sodium bromide to form an orange solution. To this was added citric acid. Fizzing started immediately, but died down quickly, and the solution turned a deeper orange. Remembering H2O2 as an oxidizing agent, some was added. The solution turned white and what little bromine was on the bottom dissolved. More citric acid was added, but the orange color did not reappear. So, to decompose the peroxide, potassium permanganate was added (~5g). The solution turned purple, then a pond scum color, as a dense oily ORGANIC settled to the bottom. That is when I left. Did I manage to make bromoform? Or is it some other chlorinated/brominated organic? I don't know. The work continues.

Rosco Bodine - 18-5-2005 at 20:49

Anybody tried the reaction of TCCA with sodium bromide solution ?

Evidently what results is sodium cyanurate
in solution , and bromine on the bottom .

See the patent GB1401120 . The file is attached to one of the recent posts in the hydrazine thread .

Polverone - 19-5-2005 at 00:16

I tried the reaction. According to the patent, the TCCA/NaCl reaction produces appreciable free halogen only when the solution is heated, strongly illuminated, or placed under reduced pressure. There isn't a lot of free halogen produced otherwise. The same seems to be true with bromine: the sodium bromide/TCCA mixture instantly turns orange on addition of water, and you can smell the bromine, but there is no liquid bromine formed, nor any visible bromine fumes. The liquid isn't the dark color expected of saturated bromine water.

On acidification with H2SO4, the reaction began with more vigor. This is not especially surprising; the TCCA seems to act like more conventional oxidizers in the production of bromine. The addition of conc. H2SO4 heated the mixture, and after a couple small pipettes of acid, the mixture was merrily fizzing out vapors. I would be interested in seeing what the effects of heating alone without acid addition are, but that will have to wait until the fumes have cleared out of my workspace.

Rosco Bodine - 19-5-2005 at 00:28

You might try putting a layer of ether
or methylene chloride on top of the mixture and setting it in sunlight .

The solvent should move the equilibrium to the right the same as heating or vacuum , and so will the sunlight .

Of course shaking it up once in awhile or stirring it should help too and the color and the bromine should end up in the solvent .

trilobite - 23-5-2005 at 15:05

Many ethers are oxidised to esters by bromine, the reaction product HBr being also able to cleave ether bonds, so methylene chloride is better.

neutrino - 24-5-2005 at 02:26

The problem here is separating the two afterwards. They're both nonpolar and have about the same boiling point. Separation would be very difficult without a good fractionating setup.

azaleaemerson - 24-5-2005 at 04:24

Just make sure you don't get bromine onto your hands. I once was helping to move a few items off a bench in a friends lab at work one day. I grabbed the capped bottle of bromine to move it to the shelves. The lid wasn't on tight and it ran over my fingers. Lots of water washing didn't help much. My fingers were actually the minor pain. The thinner skin of the webbing was eaten away. First orange, then white, now a scarred pink. This is one of those memories I won't soon forget.

neutrino - 24-5-2005 at 13:14

Don’t worry, I’m sure that most of us here know what we’re doing. Still, bromine burns are definitely something to avoid.

Saerynide - 11-6-2005 at 01:59

Nice

Looking forward to reading more about it.

ADP - 13-7-2005 at 19:04

In to route to hydrogen bromide gas the reaction occurs in the following way:

NaBr + H2SO4 --> NaHSO4 + HBr(g)

My question: When one does this reaction wouldn't the HBr react with the remaining H2SO4 and since HBr is such a strong reducer, wouldn't it be oxidized by the H2SO4 and free the bromine?

neutrino - 13-7-2005 at 19:14

Yes, that is what happens to some extent. Small samples of bromine are commonly prepared like this.

ADP - 13-7-2005 at 19:55

If that is the case, is there a more desirable way to produce HBr gas without the contamination of free bromine?

I guess one could always perform the reaction and then heat the reagents until the HBr comes out of the solution. Also perhaps the H2SO4 could be dilluted as well the the NaBr be aqueous solution

Non-oxidising acid

wa gwan - 14-7-2005 at 06:13

The use of the non-oxidising phosphoric acid and a potassium or sodium bromide salt should produce bromine free hydrogen bromide gas.

H3PO4 + KBr ---> HBr + KH2PO4

12AX7 - 14-7-2005 at 08:53

Whaddabout plain HCl?

Tim

IPN - 14-7-2005 at 10:05

Probably quite a lot of HCl impurities mixed with the HBr.

ADP - 14-7-2005 at 10:30

Wa gwan thanks for the info I should have thought of that.

mantis - 17-7-2005 at 01:54

You can produce bromine with conc. sulphuric acid and potasium bromide.
It reacts in two steps:
H2SO4+2KBr-->2HBr+K2SO4
now the concentrated sulphuric acid will oxidise the HBr to water and bromine.
H2SO4+2HBr-->Br2+SO2+2H2O
All in all you can say:
2H2SO4+2KBr-->K2SO4+2H2O+SO2+Br2
You can save oxidation agents like KMnO4, MnO2, H2O2, etc. It´s a very cheap producing methode.

[Edited on 17-7-2005 by mantis]

hodges - 24-2-2006 at 15:19

Quote:

Originally posted by rogue chemist
"Causes irritation to the gastrointestinal tract. Symptoms may include nausea, vomiting and diarrhea. May cause abdominal pain, reduced urinary output, low blood pressure, methemoglobinemia, convulsions, liver and kidney damage, and coma. Cyanosis may occur as a later symptom. Death may occur from renal failure, within 1 to 2 weeks. Estimated lethal dose is 4 grams."

Sounds like something really great to have in food products.

The MSDS must be exagerating...
Source:
http://www.jtbaker.com/msds/englishhtml/p5576.htm

I bought a loaf of "White Enriched Bread" at the supermarket recently. I happened to notice on the label:
BROMATE FREE
UNBLEACHED FLOUR
NO ARTIFICIAL PRESERVATIVES

So apparently bromates are used in baking some types of white bread.

12AX7 - 24-2-2006 at 19:14

Well, bromate is a reasonable oxidizer, no? Probably serves the same purpose as OCl, bleaching (hence unbleached flour).

Tim

mantis - 25-2-2006 at 14:41

Quote:

Well, bromate is a reasonable oxidizer, no?

yes, but it need a acid medium.

Safer Bromine Source

Flip - 5-3-2006 at 00:01

This isn't exactly OTC, but pyridinium hydrobromide perbromide makes a good brominating agent, and it is also much safer to use if you have limited space/ don't have a fume hood. The perbromide forms a rapid equilibrium with pyridinium hydrobromide and bromine, so the bromine concentration is fairly low. However, it is convenient, safe, and easy to work with. I'm not sure what it costs, but I imagine that it will be worth it. It might not be the best for working on a large scale, but I know that it's perfect for my purposes. You can run the bromination in GAA and you'll be all set.

Djerassi, C.; Scholz, C.R. J. Am. Chem. Soc. 1948, 70, 417.

Pommie - 5-3-2006 at 15:08

The only bromine source I can get hold of is 1-bromo-3-chloro-5,5-dimethylhydantoin (BrClC5H6O2N2

. Can this be used to produce bromine?

Mike.

1-bromo-3-chloro-5,5-dimethylhydantoin

Flip - 5-3-2006 at 16:43

Sure, it looks like a strong brominating agent. The trouble is, that it will also be a chlorinating agent. Interesting chemical, actually, I would very much like know if it would be useful. Some reactions don't require a specific halogen, just that the molecule be halogenated... so if a few chlorines mucking about aren't a problem.... I imagine that you could put this to use, though I am hesitant to comment on the synthetic utility of it.

bcdmh.gif - 1kB

12AX7 - 5-3-2006 at 16:56

And if nothing else, the chlorine should displace the bromine and get you pure bromine, no?

Interesting molecule, looks like a variation on TCCA I'd say?

Tim

Pages: 1 2 3 4

Bromine Source and Synthesis

Home Depot

Is it not possible

Recent post at the Hive

Bromine

oops

H2O2 method

bromine

Caution

More experimentation with bromine.

Boochit

And?

The Cl2 you bubbled in . . .

Bromine Source and Synthesis

p-methylbenzylbromide

Writeup of today

Non-oxidising acid

Safer Bromine Source

1-bromo-3-chloro-5,5-dimethylhydantoin