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patu
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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.
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thefips
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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.
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garage chemist
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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.
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hodges
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Caution
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.
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BromicAcid
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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.
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fritz
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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
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hodges
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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.
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Theoretic
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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.
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hodges
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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?
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BromicAcid
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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.
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Theoretic
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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).
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hodges
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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.
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neutrino
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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 Br<sub>2</sub> vapor.
P.S. I am the procrastinator earlier in this thread.
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BromicAcid
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More experimentation with bromine.
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.....
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hodges
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I wonder if KBrO3 can be make via electrolysis of NaBr like KClO3 can be made via electrolysis of NaCl (adding KCl, of course)?
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neutrino
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Why do you need to make the reaction anhydrous?
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neutrino
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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. 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 Br<sub>3</sub>- 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.
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Eclectic
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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
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psychokitty
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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]
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neutrino
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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.
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Saerynide
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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 
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Esplosivo
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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? KMnO<sub>4</sub> 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]
Theory guides, experiment decides.
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Saerynide
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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
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psychokitty
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Boochit
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.
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Esplosivo
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| 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:
Cl<sub>2</sub> + H<sub>2</sub>O --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]
Theory guides, experiment decides.
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