Sciencemadness Discussion Board

Bromine Source and Synthesis

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evil_lurker - 6-3-2006 at 08:11

BCDMH is definately an interesting chemical and a darn hard one to figure out how to reduce.

The only literature on it, is what is provided by the pool and spa industry which is generally written for "those not skilled in the art".

The only practical way I can think of reducing BCDMH is by adding NaOH and heating then reducing the formed halogen compounds by adding sulfuric acid... which may or may not work.

When BCDMH is placed into water, it dissolves into hypohalous acids and hypohalic ions whose equilibrium is based upon temperature.

The addition of NaOH neutralizes the acids formed and turns them into hypohalites.

Adding heat causes the BCDMH to dissolve more readily and the hypohalite salts to undergo an auto-oxidation-decomposition into halides and halates.

The reaction would be:

BCDMH + 2 NaOH = 1 NaBr + 1 NaCl + 1 NaBrO + 1 NaClO

1 NaBrO + 1 NaClO ---- heat ---> 2 NaBr + 1 NaBrO3 + 2 NaCl + 1 NaClO3

(X represents a halogen)

6NaX + 2 NaXO3 + 3 H2SO4 = 4Br

I know the equation is unbalanced but it should liberate most of the available bromine due to the chlorine liberated at the same time as the bromine... the free chlorine would undergoes substitution reactions with the bromo compounds.

It is unknown whether or not the NaOH or Sulfuric acid would react with the spent DMH.

garage chemist - 6-3-2006 at 13:48

How about just bubbling chlorine into an aqueous suspension of BCDMH?
This should result in formation of dichlorodimethylhydantoin (the Br on the BCDMH gets replaced with Cl, due to higher reactivity of Cl) and quantitative liberation of bromine. It can be distilled out of the mix (which will have to be kept cold during Cl2 addition).

However, the bromine such obtained will contain some chlorine, bound up as BrCl and also physically dissolved.
Obtaining pure bromine from this is easy: the crude bromine is stirred with a NaBr solution for a while and then distilled out of it.
This replaces all the Cl atoms in the crude bromine with Br, again due to higher reactivity of Cl.

evil_lurker - 8-3-2006 at 14:25

The problem is that BCDMH is not very water soluble...

And, as soon as it does dissolve you end up with HXO molecules.... which may or may not be of any consequence.

I landed 3 pounds of BCDMH for $10 this afternoon at Lowes on clearance (old stock from last year) and 12 oz of NaBr for another $10 (pocketbook still hurting).

Tonight I plan on doing a little expirimenting... will post the results later on...

evil_lurker - 8-3-2006 at 15:29

Interesting results... just for shits and giggles I decided to chunk a pellet into some 10 molar HCl.

Immediatly a vigorous reaction took place and it fizzed like crazy releasing chlorine gas or perhaps a combination of CL2 and perhaps some interhalogen BrCl.

As the HCl was used up, the reaction slowed down and the solution began to turn orange... I know that there has to be some bromine in there being released.

When everything slowed down the reaction turned milky orange

The reaction appears to be non-exothermic... this could be a very very convienient source of bromine or for that matter chlorine if I can just figure out how to get everything separated!



Flip - 8-3-2006 at 16:27

Good thinking with the HCl... what with the common ion effect and such. Try it with a more dilute solution and see if you get better results. The kind of reaction you describe from the 12 M solution doesn't sound very scale-able.

I don't suppose you could get away with neutralizing the solution with excess NaOH, then just distilling out the bromine?

The_Davster - 8-3-2006 at 17:04

NaOH would react with bromine making hypobromite, and if it was hot enough hypobromite woud decompose yielding bromide and bromate.

evil_lurker - 8-3-2006 at 17:51

You can't add NaOH to an elemental halogen without making the hypohalite IIRC so that is probably out.

Actually what I really need is a more concentrated solution of HCl... or so the theory goes.

Personally I want the solution to become so saturated that the bromine simply falls out of solution... then it can be mechanically separated then dried and purified over concentrated H2SO4.

The problem with the setup above is that the Cl2/BrCl "off gases" so fast that any formed Cl2 molecules dont have a chance to react with the bromine... diluting the HCl might slow down the reaction, but it wouldn't be too good for precipitating bromine.

Tommorow I might try to snag a clear champagne bottle and build a pressure reactor... I figure if I chill the HCl down to nearly freezing the pressure built up from the reaction will be well within the safety limits of the bottle... as long as the temp stays below 40 degrees the pressure inside the bottle shouldn't get over 60PSI according to my data.

Trying to contain the reaction will be interesting. I theorize that the the formed chlorine will eventually take over everything and the bromine will fall out of solution or else the mix will reach equilibrium and stop...

Anyway you go though, a setup like the above would make a great chlorine gas generator as all it would take is a needle valve and a pressure gauge added onto the stopper.

eesakiwi - 13-4-2006 at 01:36

Shit, eesakiwi just put pool Bromine (white powder, crushed up a bit) in a tall glass tube.
Poured on some HCL acid, it foamed and spit a bit so swIm poured some Parrifin oil over it & left it.
The HCL: layer kept reacting with the Bromine powder & moved down the tube, as it reacted the reaction face it got orangey, then darker, then red, then a thicker layer of red.
The Parrifin layer stopped the smell coming outta the tube & still 'floated' above the HCL layer/reaction.
Untill it got to the bottom of the tube & there was a thick layer of Bromine, & its still there.
This was all about 1+ ysr ago.

SwIm'd just put the powder in the tube, pour over it a 15mm layer of Parrifin oil, & then use a syringe to inject the HCL into under the Oil layer.

evil_lurker - 17-6-2006 at 16:20

Tonight is the night.

I'm going to attempt a 2 mole reduction and isolation of Br2 using NaBr using H2O2 and H2SO4 which have shown to be very good in tests. All methods using the BCDMH have been pretty much fruitless. KMnO4/H2SO4 tests have been less than desirable along with any method involving chlorine. The former suffers from the Br bonding with the oxidant and the latter suffers from equilibrium issues.

Reactants are as follows.

Rooto brand sulfuric acid drain opener
Salon Care 40 volume clear developer (12% H2O2 by volume)
Dupont spa start up/reserve sodium bromide

The setup I'm using involves a 1 liter FB flask, claison adapter, 500ml addition funnel, 20 ml Barret style moisture test reciever with overflow (cadillac of dean stark traps), 400mm high efficiancy coil reflux condensor.

The idea is to drip the H2SO4 into a mixture of pre-dissolved NaBr in H2O2 similar to what Bromic Acid did on his website the first time around. But, instead of trying to collect the highly volatile condensed bromine in an ice cooled flask, it will be collected in the barret reciever. This way any fumes that are re-emitted will simply rise back up into the condensor and be re-condensed preventing vapor losses similar to what Bromic was experiencing. When its time to empty, the drain cock on the barret is opened and the contents are drained into an ice cold container with PTFE top.

The reaction pilfered off of Bromic's site is as follows:

H2O2(l) + 2NaBr(aq) + H2SO4(l) ---> Br2(g) + 2H2O(l) + Na2SO4(aq)

So when you figure it up, 100ml of 12% H2O2 contains 1/2 a mole of H2O2. So to reduce 2 moles of NaBr, 200ml of H2O2 is required. Since the rooto causes some catalytic reduction and the solution would be right up near its saturation point, a slight amount more H2O2 will be added, say another 25ml or just enough to get everything dissolved. Not very much H2SO4 is needed, but will be added until the solution is completly acidic.

Well I'm off to hook everything up. Wish me luck. Will take pics.

Jdurg - 17-6-2006 at 18:46

Good luck! Bromine is some really remarkable stuff, and any time you can make some on your own instead of having to purchase a sample of it is a good thing. (Shipping costs are kind of nuts. Many carriers don't like shipping the stuff).

I bought a good deal of bromine off of E-Bay a few years ago when you were still able to buy things like elemental bromine, sodium, potassium, rubidium, phosphorus, etc. from there. The bromine was shipped in a normal amber glass bottle with a PTFE cap. This bottle was in a plastic bag which was then put into a LARGE metal can surrounded by vermiculite. (The bottle was a 50 mL bottle as my purchase was of 100 grams of Br2). Anyway, after about a week of having the bromine I noticed that everything in the area the can was in was starting to rust severely. I opened up the can and saw that the inside was HEAVILY corroded. The Br2 had eaten right through the vial's cap and was starting to leach out! :o I called up the person I had purchased it from and they VERY quickly overnighted a new bottle to store the bromine in. (This was one specially designed to contain corrosive materials, and was basically pure Teflon).

I will never forget the smell of the bromine as I was transferring it. It reminded me of a skunk that took a bath in concentrated bleach. Nasty stuff. I made the transfer outside and was happy to be upwind of the stuff. Thankfully, I didn't have any further problems with it. I eventually got a sample of about 5 mL of it sealed in a glass ampoule for my element collection and traded the rest of it away for other elements I needed. (Like some Ir and Os metal).

In high school, I also learned first hand that bromine will leach into glassware if given a chance and remain there for MANY years. I was working on a lab in chemistry class which required some glass tubing to be bent. My teacher had me go into the supply room and get some glass tubing that was in one of the drawers. Apparently the previous chemistry teacher used it to generate and condense elemental bromine, but my then teacher did not know that. The glass looked perfectly fine to me so I went and brought it out. My lab partner started to hold the glass in the bunsen burner flame to begin the bending process when the entire tube turned slightly yellow, then orange, then red, then red-brown. A reddish-orange gas began to eminate out of it and up towards the ventillation duct in the ceiling. Stupid me was in between the glass and the duct and got a very large lungful of dilute Br2 vapor. I think I was coughing for about two weeks after that. heh. I was just shocked that the Br2 could get trapped in the glass for such a long time and not leach out.

evil_lurker - 17-6-2006 at 20:21

Wow. All I can say is wow.

Bromine vapor is some tough stuff to try and condense! Thank goodness I had my respirator on... fumes were going everywhere.

DO NOT, I REPEAT DO NOT TRY TO DISTILL BROMINE WITHOUT A RESPIRATOR ON, PREFERABLY ONE THAT IS AIR LINE SUPPLIED!

I decided to make some last minute changes, one being to up the batch size to 4 moles (408 grams) and use ACS sulfuric acid.

All in all I managed to catch 153 grams of wet reddish black bromine. I was hoping to get a yield of around 80% but I figure my yield will be about 38-43% of theoretical after the crude Br2 is dried over sulfuric acid. It just didn't want to happen for me tonight. Much was lost out the top of my reflux condensor (vent hose is red), and still there is a bunch in solution.

What went wrong:

Added acid just a little too fast (was very hard to see what I was doing due to condensed bromine water plus there is a large hole in my addition funnel)... that resulted in losing quite a bit out the top of the condensor.

The H2O2 will begin to decompose as soon as the sulfuric is added resulting in loss of oxidant before all the Br2 is released from solution. After I added the acid, I decided I would try and add some more H2O2 and got another 10 or so mL of product.

Things I would do next time:

Add some salt of some sort to lower the cooling temp down as low as possible.

Change over to simple distillation, put receiving flask under ice use reflux condensor on vacuum takeoff connected by short hose.

Possibly add some H2O2 to the acid to reduce volatility and increase available oxidant.

[Edited on 18-6-2006 by evil_lurker]

Brsetup.jpg - 87kB

evil_lurker - 17-6-2006 at 20:25

Bromine in receiver.

bromine.jpg - 56kB

Jdurg - 18-6-2006 at 07:34

Very nice! It's always more fun to make the elements yourself than it is to buy them. The sample of chlorine I have in my element collection is one of my favorites because I made it myself from calcium hypochlorite and hydrochloric acid. Whenever I get around to it, I hope to make some amorphous silicon through a reaction between magnesium powder, sand, and HCl.

evil_lurker - 18-6-2006 at 15:38

Well its nice, but yeilds are still a bit on the low side. I'm hoping for a 60% or better yield.

I'm going to try it again this afternoon. Right now I am adding H2SO4 drop by drop into a saturated solution of NaBr. My reasoning is since the oxidant for whatever reason gets reduced/lost during the process, it would be better to add it to a solution of HBr than to do it the other way around. Plus, since the H2O2 is more diluted, it will be easier to contro the rate of addition and give my condensor time to catch up.



[Edited on 18-6-2006 by evil_lurker]

evil_lurker - 18-6-2006 at 16:39

MORE LIKE IT!

This time I did a 3 mole batch.

Using the following:

3 moles NaBr were dissolved in just enough hot water to make a saturated solution.

1.5 moles of H2SO4 were slowly dripped in at a sufficiant rate to prevent HBr gas loss.

Heat was added so that the temperature of the HBr solution was over the boiling point of Br2 and left on until the end of the distillation. 12% H2O2 was slowly added at a rate the condensor could keep up with the bromine coming over and was discontinued when it appeard the Br2 began to slack off. Distilation was allowed to run for a few more minutes then discontinued

Both the cooling water and the reciever flask had CaCl2 added to reduce their temperatures. Very little bromine vapor escaped this way, and in fact I could see only a trace of Br2 vapor in the top by the vacuum exit.

Yeild has not been measured as of yet, but looks much much better than the last run.

[Edited on 19-6-2006 by evil_lurker]

BrSetup2.jpg - 106kB

evil_lurker - 18-6-2006 at 16:47

Bromine in bottom of condensor...

Brcondbtm.jpg - 95kB

evil_lurker - 18-6-2006 at 16:51

And now the results... a bunch of wet bromine!

Total yield is approx 166 grams for this run. Based on available bromine and deducting for water carried over, I'd guess that my percentage is somewhere in the 60-65% range based on the starting bromine.

Thats still not as quite as high as I'd prefer, but all things considered it would be acceptable to me.

I know its unscientific not to be doing all the measurements, but if I had to guess I'd say there is around 300 grams of Br2 sitting in my deep freezer.

Which by the way, when I opened the bottle it was under slight pressure even though the bromine had partially solidified!



[Edited on 19-6-2006 by evil_lurker]

bromine2.jpg - 97kB

Jdurg - 18-6-2006 at 17:21

Wow! That's still quite a bit of bromine there. Nice job! You may want to consider keeping it stored in a dedicated freezer. As a solid, Br2 is less likely to eat through whatever container you plan on storing it in. If it's kept in an amber glass bottle with a Teflon lid and is tightly sealed, you can freeze it solid and in the case that there is a power failure the bottle will still keep it relatively safely.

Magpie - 5-7-2006 at 10:28

To celebrate Independance Day I decided to make bromine for the first time. After reading through this thread it seemed wise to just make a small amount at first. Also I wanted the synthesis to be under good control. Hermann Rymmp's method using MnO2 (provided earlier by jubrail) seemed perfect.

I added pool spa NaBr and pottery grade MnO2 (which is very fine) to a 250mL RBF then added the dilute H2SO4 (Rooto grade). I could see a haze of orange gas at first. Using a bunsen burner I began to heat at low heat until drops of dark red liquid slowly began to roll down the condenser. I only made about 0.5mL but it was still exciting to see it form for first time.

See picture:

[Edited on 5-7-2006 by Magpie]

[Edited on 5-7-2006 by Magpie]
E.b.Chemoleo: reduced resolution of pic to make thread readable.

[Edited on 30-1-2007 by chemoleo]

Br2.jpg - 65kB

evil_lurker - 5-7-2006 at 13:08

H2O2 from Sally Beauty Supply is the way to go on the oxidant. Not so cheap, but readily available, and no stinking hazmat fees!

I found a deal on NaBr on ebay, $5.48 per pound with $6.00 flat shipping regardless of quantity. I ordered 6 pounds, and by the time I'm done I hope to have about 1.5 kilograms of bromine to play with!

Magpie - 5-7-2006 at 14:12

1.5 kilograms! You (and Bromic Acid) must really love this stuff. How do you store it? I put mine in a small Boston round bottle with Teflon lined cap. But even that made me uneasy as it is so hot in my lab. So I stuck the bottle in my freezer. It is now nicely frozen. :D

For my planned synthesis of bromobenzene I only need about 15 mL. So will have to go back into production on a larger scale.

evil_lurker did you decide that beauty shop H2O2 is OK for this application. Earlier you had some concern about buffers. Beauty shop H2O2 is readily available in drug stores and would make a convenient oxidant. MnO2 is too messy for large scale work, I think.

My bottle of "Clairol 20" lists water, H2O2, and phosphoric acid as the only ingredients. I wouldn't think phosphoric acid would be a problem here.

[Edited on 6-7-2006 by Magpie]

evil_lurker - 5-7-2006 at 20:19

Concerning the buffers, probably not. I'm not for sure, but after I rinsed out the reaction flask and left water standing in it for a coupla days some unknown "snot" formed inside. That left me to believe that the buffer to be some sort of organic compound which would react with the bromine and reduce yields. But then I got to thinking about it, all the organics that I know of and hydrogen peroxide don't mix very well. :P

I use the 40 volume clear myself. Its like $4.99 for a quart or so and is good for several moles. 20 volume is only double strength OTC hydrogen peroxide... not much use in my book... simply not concentrated enough.

For storage I been using the amber glass safety coated PTFE lined cap bottles from www.cynmar.com in the freezer. They seem to work well and are NICE. The safety coating makes them a lot more durable... they don't even "clink" when they come into contact with each other.

As far as why I like bromine so much, there is a lot you can do with it. Substitutes for chlorine very nicely and is much easier to handle.

The_Davster - 6-7-2006 at 18:58

Everyone is playing with bromine these days it seems!:)

I did the bromate, bromide, acid method tonight. The ammounts were such that as a maximum I could get 1g of bromine.
Diluted H2SO4 was added to a solution of bromate and bromide while being heated over an alcohol burner. I collected around 0.2 ml or so of bromine, with double the ammount of water floating above it.

Did I mention I managed to do this indoors with only a window open?:D Only a slight bromine smell was noticed near the end, before the system was filled with NaOH solution, and the apparatus disassembled under water, and the distillation flask containing the most bromine was imersed in the toilet bowl and flushed several times dispersing the bromine solution.
Unless of course it deadened my smell, in which case Ill be dead tomorrow:o.

EDIT: I kinda like the smell, seems similar to iodine.

[Edited on 7-7-2006 by rogue chemist]

Magpie - 6-7-2006 at 20:11

I did some test tube scale bromine making today:

1. In the 1st I dissolved 0.5g NaBr in 1 mL Clairoxide 20 (6% H2O2) then added about 0.25mL 86% sulfuric acid (Rooto). Bromine production was good. I added heat to get as much bromine produced as possible.

2. The 2nd test was with 0.25g NaBr and 1 mL 86% sulfuric acid. Bromine production started right away but was weak. Even added heat did not help much.

Just out of curiosity I calculated the potentials to produce Br2(l) from Br- using various oxidants in acid. Here's the results:

1. MnO2, Br-, and H+................+0.165V
2. MnO4-, Br-, and H+..............+0.445V
3. H2O2, Br-, and H+................+0.705V

Magpie - 8-7-2006 at 16:06

I have finished making the bromine (10-15mL) that I need for my synthesis of bromobenzene. I plan to use Br2 and iron filings per my old school procedure.

The procedure calls for dry glassware. I have noted a bit of ice crystals in my bromine bottle. Vogel says this water can be removed by washing with con. sulfuric acid. All I have readily available is my Rooto which is 86% vs the 96-98% of con. sulfuric. I hate to find out by "trial & error" when working with all my bromine. Does anyone know if the 86% would be of sufficient strength?

Duh. Maybe I should just pour the bromine off the ice.

[Edited on 9-7-2006 by Magpie]

[Edited on 9-7-2006 by Magpie]

neutrino - 9-7-2006 at 11:03

Some will still remain dissolved. Imagine crystallizing a salt from a solution and removing the crystals. Some salt will remain in the resultant solution.

Why not boil down some acid to concentrate it? With volumes this small, it shouldn't be too hard.

The_Davster - 9-7-2006 at 11:33

I was thinking the other day while distilling bromine, that if the exit of the condensor was run through a tube full of calcium chloride, that the bromine would come out much dryer. Mainly for convenience as it would reduce the handling of the bromine.

I did the pour off ice thing yesterday. I got back around half of the bromine I started with. My bromine may have been damper than yours however.

Magpie - 9-7-2006 at 15:05

I am planing on preparing some con. sulfuric acid from my Rooto. Why should I be the only one who hasn't yet had the fun of doing this.

I don't think I have too big a fraction of ice in my bromine. But I may be surprised. I'll find out when I pour it into my con. acid.

Zinc - 12-7-2006 at 12:25

Today I made some bromine. To 6% H2O2 I added some KBr and diluted sulfuric acid (battery acid). The solution first turned green yellow than yellow than yellow orange orange than orange than orange red and now it is still orange red. During the evolution of bromine the solution heated up. I don't know how much reactants I used since I didn't measure anything.

Zinc - 13-7-2006 at 02:33

Today the solution is yellow. Probbably most of the bromine has evaporated.

Zinc - 13-7-2006 at 09:14

Today I made bromine again. I made a saturated solution of KBr ant to it I added some 12.5% Sodium hypochlorite solution (bleach). Then the solution first turned green because the bleach was green. A few minutes after that the solution turned from green to yellow and the smell of bromine appeared. Then to it I added 19% HCl. Instantly the solution turned from yellow to red. 2-3 minutes after adding the HCl I noticed some bromine vapour above the liquid.

Can this mixture be keept or does the bromine ned to be distilled imidiently out of the mixture? If it can be keept how long can it be keept?(1 year). I have to keep it because I do not have a distillation apartus right now and I t think I will get one in the near future.

[Edited on 13-7-2006 by Zinc]

hodges - 13-7-2006 at 14:28

It can be kept if you can keep the bromine from evaporating. Bromine evaporates much more easily than water. The trick would be finding something to keep it in. Bromine attacks plastic and metals. You would need a glass bottle - not sure what you could use for a cap. A plastic cap might work for a few days.

Hodges

garage chemist - 13-7-2006 at 16:42

Bromine slowly reacts with water. Bromine water will slowly decompose into HBr solution and oxygen. I had some concentrated Br2 solution in a tightly closed bottle (teflon cap) and after a few weeks the color of the solution had lightened up very noticeably.
The decomposition can be slowed down by protecting the solution from light, as light speeds up the reaction of bromine with water.

To store your bromine, you have to isolate it by distillation, dry it by shaking with concentrated H2SO4 and redistill in a dried distillation setup.

Magpie - 13-7-2006 at 16:55

Also see:

http://www.sciencemadness.org/talk/viewthread.php?tid=5572#p...

The_Davster - 31-8-2006 at 19:15

I was making a few mL's of bromine today in the garage, not distilling it, just sucking a few drops off the bottom of the reaction vessel. I wore a gasmask and opened the big garage door when the fumes built up too much. Could smell the bromine on my clothes after. According to someone else, there was a distinct 'chlorine like' smell in the garage, and I don't notice a thing. Later in the day I was using dilute ammonia when I realized 'oh shit, I can't even smell it'.
Halogens don't seem too good for ones sense of smell. I cannot even smell iodine anymore.

[Edited on 1-9-2006 by rogue chemist]

woelen - 1-9-2006 at 01:09

It is temporary. Your sense of smell will return, I once had this while working with Cl2. It did not really irritate me (concentration was low), but the smell for a somewhat longer time apparently did suppress my sense of smell. Since then I have become more careful.

Jdurg - 4-9-2006 at 07:15

Heh. Another thing to keep in mind. Even if you seal bromine into a glass ampoule, over time it will eventually eat through the weak point and start to leach out. I had my bromine in a sealed glass ampoule and after about two years it has started to leach out. I can tell because the smell of it is quite strong in the cabinet I have it stored in.

1-Bromo-3-Chloro-5,5-dimethylhydantoin

Al Koholic - 30-10-2006 at 10:24

This substance is easily found as "brominating tablets" and is used typically in pools or spas as a source of HOBr for decontamination.

I've been thinking of ways to utilize this as a bromine source and have done a few test tube scale reactions to show the liberation of Br is feasible.

The compound is only sparingly soluble in water but when placed in a test tube and boiled, a chlorine like odor is first noted (hypohalous acids) with much solid remaining undissolved. When NaNO2 is added to this solution, there is first the evolution of red gas which quickly dissolves in the water and creates a red solution. After a minute or so, the solution becomes clear with the dissolution of the initial undissolved substance (which releases colorless gas as it dissolves). Eventually a clear solution is obtained that shows slow evolution of colorless gas.

I am curious as to what is happening here. The in situ formation of ClNO2 and BrNO2 is possible with their quick reaction with excess NO2- and formation of transient N2O4.

Is is possible there is diazotization occuring on the nitrogens of dehalogenated 5,5-dimethylhydantoin? Could the colorless gas be N2 arising from the decomposition of this diazo compound? If so what would be the product?

1-Bromo-3-Chloro-5,5-dimethylhydantoin has potential as an interesting source of bromine IMHO.

I guess I'm just not certain what happens with hypohalous acids and nitrite ions in aqueous media.

Nicodem - 30-10-2006 at 11:03

You would need to reduce 1-Bromo-3-Chloro-5,5-dimethylhydantoin with KBr, NaBr or other bromides in the presence of an acid in order to efficiently use it as a bromine source:

1-Bromo-3-Chloro-5,5-dimethylhydantoin(s) + 3KBr(aq) + H2SO4(aq) => 2Br2(l) + K2SO4(aq) + KCl(aq) + 5,5-dimethylhydantoin(s)

Using this exact stoichiometry is very important else you get other pathways reduce the yield of Br2.
If you want only the Br in 1-Bromo-3-Chloro-5,5-dimethylhydantoin you could reduce it with other reductants instead of bromides, but then you would obtain only half mol Br2 per one mol 1-Bromo-3-Chloro-5,5-dimethylhydantoin. For example, like you did with NaNO2, forming NaCl, NaNO3 and Br2. Sulfites and similar reducing agents would also work. However, like I said, it would be a waste of the oxidative power to do so while the reagent can be used to oxidize bromides instead.

Anyway, this has been discussed ad nauseam already. There was no need to open a new thread. UTFSE!

Mason_Grand_ANNdrews - 11-11-2006 at 08:50

I would think when bromine is chilled it can be stored in a glass flask. I have made a small HTML to the bromine and
chlorine synthesis. I hope it looks not to simple.
The HTML contains a chapther to the chlorine production per
electrolysis. I will not open e extra topic because of the
document. Somewhat a probelm is to prepare the layer in the
electrolysis cell. Have someone a idea to design the layer ? I
would guess you can produce bromine with this method. (NaBr in 5% H2SO4)

[Edited on 11-11-2006 by Mason_Grand_ANNdrews]

Attachment: Cl2-Br2.zip (571kB)
This file has been downloaded 854 times


gil - 22-11-2006 at 19:54

hi. I haven't got a clue 'bout chemistry, but I love this site! As I like to expand my knovledge radially, as far as I can get. I can grasp most "basic" concept, due to my backgroun, elsewhere I'm learning everyday STOP
Just for the press, sodium bromide was used MANY years ago as mild sedative/anaphrodisiac to prevent trooper gettin' crazy for sex abstinence! It was served insice milk/caffe/tea for breakfast.

daryl - 29-11-2006 at 21:58

I used to react 2KBr + 2HNO3 -> 2KNO3 + Br2

I forget where the hydrogen went, probably excess water?

This was done in a glass retort which was heated in boiling water.

The distilled Br2 was collected in a test tube sitting in ice.

This was very effective. But you must empty out the potassium nitrate solution before it cools down as the KNO3 will crystalize in the retort.

On occasions, I noticed that when the KNO3 solution cooled, the solution crystalized and there was very little water left.

This is also a handy source of KNO3!

daryl

The_Davster - 29-11-2006 at 22:10

In order for you to be oxidizing bromide to bromine the nitric has to be getting reduced. NO/NO2 was likely the result, completly unnoticable amidst the bromine vapours. Your bromine distillate may have contained an ammount of nitrosyl bromide as a result of this.

woelen - 30-11-2006 at 01:09

Yes, I have done this experiment also. In fact, I found it to be unsuitable for making bromine of accetably purity. The bromine is increadibly impure and difficult to clean. I think that you obtain a mix of Br2, ONBr, NO2 and HNO3. When you add a drop of that bromine to water, it makes a very acid solution, while pure bromine should make an almost neutral solution.

Cleaning the Br2 is hard, because the impurities also are dark brown in the gas phase, and their boiling points also are fairly close (IIRC NO2: 26 C, Br2: 57 C). Repeated fractional distillation may give you purer Br2, but then there are easier ways to obtain relatively pure Br2.

Fleaker - 30-11-2006 at 10:36

I have probably mentioned this earlier in this very thread, but I made it via oxidation by peroxysulfuric acid (caro's acid). The reaction is exothermic, enough so that the bromine readily distills over. An excess of sulfuric acid prevents much water being carried over, but the resultant bromine must still be shaken in a separatory funnel with concentrated sulfuric to rid it of residual water (as well as separate any HBr)

[Edited on 30-11-2006 by Fleaker]

12AX7 - 30-11-2006 at 11:47

Quote:
Originally posted by woelen
When you add a drop of that bromine to water, it makes a very acid solution, while pure bromine should make an almost neutral solution.


Would washing it be any faster than fractional distillation, then? Water has a good affinity for HBr and HNO3, while the Br2 would tend to oxidize the other bits (NO2, ONBr?), or they would oxidize HBr... The result being some disproportion between Br2 <--> HBr and lower NOX <--> HNO3, if any of them be particularly favorable...

Tim

your friend electrolysis

roamingnome - 30-11-2006 at 12:09

http://www.crscientific.com/article-bromine.html


it slides right off the graphite!!!!

verode - 1-1-2007 at 09:12

Quote:
Originally posted by daryl
I used to react 2KBr + 2HNO3 -> 2KNO3 + Br2

I forget where the hydrogen went, probably excess water?

This was done in a glass retort which was heated in boiling water.

The distilled Br2 was collected in a test tube sitting in ice.

This was very effective. But you must empty out the potassium nitrate solution before it cools down as the KNO3 will crystalize in the retort.

On occasions, I noticed that when the KNO3 solution cooled, the solution crystalized and there was very little water left.

This is also a handy source of KNO3!

daryl

you can make it with cold conc. HNO3 just add some NaNO2
and reaction starts at once even at 0ºC you can get bromine from the bottom of flask

Filemon - 4-1-2007 at 15:52

Quote:
Originally posted by evil_lurker
BCDMH is definately an interesting chemical and a darn hard one to figure out how to reduce.

The only literature on it, is what is provided by the pool and spa industry which is generally written for "those not skilled in the art".

The only practical way I can think of reducing BCDMH is by adding NaOH and heating then reducing the formed halogen compounds by adding sulfuric acid... which may or may not work.

When BCDMH is placed into water, it dissolves into hypohalous acids and hypohalic ions whose equilibrium is based upon temperature.

The addition of NaOH neutralizes the acids formed and turns them into hypohalites.

Adding heat causes the BCDMH to dissolve more readily and the hypohalite salts to undergo an auto-oxidation-decomposition into halides and halates.

The reaction would be:

BCDMH + 2 NaOH = 1 NaBr + 1 NaCl + 1 NaBrO + 1 NaClO

1 NaBrO + 1 NaClO ---- heat ---> 2 NaBr + 1 NaBrO3 + 2 NaCl + 1 NaClO3

(X represents a halogen)

6NaX + 2 NaXO3 + 3 H2SO4 = 4Br

I know the equation is unbalanced but it should liberate most of the available bromine due to the chlorine liberated at the same time as the bromine... the free chlorine would undergoes substitution reactions with the bromo compounds.

It is unknown whether or not the NaOH or Sulfuric acid would react with the spent DMH.


Would a simple substitution reaction work?

Cl2 + NaBrO3 => Br2 + NaClO3

I believe that NaBrO3 react with the HCl. That reaction takes place? I don't believe that it produces this reaction because it detach a gas:

NaBrO3 + HCl => NaCl + HBrO3

[Edited on 4-1-2007 by Filemon]

woelen - 5-1-2007 at 01:40

Cl2 + NaBrO3 => Br2 + NaClO3
This reaction does not occur. Chorine replaces bromine from bromide, but not from bromate, because bromate is a stronger oxidizer than chlorine.

NaBrO3 + HCl => NaCl + HBrO3
This also does not occur. HCl reacts with bromate, forming chlorine gas, bromine chloride and water. The reverse also happens, an equilibrium is reached with bromate, bromine, chlorine and water. In fact, the precise conditions and reaction are fairly complex.

PainKilla - 5-1-2007 at 12:52

Trichloroisocyanuric acid is a most excellent source of chlorine, and I found that this process works quite well for generating bromine: (I kept in moles for easier adjusting of ratios etc)

.5 mol C3Cl3N3O3 (crushed up very finely)
2.75 mol NaBr in ~300ml H2O
~275ml 20% HCl

In a 1L round bottom flask set up for distillation and equipped with an addition funnel containing the HCl, add NaBr solution. Add TCCA to this solution. With stirring, begin to heat the solution (to ~55C) while very slowly adding the HCl. The bromine will begin to gradually distill over. Continue addition of HCl and distill until no more bromine is carried over.

Yields should be near quantitive (~1.25mol Br2)


The rate of bromine generation depends on HCl addition, so if your condensor is less than adequate, adding the HCl slowly works well as compensation. The procedure works very well, some other things to note are that the TCCA must be finely crushed, otherwise it takes a long time to react and results in a buildup of Br2. It's also quite a bit easier to stir when fine. It is also very important to have good stirring because without it, the bromine forms a layer in the reaction flask and once heat is applied, you begin to have an extreme amount of bromine evaporating that the condensor can't handle (well mine can't anyway). Other than that, this is a very straightforward procedure.

Some things I have yet to test is adding H2SO4 to the receiving flask in order to dry the bromine and keep it at least somewhat less volatile. I should be retrying the procedure soon and will report with more accurate measures.

EDIT: Forgot to put in yield. EDIT2: Fixed error.



[Edited on 7-1-2007 by PainKilla]

JohnWW - 5-1-2007 at 15:59

I wonder how pure that Br2 would be, though. As well as oxidizing Br- to Br2, the Cl2 could also result in the formation of BrCl, and oxidize further to BrCl3 or the BrCl4- anion.

Tsjerk - 5-1-2007 at 16:18

And what would be the solubility of chlorine in bromine? I can imagine that some chlorine would bubble out and dissolve in the condensor.

Nicodem - 7-1-2007 at 03:56

Quote:
Originally posted by PainKilla
.5 mol C3Cl3N3O3 (crushed up very finely)
1.25 mol NaBr in ~300ml H2O
~275ml 20% HCl


What is the point in using such an enormous excess of trichloroisocyanuric acid (TCCA)?

The redox reaction for the oxidation of NaBr with TCCA is:
TCCA + 6NaBr + 3H(+) => cyanuric acid + 3Br2 + 3Cl(-) + 6Na(+)

So you used more than a double excess of TCCA which quite possibly caused the formation of lots of BrCl and Cl2. I recommend using a 5 to 10% excess of NaBr instead. This would efficiently prevent any interhalogen compounds distilling over. There is no need to use so much HCl as the acid for the reaction. Use just a bit more than the required amount (~0.6 equivalent in relation to NaBr instead of ~1.65 that you used). Using a non volatile acid like H2SO4 would also be preferred to avoid HCl contamination of the distillate. Mind that only water is easily removed from the Br2 distillate with a conc. H2SO4 wash, but removing BrCl, Cl2, HCl and such is not easy (so preventing contamination is wiser than repairing what is already done).
Also, as I already said several posts above, using 1-bromo-3-chloro-5,5-dimethylhydantoin as the oxidant is more rational as you get to use the Br equivalent from it as well (and not only from NaBr). Otherwise such methods have no advantage over simply using H2O2/H2SO4 as the oxidant (except maybe if H2O2 is getting difficult to acquire).

woelen - 7-1-2007 at 07:25

I actually like the method of Painkilla (with your modification of using a little excess NaBr), because TCCA is very cheap and easy to obtain. Over here in Europe, 1-bromo-3-chloro-5,5-dimethylhydantoin is not easy to obtain, swimming pools never are brominated over here (in NL it even is forbidden).

Filemon - 7-1-2007 at 10:43

Quote:
Originally posted by PainKilla
Trichloroisocyanuric acid is a most excellent source of chlorine, and I found that this process works quite well for generating bromine: (I kept in moles for easier adjusting of ratios etc)

.5 mol C3Cl3N3O3 (crushed up very finely)
1.25 mol NaBr in ~300ml H2O
~275ml 20% HCl

In a 1L round bottom flask set up for distillation and equipped with an addition funnel containing the HCl, add NaBr solution. Add TCCA to this solution. With stirring, begin to heat the solution (to ~55C) while very slowly adding the HCl. The bromine will begin to gradually distill over. Continue addition of HCl and distill until no more bromine is carried over.

Yields should be near quantitive (~1.25mol Br2)


The rate of bromine generation depends on HCl addition, so if your condensor is less than adequate, adding the HCl slowly works well as compensation. The procedure works very well, some other things to note are that the TCCA must be finely crushed, otherwise it takes a long time to react and results in a buildup of Br2. It's also quite a bit easier to stir when fine. It is also very important to have good stirring because without it, the bromine forms a layer in the reaction flask and once heat is applied, you begin to have an extreme amount of bromine evaporating that the condensor can't handle (well mine can't anyway). Other than that, this is a very straightforward procedure.

Some things I have yet to test is adding H2SO4 to the receiving flask in order to dry the bromine and keep it at least somewhat less volatile. I should be retrying the procedure soon and will report with more accurate measures.

EDIT: Forgot to put in yield.


[Edited on 5-1-2007 by PainKilla]


As much as time takes the reaction?

PainKilla - 7-1-2007 at 10:55

Sorry everyone, Nicodem is right, I actually messed up in writing out the procedure. The proper amounts should be a slight excess of TCCA, but only by .25mol less of NaBr. In other words, 2.75mol NaBr. I use the slight excess of TCCA because I want to ensure that all of the NaBr is used up (harder for me to get NaBr than TCCA).

As far as contamination goes, isn't BrCl unstable at those temperatures? All of the other contamination is solved without problem by redistilling the bromine, and I have noticed no problems in using the bromine for synthesis.

It may be true that other methods aren't any more or less effective but this is a very quick and easy procedure to do and generates large amounts of bromine without any sort of effort, and for those in the US, the chemicals are fairly easy to obtain (and cheap).



[Edited on 7-1-2007 by PainKilla]

woelen - 7-1-2007 at 13:28

Using excess NaBr leads to double loss of bromine. The excess bromide is not reacted, but the excess bromide also keeps behind quite some bromine, due to formation of the complex ion [Br2.Br](-), the so-called tribromide ion. Bromine dissolves MUCH better in solutions of a bromide, than in plain water.

On the other hand, BrCl is not that unstable. If it is formed, then it certainly will be in the bromine, it simply is dissolved in the bromine. It will be hard to remove, once you have it in the bromine.

So, whether you take excess bromide or excess TCCA depends on what you want to do with the bromine. In the past I have made bromine several times, but I always used it for somewhat more spectacular experiments, such as reacting it with Al, Mg, Na and red P. For all those experiments, the presence of some BrCl is not an issue at all.

But if you want to use the Br2 for synthetic purposes (e.g. brominating some organic), then you do not want the BrCl, because that always will lead to very hard to remove impurities of chlorinated compounds. BrCl chemically speaking resembles Br2 very much, but it is somewhat more reactive.

apidej - 7-1-2007 at 18:54

good

S.C. Wack - 7-1-2007 at 20:11

Since I have an old photocopy of a relevant 1 page article - "Methods for preparing aqueous solutions of chlorine and bromine for halogen displacement reactions", where a saturated solution of bromine is made from TCCA in that way, among other things - here it is, even though I think that most will pass on this as a preparative method.

Attachment: jce_64_156_1987.pdf (58kB)
This file has been downloaded 1769 times


YT2095 - 11-1-2007 at 06:16

I just followed a synth Woelen gave me, using KBr + KBrO3 + HCl as a "one pot" reaction, I had an instant reaction and bromine was created, although it wasn`t until it spent a short time in the fridge cooling that I actualy got a blob of Br2 at the bottom (about match head sized).
I`ve had a read through this thread and distilation use etc... and also Bromine water.
I decided to make Bromine water with the gas given off, so using a double boiler I heated the rxn vessel and caught the gas under water, naturaly it changed color, but then eventualy droplets of Br2 were falling to the bottom, cut a long story short I now have a 4mm blob of Br2 in the bottom of a test tube with about 2cm of bromine water over it.
also I did the RP rxn with the spent vessel as there was still a little Br2 gas in there, RP instantly ignites on contact with it, the same as with Cl2 gas.
it`s also a nice way to clean your vessels afterwards :)

Thnx for the synth woelen and the extra data I needed from the rest that have posted in this thread, on behalf of myself and my Element collection ;)

garage chemist - 11-1-2007 at 06:26

A slight excess of NaBr in the TCCA reaction does of course keep some of the bromine in solution as the tribromide, but this does not matter if the bromine is distilled out of the mix, as one must always do if any useful yield of bromine is to be achieved.
Extracting the blob of liquid bromine that frequently forms under the mix and discarding the aqueous layer is essentially the same as dumping half of your NaBr down the drain.
Bromine does dissolve in water to a degree that cannot be neglected, even if bromide ions are absent. Distillation is always necessary for any serious preparation of bromine.

However, maybe it isn't that bad an idea to not use an excess of NaBr and get BrCl- containing bromine as the first product.
BrCl and dissolved chlorine in the crude bromine can be removed completely by dissolving it in NaBr solution and distilling the bromine out of it. Cl and BrCl react with NaBr to form additional bromine.
The NaBr solution can be used again and again for this purpose as long as it still contains bromide (test with HCl + H2O2).
This would give an overall better yield of bromine since all of the bromine in the NaBr gets extracted in the first step.

[Edited on 11-1-2007 by garage chemist]

YT2095 - 11-1-2007 at 11:56

I`ve just repeated this synth again, scaled up X10 and now hav a good 6mm of the test tube bottom full of Br2, however, I don`t know what`s happened but it would seem you cannot simple resume where you left off (so to speak), plenty bromone was "spilling" over the sides of the receiver tubs and into the conical flask holding the ice water, much went into the air too and I was forced to put my gas mask on!:o

I have no idea why it went this way at all??? everything was exactly the same as before other than being scaled up, the only Logical thing I can think of is that it overwhellmed the absorbtion capacity (surface area) of the liquid as it was being produced too fast.

all`s good in the Lab now, just my hair and clothes stink of Br2 now! my wife is just gunna LOVE ME in bed 2night! :(

woelen - 11-1-2007 at 12:05

Scaling up of an exothermic reaction usually leads to different results. Speed of transfer of heat grows quadratically (area), while production of heat grows cubically (volume). So, the same reaction will lead to higher internal temperatures in the mix, so I certainly can imagine that you obtained copious amounts of bromine vapor from the warmer mix.

Filemon - 14-1-2007 at 04:38

That solubility has in water ClBr?

woelen - 14-1-2007 at 09:36

Similar to that of bromine. It also reacts with water, leading to formation of bromic acid, bromine, and chlorine. There is a fairly complicated equilibrium, BrCl when added to water leads to a complicated mix of different chemicals.

indigofuzzy - 15-1-2007 at 01:26

Ok, 7 pages is a lot to sift through, but has electrolysis been brought up?

I know that when electrolyzing a solution of NaCl, H2 forms at the anode, and Cl2 at the cathode, resulting in NaOH in solution. Would the same work with NaBr? Would you get H2 at the anode and Br2 at the cathode? If so, it sounds like a much less messy, and vastly simpler method.

woelen - 15-1-2007 at 04:28

H2 forms at the cathode, not at the anode. Br2 forms at the anode.
But isolating the Br2 from the solution is a real pain. For a home-lab, electrolysis definitely is not the best way to make bromine.

You can use electrolysis indirectly, first make KBrO3 by means of electrolysis (I made another topic on that subject) and then react the KBrO3 with KBr and concentrated acid. That is a very convenient method of making Br2.

Filemon - 15-1-2007 at 09:17

what reaction does it produce: NaBrO3 + Cl2 =>?

YT2095 - 15-1-2007 at 09:46

I expect you`ll get plenty side reactions doing it that way, leaving a messy outcome with a shopping list of products.
don`t forget you`ll need these in a state where they can combine, as Cl2 would be gas the other must be a liquid, and since water would be choice, Messy:(

[Edited on 15-1-2007 by YT2095]

woelen - 15-1-2007 at 11:09

Quote:
Originally posted by Filemon
what reaction does it produce: NaBrO3 + Cl2 =>?

No reaction: bromate does not react with chlorine, not in the solid state, nor when both are dissolved in water.

[Edited on 15-1-07 by woelen]

YT2095 - 15-1-2007 at 11:22

or of course it do nothing at all like Woelen said :P

Nerro - 15-1-2007 at 11:49

It's not likely that chlorine would donate electrons to the Br in the bromate.

Which makes me wonder,

how well would this work with I2?

woelen - 16-1-2007 at 05:19

Bromate can oxidize I2. I2 and bromate can react, giving iodate and Br2. I do not expect this to work well in the solid state, but in solution, certainly when some acid is present I'm quite sure the following happens:

I2 + 2HBrO3 ---> Br2 + 2HIO3

I did not try with my bromate, but what I did do is reacting potassium chlorate in acidic solution with I2. If you do that, then the I2 quickly dissolves and the liquid becomes light green. Chlorine gas is evolved.

[Edited on 16-1-07 by woelen]

Antwain - 6-10-2007 at 09:24

Interesting thread. I have 1kg of bromochlorodimethlyhydantoin and want to make bromine from this.

The way I see it my options are throw in NaOH and boil it (slow addition may be better, not sure yet) to make NaCl:NaClO3:NaBr:NaBrO3 = 2:1:2:1. Then slowly add HCl to the solution produced (minus the dimethylhydantoin, solubility 0.1g/100mL) while heating, to produce Br2 with perhaps some BrCl. Or I could just add HCl to the BCDMH and get the bromine and halogens from that. If bromine is liberated first then the amount of HCl added could be tailored so as to not liberate large quantities of chlorine.

Q1. Will concentrated, acidic halogens react with dimethylhydantoin to decrease my yield? ie. do I need to do an alkaline extraction first?

Q2. If the addition of HCl is slow will I be getting *mostly* Br2, BrCl, Cl2 or all of the above? If all of the above, then how the heck do I handle 100g+ of chlorine :(

Q3. Since at least some BrCl will be formed how do I remove this from the Br2? I was thinking that by refluxing it up a column for ages I could let it escape as Cl2, but using tap water to cool the column BrCl will probably be lost too (bp.'s of Cl2, BrCl and Br2 are -34*C, 5*C and 59*C respectively). I am assuming that an equilibrium exists Cl2 + Br2 <---> 2BrCl.

Any help appreciated. ;)

[Edited on 7-10-2007 by Antwain]

Nicodem - 6-10-2007 at 10:16

Quote:
Originally posted by Antwain
Interesting thread. I have 1kg of bromochlorodimethlyhydantoin and want to make bromine from this.

The way I see it my options are throw in NaOH and boil it (slow addition may be better, not sure yet) to make NaCl:NaClO3:NaBr:NaBrO3 = 2:1:2:1. Then slowly add HCl to the solution produced (minus the dimethylhydantoin, solubility 0.1g/100mL) while heating, to produce Br2 with perhaps some BrCl. Or I could just add HCl to the BCDMH and get the bromine and halogens from that. If bromine is liberated first then the amount of HCl added could be tailored so as to not liberate large quantities of chlorine.

Q1. Will concentrated, acidic halogens react with dimethylhydantoin to decrease my yield? ie. do I need to do an alkaline extraction first?

Q2. If the addition of HCl is slow will I be getting *mostly* Br2, BrCl, Cl2 or all of the above? If all of the above, then how the heck do I handle 100g+ of chlorine :(

Q3. Since at least some BrCl will be formed how do I remove this from the Br2? I was thinking that by refluxing it up a column for ages I could let it escape as Cl2, but using tap water to cool the column BrCl will probably be lost too (bp.'s of Cl2, BrCl and Br2 are -34*C, 5*C and 59*C respectively). I am assuming that an equilibrium exists Cl2 + Br2 <---> 2BrCl.

Any help appreciated. ;)

[Edited on 7-10-2007 by Antwain]

If you would have read the last two or three pages of this thread you would have seen all your questions were already answered.
According to the redox equation you need exactly 3/2 mol Na2SO3 for each mol of 1-bromo-3-chloro-5,5-dimethylhydantoin (in acidic media). But like it was already explained, reducing 1-bromo-3-chloro-5,5-dimethylhydantoin with anything but KBr (or some other bromide) would be quite irrational as a source of Br2.
Your idea of first "hydrolyzing" the bromochlorodimethylhydantoin by halate disproportionation in basic media followed by another disproportionation in acidic media is a bit irrational to say the least. I see no rationale on why to do that. Not to mention that the first step would be terribly slow since the bromochlorodimethylhydantoin with hydroxide is reversible with the equilibrium laying far to the left (thus the concentration of hypobromite and hypochlorite would be very low at any given time).

garage chemist - 6-10-2007 at 10:19

Antwain, thats not the way to go about it. You can't separate bromine and chlorine by physical means. Use chemistry.

You can, for example, make use of the fact that H2O2 will liberate bromine from HBr, but not chlorine from HCl. So boil your BCDMH with NaOH, acidify slightly, reduce both chlorate and bromate with SO2 (from K-metabisulfite), acidify strongly and add H2O2. Then only bromine will be set free. Distill off. Free from chlorine traces by stirring with NaBr solution and distilling from it.

Nicodem - 6-10-2007 at 10:45

But why? It simply does not make any sense when you can simply reduce BCDMH directly to Br2 and chloride ions. What role would boiling it with NaOH have?

Antwain - 6-10-2007 at 10:58

@ garage chemist- thanks, that seems like an acceptable way of doing things. Would you expect that 6% H2O2 would be up to the task or do I need to find some 30%? Alternately, if i was to reduce some of it and then throw the mixture of MBr and MCl back in with BCDMH in the right proportions and acidify with H2SO4 will this give me chlorine-free bromine? Since I believe that the answer to that is no, would the mixture that came across be able to be reacted with some more mixture of alkali metal halides to produce only bromine?

@nicodem - I said help would be appreciated, so you can assume that your sarcasm is unappreciated. If MY questions had been answered then I would not have asked them. That is all.

edit- a hydrolysis is probably not necessary in the case of a subsequent reduction. I only have a few grams of bromide and cannot easily get more here. I was exploring other options, which have clearly lead to dead ends.

[Edited on 7-10-2007 by Antwain]

Nicodem - 6-10-2007 at 11:03

Sarcasm? WTF? Where? :mad:
I answered your questions and even told you where you can find more detailed information. I even took time to calculate the amount of bisulfite needed. For this I get to be accused of sarcasm? You really should think more about your ignorance and its origins.

Yes, your proposal for bromine generation is a great idea! (<- that is sarcasm from my side)

Antwain - 6-10-2007 at 11:19

You are right, it is not sarcasm exactly, just general not-niceness. Whilst you clearly know what you are talking about, your post(s) have not helped me at all towards my goal. Did you think that maybe if I am not going to use the bromide + BCDMH method which HAS been used successfully then maybe I had a good reason, such as having no bromide. If you were trying to suggest to me that I could make bromide (which I could then oxidise with BCDMH) by reducing BCDMH then you failed completely to communicate that idea in your first post. If that is not what you were suggesting then I still don't know what you meant. It is one thing to tell someone that they are wrong, it is another to do so in a condescending and confusing way.

[Edited on 7-10-2007 by Antwain]

Nicodem - 6-10-2007 at 11:52

If you didn't understood then say so. I will gladly explain if I can. But accusing me of sarcasm just because I tried to help you with a short post instead of explaining basic redox chemistry in a lengthy post is not being nice from your side. Moreover, I directed you to other posts where you could have found a lengthier explanation if you needed it!

Anyway, what I was saying in my reply is that if you do not want to use KBr as the reducing agent you can use sodium or potassium sulfite or bisulfite in acidic media instead. This can come handy in cases where you can not obtain KBr (like in your case) but is not so rational chem-wise since you would end up with less Br2. Another problem with the (bi)sulfite is that you need to use the exact amount needed to reduce BCDMH to chloride and Br2. Using too much would result in reduced yields since Br2 gets reduced by SO2 as well. You would not encounter this problem if using KBr since even it used in excess it can not reduce Br2, but utmost dissolve in the form of the unstable Br3(-) complex anions.

garage chemist - 6-10-2007 at 15:56

Nicodem is right with the equilibrium BCDMH + 2 NaOH <---> NaOBr + NaOCl + DMH being far on the left side. That would speak against the NaOH treatment.
So direct reduction with 3/2 mol Na2SO3 per mol BCDMH would be much better, followed by acidification with HCl. If the amounts were right, only bromine will be liberated... it would be better to use some excess of Na2SO3 and then oxidise the leftover bromide ions to Br2 with H2O2 after acidification.

[Edited on 7-10-2007 by garage chemist]

Sauron - 6-10-2007 at 20:35

If I did the sums right, a Kg of the hydantoin is about 4 moles. So you ought to get 2 mols of Br2 (four gram-atoms Br) out, or about 320 g. That is only about 100 ml, it is dense stuff.

Don't throw away the dimethylhydantoin, it is easily regenerated to the haloamine.

Antwain - 7-10-2007 at 01:32

Those are the values I got too. I used Na2S2O5 as the reducing agent.

I am not sure that I will ever attempt this again. It was a complete failure (except that I got bromine). My lab book was destroyed in the chaos, however I believe I used 254g of metabisulfite and 400 and something g of BCDMH. Either way it was correct to within 0.1g, 4:3 molar ratio.

I added some water to the 1L RBF connected to a 24/40 water condenser with a good flow and a 250mL RBF in an ice bath. The reactants were already in it and had changed colour slightly to a slight red/orange.

The reaction started slowly and vapor was coming out the pressure vent, not much but it smelled. Then liquid started to condense in the condenser then run into the flask. By now much vapor was coming out the vent so I figured that lightly placing something over it would help. It was not SEALED, but still it blew out a plug on the 1L flask. Despite being emersed in a water bath it continued to react faster and faster. Somehow something blew out (I don;t know what because I was unable to enter the shed by this point) but somehow bromine and or solution got into an oil bath. White fumes went everywhere. Everything turned from bad to shit in an instant. I took a deep breath and grabbed some sodium sulfite that was very fortunately at the very front of my cupboard (it was so white you couldn't see and it was a lachrymater). I turned the fuse off at the box and started spraying torrents of water blindly into my shed then went in to get it to spray into the flask (better dilute on the walls than still reacting). This is when I got a lungfull of something nasty. I threw the sulfite solution into the bromine which had distilled (it was full of water which had blown through the condenser, but there was 20-40mL there even though I had added no acid). For the next half hour I was seriously contemplating going to the hospital but decided that after half an hour the damage had been done. It probably isnt permanent but I still have a slight wheeze :(

I don't know why it reacted so quickly and violently. Thank god my neighbors were still in bed cos you would have noticed this smell 100m downwind at its peak. I have not been so scared in many years, especially when all hell had broken loose and I knew that the reaction was <10% complete and when initially I was panicking and couldn't breathe at all. Thank god I had the sulfite at the front of the shelf, if it had been at the back things would have been much worse.

I am at a loss as to the speed of the reaction. Earlier in this thread it was claimed that the reaction of BCDMH with bromide was not exothermic, so I don't necessarily see why this should be. Maybe oxidising the metabisulfite made bisulfate, that would explain it (duh). In hindsight, adding the metabisulfite in the form of solution, slowly would have been much smarter.

To be honest the main reason I am posting this humiliation is so I can tell you that if you are a n00b or a kewl or generally someone who knows nothing about chemistry, then don't try this at home YOU WILL PROBABLY DIE!

IF, and thats a big IF, I ever try this again it will be scaled down several fold and there will be dry ice. I think I am going to take a break from practical chemistry after this, at least until uni is over. Except tomorrow when I will have to spend 3 or so hours cleaning out my waterlogged shed, waterlogged heating mantle, waterlogged graphics calculator. :(

woelen - 7-10-2007 at 09:58

Antwain, sad to read about this (near) accident. Most important is that you personally have not been seriously injured. Please don't throw the towel in the ring, I like your contributions and the things you manage to do, even in the difficult social climate of Australia.

Consider this event as a wise lesson. The goddess of chemistry is beautiful, but she must be handled with care and don't make her pissed off ;).

Nicodem - 7-10-2007 at 12:22

Antwain, I'm sorry to hear about the disaster and wish you had no permanent lung damage or other consequences.
However, the kind of accident you had is a perfect example of malpractice. You actually did everything wrong in respect of preventing such a disaster. So for the newbies who might do experiments with powerful oxidants I though to emphasize what is a definitive not to do:

1.) Never ever do exothermic reactions before testing them on a 10 mmol scale! When scaling up always consider that exothermic reactions do not scale up smoothly at all.

2.) Never ever put the reactants together and then add the solvent. This will always cause a runaway even if the reaction is only slightly exothermic.

3.) At least one of the reagents needs to be in solution and the second one need to be added slowly in small portions/dropwise with temperature monitoring and efficient cooling bath.

4.) If a reaction does not appear to start immediately do not heat or add more catalyst/acid/base or whatever the reaction mechanism requires to speed it up. Have patience.

5.) Always use theory to evaluate the expected amount of heat or gasses evolved in the reaction and think of all possible influences any of the parameters will have on the reaction proceeding and possible side reactions.



Edit: I forgot to mention that since HCl is one of the products the reaction actually needs no added acid to proceed to Br2. It is practically autocatalytic.

[Edited on 7/10/2007 by Nicodem]

Antwain - 8-10-2007 at 08:29

Yes I have embarrassed myself greatly with this one but fortunately after coughing up some questionable stuff my lungs appear to have returned to normal. There was a reason why I chose to do it on this scale, although clearly it was a bad idea. The solubility of Na2S2O5 is not sufficient in water for my to have completed this reaction in my largest vessel otherwise, and I was reluctant to take it apart and recharge it to complete the procedure. I don't, in hindsight, think it was thermal runaway. I believe that it was agitation by gas causing greater mixing leading to greater production of gas.

I will be out of the lab for the next several weeks as I do not have the time anyway, but I will be back after that. Ok, lets suppose I want to do this properly, later. I am now thinking that running metabisulfite into hot but not boiling, wet BCDMH may be bast. The reaction will be quick at that temperature, and the rate can be moderated by addition of reducing agent, with external cooling with a water bath if the mixture heats up or heating with a mantle if it is not exothermic enough. This can then pass through a condenser to a large enough flask immersed in ice water, possibly with a few mL of H2SO4 in it. The gas outlet from the condenser can be put through tubing (which will be attacked but not destroyed) and then into a sintered bubbler ( I need to get one anyway really) which will bubble through sodium sulfite or metabisulfite solution. Do you see a problem with this method nicodem?

Also, Whilst I was able to come up with the ratio BCDMH : Na2S2O5 = 3:4 myself, I couldn't balance the hydrogen ions needed and liberated, and I didn't account for the fact that oxidising Na2S2O5 + H2O + 2[O] ---> 2NaHSO4 which is acidic! can you help me with the FULL balanced equation for this. Or even just tell me that I do not need to add any acid which would be as useful. Thanks. Actually one more thing... If it makes HCl, should I be running a stoichiometric mixture of metabisulfite and NaOH into the solution to prevent this from distilling over as well, or will this also prevent the reaction from working.

Engager - 11-10-2007 at 18:22

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).

Photo with bromine made by this reaction is shown below:



[Edited on 12-10-2007 by Engager]

brom.jpg - 64kB

Antwain - 11-10-2007 at 19:50

Well if I had KBr, and KBrO3 instead of neither, then I am sure that is the method I would use ;)

Siddy - 11-10-2007 at 19:54

A percent yield for that post would be nice.

Also, where can you buy potassium bromate?

haha, guy above beat me.

[Edited on 11-10-2007 by Siddy]

12AX7 - 12-10-2007 at 03:22

"Almost quantitative" = 90% range.

woelen - 12-10-2007 at 05:09

Quote:
Originally posted by Siddy
A percent yield for that post would be nice.

Also, where can you buy potassium bromate?

haha, guy above beat me.

[Edited on 11-10-2007 by Siddy]

Potassium bromate you can make yourself. I posted a thread about this on sciencemadness. For your convenience I will post the method here. For comments and further discussion of this, please search the forums.

http://woelen.scheikunde.net/science/chem/exps/KBrO3_synth/i...

Using KBrO3 and KBr for making Br2 indeed is very nice. This reaction runs very smoothly, it only is slightly exothermic and has very good yield. The only issue is that quite some bromine may remain dissolved in the water, this has to be taken out by distillation.

I expect that engager will have a yield of approximately 75% if he only took the blob of liquid bromine and did not distill the bromine, dissolved in the water.
Solubility of bromine in pure water is appr. 3 gram per 100 ml. When H2SO4 is dissolved, and no excess bromide is present, then it is somewhat less.

[Edited on 12-10-07 by woelen]

Engager - 12-10-2007 at 07:24

Quote:
Originally posted by woelen
I expect that engager will have a yield of approximately 75% if he only took the blob of liquid bromine and did not distill the bromine, dissolved in the water.
[Edited on 12-10-07 by woelen]


Yes, you are correct. Yeild is near to quantative then bromine dissolved in water is recovered. I usualy do this after large ammount of bromine water is accamulated, from several preparations of bromine. By the way the byproduct - bromine water is quite usefull for many experiments, such as preparation of CBr4 from acetone by action of hypobromite (made by dissolving NaOH in ice cold bromine water), or preparation of isocyanogen tetrabromide from sodium 5,5'-azotetrazolate. So, don't waste bromine water it can be quite usefull.

[Edited on 12-10-2007 by Engager]

Nicodem - 13-10-2007 at 02:03

Quote:
Originally posted by Antwain
Also, Whilst I was able to come up with the ratio BCDMH : Na2S2O5 = 3:4 myself, I couldn't balance the hydrogen ions needed and liberated, and I didn't account for the fact that oxidising Na2S2O5 + H2O + 2[O] ---> 2NaHSO4 which is acidic! can you help me with the FULL balanced equation for this. Or even just tell me that I do not need to add any acid which would be as useful. Thanks. Actually one more thing... If it makes HCl, should I be running a stoichiometric mixture of metabisulfite and NaOH into the solution to prevent this from distilling over as well, or will this also prevent the reaction from working.

The reduction should be carried by carefully adding BCDMH in small portions into a stirring, less than 10%, solution of the sulfite (for the exact amounts see the reaction stoichometry) in a flask immersed in an ice bath. Since the reaction media gets more and more acidic it is best to use Na2SO3 instead of bi- or metasulfite to yield the neutral Na2SO4 instead of acidic NaHSO4. This is only important as to prevent too much SO2 escaping from the reaction mixture thus ruining the exact stoichometry. It is therefore advised to add 1/2 of Na2CO3 per every NaHSO3 used if one has no Na2SO3 available. It is also important to add BCDMH to the sulfite as the opposite order gives rise to Cl2 formation (before all BCDMH is reduced to 1-bromo-DMH, there is Cl2 equilibrating in the reaction mixture).
Bellow are the relevant redox reactions (I hope I calculated all the electrons correctly, otherwise someone please correct me).

BCDMH_reductions.gif - 13kB

Antwain - 13-10-2007 at 10:08

@Nicodem- thanks for that. I was concerned about the acid, but it completely escaped me that I would be losing SO2 as well. At 0*C will the bromine have a low enough vapor pressure for this reaction to be done outside a fume hood, of course the BCDMH will have to be added slowly to prevent the temperature rising.

Antwain - 15-10-2007 at 00:38

I couldn't stay away from the lab. I dissolved 94.5g of Na2SO3 in ~400mL of water and cooled in ice water. I added added 121.6g of BCDMH in small portions with stirring. During this time 24 ice cubes were consumed (not a good measurement, but it gives some idea), and a further 12 ice cubes were melted in the following few minutes. At all times there was spare ice in the bath.

I didn't want bromine liquid per se, so a solution of 28.4g of KOH in ~150mL of water was prepared in a Erlenmeyer and a liebig was set up with a small length of disposable plastic tube to dip into the flask. For a long time only vapour came across but when i ramped the temperature up finally bromine started condensing. The tube was attacked and discoloured, but this was not a concern and it seems that the very small amount of crap will be able to be filtered out. Heres the shit part.... Not only did not all of the bromine distill across (most of it came) and the reaction flask is still somewhat red, but also the collection flask had to have extra KOH added (a few g) to stop it from turning polyhalide coloured and smelling bad. I should also point out that after bromine stopped distilling across (at a temp of 58-63*C) I had to remove the thermometer so I could stick a pipette connected to an air pump down and put roughly a bubble a second through the flask to purge the vapour above the solution. This did get much more of the red colour out. So.... Not all the bromine distilled and yet the base was inadequate.... Clearly HCl came across too.

Can someone tell me why my bromine was so reluctant to leave solution, and suggest a modification to fix it next time. Also it may be prudent to collect liquid bromine next time, possibly beneath H2SO4, so that all the HCl buggers off.

kilowatt - 15-10-2007 at 01:09

I too have extracted bromine in dilute form the electrolysis of NaBr in a separatory funnel. I plan to distill it with H2SO4 to remove water and condense pure bromine, though I am not sure of the exact composition of the liquid precipitate I got. I guess it is just dilute bromine/hypobromous acid.

Anyway, more recently I have thought of a method using a rather well known set of reactions to obtain any of the halogens except fluorine in highly purified form with very easy decompositions. By adding sodium bromide to a solution of silver nitrate or sulfate (the latter would be less preferred due to its very low solubility, but it could work especially if you have trouble getting/making nitric acid), one can precipitate very pure silver bromide (or chloride or iodide). As the silver halides are photosensitive and decompose to pure silver and halogen, the halogen can be isolated from the decomposition flask with a simple distillation. The silver powder left behind can be easily acidified and reused as silver is an expensive metal and will likely be in limited quantity

NaBr + AgNO3 --> AgBr + NaNO3
2AgBr -(light)-> 2Ag + Br2

The recycling of the silver and also the nitrate (or sulfate) ion could likely be simplified by using copper bromide as an intermediate, as the copper salts are easily regenerated to acid.

NaBr + H2SO4 --> NaHSO4 + HBr
2HBr + Cu -(H2O2)-> CuBr2 + H2
CuBr2 + 2AgNO3 --> 2AgBr + Cu(NO3)2

2Cu(NO3)2 -(170°C)-> 2CuO + 4NO2 + O2
4NO2 + O2 + 2H2O --> 4HNO3
HNO3 + Ag --> AgNO3

CuO + 2HBr --> CuBr2 + H2O
repeat CuBr2 + 2AgNO3 --> 2AgBr + Cu(NO3)2

Alternatively:

CuBr2 + Ag2SO4 --> 2AgBr + CuSO4

CuSO4 -(650°C)-> CuO + SO3
SO3 + H2O --> H2SO4
H2SO4 + 2Ag --> Ag2SO4

CuO + 2HBr --> CuBr2 + H2O
repeat CuBr2 + Ag2SO4 --> 2AgBr + CuSO4

Nicodem - 15-10-2007 at 07:01

Quote:
Originally posted by Antwain
Can someone tell me why my bromine was so reluctant to leave solution, and suggest a modification to fix it next time. Also it may be prudent to collect liquid bromine next time, possibly beneath H2SO4, so that all the HCl buggers off.

Quite surely some HCl distilled over, but that might not have been the main reason your 28.4g KOH did not suffice to trap all Br2. It appears to me that you have not considered that KOH pellets are only about 85% KOH with the rest being water (the amount you used is actually 0.43 mol and not 0.5 mol as you intended). You would actually need at least 33g KOH for complete neutralization. Also, it is very difficult to distill or extract all bromine from water solutions due to various equilibriums and interactions in the system (bromine disproportionations, bromide/bromine complexations and DMH+Br2 reversibility). There will always remain enough for the reaction mixture to remain bromine-like colored. Just ignore that and judge its presence by the vapors color instead.

not_important - 15-10-2007 at 08:04

The decomposition of silver halides is neither fast nor complete. Plus unless you have the halide as extremely small particles the outer layers will shield the inside of the halide particle from light.

If you have simple halide slats, such as NaBr, to start with, just use the halide salt + MnO2 + H2SO4 distillation to obtain the free halogen. Even low grade MnO2 will work, if it's from old batteries just wash it well with water to remove chlorides. This eliminates the problems with recovering bromine from water solutions.

In the case of organic sources such as spa "bromine" tablets, then an intermediate isolation step is useful.

An very old industrial method used to recover dilute bromine vapours was to absorb in alkali solution, evaporate to dryness and then heat strongly to decompose bromate, then extract with methanol which left most of the chlorides behind. The extraction trick works even better with NaI, for which acetone can be used as a solvent.

To remove chlorine from bromine or iodine, or bromine from iodine, distilling the halogen from some of its sodium or potassium salt will do the job. The salts remaining behind can be repeated used until they are mostly the chloride, at which point the extraction with methanol can be used to recover the last of the bromide or iodide from the chloride.

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