Sciencemadness Discussion Board - Bromine from NaBr

Bromine from NaBr

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longtermmadscientist - 13-5-2012 at 13:13

I've been thinking of a simple way of producing a few ml of elemental bromine using sodium bromide. Could conc. HNO3 be used on NaBr solid or solution? I've been researching up on this, and there's a lot of talk of H2SO4 but it seems it's not a strong enough oxidant to produce Br2 in good yield, and you get a lot of HBr.

My concern is that I will produce a lot of NO2, which, being the same colour as Br2, could lead to some confusion as to how much Br2 has been produced. Also, I'd like to know whether the bromine would contain more that a trace of NO2. Water impurity can be extracted by conc. H2SO4, but what about removing the NO2, if this turns out to be an issue?

simba - 13-5-2012 at 13:41

Why don't you just use sulfuric acid? You said you need only a few ml, so yields are not important right?

elementcollector1 - 13-5-2012 at 13:43

I would recommend 98% H2SO4, added slowly to very cold NaBr solid. If you really wanted to go the extra mile, you could use dry ice so that any bromine formed is frozen, not likely to sublimate. At least, I think that would work (hint: don't use glass. Go for ceramic.), but it might end up that the solid bromine formed would prevent further reaction, etc.
Alternative: Trap reaction inside a corked gigantic flask (large enough for the bromine gas to spread out without increasing pressure to dangerous levels and add the H2SO4 through a one-way funnel (or add it quickly and stopper it off).

longtermmadscientist - 13-5-2012 at 14:06

Quote: Originally posted by simba

Why don't you just use sulfuric acid? You said you need only a few ml, so yields are not important right?

I'm dealing with a small quantity of NaBr, using a small vessel, as is necessary due to limited resources. I plan on using a pipette to draw- off the bromine from the underlayer (which I'm hoping will form, as Br is denser than HNO3). What I don't want to happen to get so little bromine that I can't pipette more than miniscule quantities of it up without scooping up water with it. I'm guessing, and only guessing, here, that HNO3 would give a higher yield than H2SO4, but there's very little info out there, it seems, on using the former acid for producing Bromine from bromide.

longtermmadscientist - 13-5-2012 at 14:09

Quote: Originally posted by elementcollector1

I would recommend 98% H2SO4, added slowly to very cold NaBr solid. If you really wanted to go the extra mile, you could use dry ice so that any bromine formed is frozen, not likely to sublimate. At least, I think that would work (hint: don't use glass. Go for ceramic.), but it might end up that the solid bromine formed would prevent further reaction, etc.
Alternative: Trap reaction inside a corked gigantic flask (large enough for the bromine gas to spread out without increasing pressure to dangerous levels and add the H2SO4 through a one-way funnel (or add it quickly and stopper it off).

Thanks, refridgeration is definitely something I'll be looking into, due to the exothermic nature of this reaction, or a similar reaction using HNO3. Do you get a good yield with the method you suggested here?

elementcollector1 - 13-5-2012 at 14:50

I have no idea, I just got it from a youtube video.
http://www.youtube.com/watch?v=QkTu7yKSyzg
You can clearly see the enormous amount of bromine lost as vapor as the reaction predictably heats up.

AJKOER - 13-5-2012 at 16:20

Actually if one doesn't have 98% H2SO4, I would argue that Bromine can be released by the action of NaOCl and vinegar (in place of HOAc, one can also use dilute or concentrated HCl), which forms dilute HOCl (or Cl2 with conc HCl), on NaBr.

Some chemistry (pardon the complexity):

NaOCl + HOAc --> NaOAc + HOCl

HOCl + NaBr --> HOBr + NaCl

3 HOBr --> HBrO3 + 2 HBr

Or:

5 HOBr ---> HBrO3 + 2 Br2 + 2 H2O (on heating, low pH)

As:

HBr + HOBr <--> Br2 + H2O (the reverse reaction is much slower than occurs with chlorine)

On net we have therefore approximately:

NaOCl + HOAc + NaBr ---> NaOAc + NaCl + 1/5 HBrO3 + 2/5 Br2 + 2/5 H2O

Approximately owing to the minor secondary reaction:

2 HOBr --> 2 HBr + O2 (g)

and the post reaction additional of H2O2 may remove any free HBr:

2 HBr + H2O2 --> Br2 + 2 H2O

(Source: https://docs.google.com/viewer?a=v&q=cache:5FnZ5VbTCUAJ:... )

Note, at this point in the reaction with acetic acid or dilute HCl, some the Bromine is locked away as bromate. With strong HCl, the reaction of Chlorine gas on aqueous NaBr proceeds much more simply as:

2NaBr + Cl2 → Br2 + 2NaCl

Here is a patent that references these reactions and add more detail on the reaction of Bromine in water and the influence of pH:

http://www.patentstorm.us/patents/5516501/description.html

Here is a YouTube video on the NaOCl/HCl/NaBr path to Bromine:
http://www.youtube.com/watch?v=BdrM9oj-Cyc&feature=relat...
-----------------------------------------

Now, before someone attempts to criticize the HOCl synthesis over the possible loss of some of the Bromine to Bromate, one may be able to recover the loss by adding more NaBr and an excess of acid paralleling the reaction with Iodine and Iodate. In particular, per "Synthetic inorganic chemistry: a course of laboratory and classroom study ...", by Arthur Alphonzo Blanchard, Joseph Warren Phelan, page 232:

"Iodate and iodide ions alone have no action on each other, but with hydrogen ions present a mutual oxidation and reduction of the iodine takes place.

6H+ + 5 I- + IO3- -> 3H2O + 3 I2

No oxidation or reduction of the hydrogen occurs, but the hydrogen ion is used up, which explains why the presence of acid is necessary to make the reaction take place." where I would assume a similar reaction based on Bromine. Link:
http://books.google.com/books?ei=E1ewT5P1IOaR6gHf9fnWAQ&...

EDIT: Found a related equation for Bromine reference. To cite two reactions from "White's Handbook of Chlorination and Alternative Disinfectants" by Geo. Clifford White, page 860:

KBrO3 + 2 KBr + 6 HCl --> 3 BrCl + 3 KCl + 3 H2O

and: BrCl + H2O --> HOBr + HCl

Link: http://books.google.com/books?id=9_2idzksARMC&pg=PA859&a...

[Edited on 14-5-2012 by AJKOER]

[Edited on 14-5-2012 by AJKOER]

woelen - 13-5-2012 at 22:31

All methods, described above are utter crap for making Br2.

I myself tried adding HNO3 to solid NaBr. It does not work. The reaction at first does not seem to start at all, you must have some NO2 in the mix already to get it started. That can be done by adding a small pinch of Na2SO3 or NaNO2. Once this is added you get bromine, but also copious amounts of NO2 and I think that the bromine will be heavily contaminated with this. I only obtained a few small droplets of Br2 from this reaction though. One of the big problems is that NaBr hardly dissolves in conc. HNO3, it becomes covered by an insolutable layer of NaNO3, which prevents further reaction.

Adding conc. H2SO4 to solid NaBr is even worse. The stuff starts foaming a lot and you get HBr-gas, which is contaminated with Br2 and SO2. But the yield of Br2 is so low, that you hardly can speak of more than serious contamination. No visible liquid droplets of bromine are formed, you only see gaseous Br2.

Adding bleach to NaBr does give Br2, but it is so dilute that it does not separate from the aqueous solution. Vinegar as acid is equally bad, it dilutes even more and this acid is weak.

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

Some decent ways of making Br2:

- Electrolysis of a concentrated solution of NaBr, such that 1/6-th part of the bromide is converted to bromate and then adding a slight excess of 50% H2SO4 or saturated solution of NaHSO4. I did this myself. It works very well. You need a graphite anode and unfortunately, you need a pinch of chromate or dichromate in order to have a decent current efficiency. Here follows the write-up:

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

Another option:

- Make a concentrated solution of NaBr and for each mole of NaBr add well over 1 mole of appr. 50% HNO3. Allow mixing of the solution and allow crystals of NaNO3 to separate from the liquid. This can occur if the acid is too concentrated. Instead of using HNO3 you can also use HCl. In that case use at least 30% HCl, otherwise the solution is too dilute. Do not use H2SO4, that will lead to a terrible mud of insoluble CaSO4 in the following step of the process!

- For each gram of NaBr take half a gram of calcium hypochlorite with 70% chlorine content (for swimming pools). If the chlorine content is less or more, then adjust the computation. Do not use more than half a gram of less than half a gram. Both excess and too little will result in loss of bromine.

- Carefully add the solid to the solution of NaBr/HNO3. Do this in small increments!!! When all of it is added, bromine will settle at the bottom.

I like the electrolysis process most. The bromine, produced in this way is very pure. The bromine, made with hypochlorite almost certainly contains some chlorine in the form of BrCl. This is invisible though. For fun energetic experiments with bromine, or when the bromine is stored as a sample in an ampoule this chlorine content is no issue. If the bromine is used for bromination in organic syntheses, then the presence of the chlorine may be disturbing.

weiming1998 - 14-5-2012 at 00:31

What about bubbling Cl2 using a simple gas generator through concentrated NaBr solution? Bromine has a limited solubility in water, and after some time, a layer of liquid bromine will form at the bottom, which you can then purify by distillation, then chilling with a cold ice bath.

woelen - 14-5-2012 at 02:35

Yes, reacting Cl2 with a solution of NaBr also works, but I found the practical setup for this rather cumbersome. Simply bubbling Cl2 through a solution did not work very well for me. The Cl2 reacts too slowly, a lot of it just bubbles through the solution and some Br2 is drawn out of solution as well. I used a normal beaker and that was not a great success. Maybe if you use a long thin glass bottle or a longdrink glass and insert the bubbles of Cl2 at the bottom and disperse these bubbles as very fine bubbles you might have better results.

Another option might be to use an inverted funnel and put its rim just under the liquid surface. Initially, air will be driven away and bubbles of air escape along the rim of the funnel, but as the concentration of Cl2 is rising, more and more Br2 will form at the surface of the liquid under the funnel and the Br2 will sink to the bottom. Generation of Cl2 should be slow though in this setup. If generation of Cl2 becomes too slow, then liquid will be sucked up in the funnel, but this causes the liquid level to go down and if the funnel is put at such a height that it looses contact with the surface of the liquid when it is sucked half full, then you will not experience suck-back of liquid into your chlorine gas generator.

I'm sure that with some experimenting you will get it working, but I never bothered after my initial bad experiences with it. The other methods worked more easily for me and I felt no need to try again. I know of other people over here, who succesfully applied the bubbling of Cl2 through bromide solution, so it can be done.

Endimion17 - 14-5-2012 at 06:25

woelen, what about concentrated hydrogen peroxide and sulphuric acid with concentrated solution of alkali bromide? It works quite well.

woelen - 14-5-2012 at 06:45

Yes, the latter also works quite well. Actually, there are so many methods of making bromine which work quite well. Any decent oxidizer, combined with a strong acid and a soluble bromide does the job. The only thing is that you need solid chemicals or at least highly concentrated solutions if you want to have bromine separate from the aqueous layer.

I made Br2 with H2O2, KMnO4, Ca(ClO)2, NaClO3, KBrO3 as oxidizer. Electrolysis is another option, as written above. The following are not suitable and do not work at all or only marginally: nitrates, perchlorates, conc. H2SO4, persulfates. These oxidizers are not strong enough or react too sluggishly. TCCA also is not really suitable, due to the formation of a mud of insoluble cyanuric acid. I expect that dichromates also work reasonably well, but some heating or patience may be needed with them.

Zan Divine - 14-5-2012 at 07:10

All factors considered, I favor a warm, stirred mixture of NaBr & KMnO4, treated dropwise with conc. H2SO4, which will quickly warm to a temp above the bp of Br2. The vapors are condensed, washed, dried over NaBr & distilled.

There are no interhalogens as with Cl2. This is operationally simple and it is considerably cheaper than H2O2, which is important for larger batches. Look at Wiki under bromine for the stoichiometry.

[Edited on 14-5-2012 by Zan Divine]

barley81 - 14-5-2012 at 08:44

Has anyone tried manganese dioxide? In <a href="http://books.google.com/books?id=OkUAAAAAYAAJ&printsec=frontcover#v=onepage&q=bromine&f=false">this book</a>, there is a procedure which uses manganese dioxide on page 124. Since manganese dioxide can be used to make chlorine gas, and since bromine is less volatile than chlorine and is not as good an oxidizer, this might be a good method.

[Edited on 14-5-2012 by barley81]

Endimion17 - 14-5-2012 at 08:49

Quote: Originally posted by woelen

Yes, the latter also works quite well. Actually, there are so many methods of making bromine which work quite well. Any decent oxidizer, combined with a strong acid and a soluble bromide does the job. The only thing is that you need solid chemicals or at least highly concentrated solutions if you want to have bromine separate from the aqueous layer.

I made Br2 with H2O2, KMnO4, Ca(ClO)2, NaClO3, KBrO3 as oxidizer. Electrolysis is another option, as written above. The following are not suitable and do not work at all or only marginally: nitrates, perchlorates, conc. H2SO4, persulfates. These oxidizers are not strong enough or react too sluggishly. TCCA also is not really suitable, due to the formation of a mud of insoluble cyanuric acid. I expect that dichromates also work reasonably well, but some heating or patience may be needed with them.

H2O2 and H2SO4 (both concentrated) represent a quite clean method, wouldn't you agree? The only thing left to remove is water. Bromine can be pippetted out and water traces removed by some H2SO4 and distillation. IMO, the best option for getting pure bromine suitable for organic synthesis.

Alkali bromide really should be in a concentrated solution because if solid, it won't react completely as a cake of bromide and sulphate is formed.

longtermmadscientist - 14-5-2012 at 12:37

OK, thanks, guys, this has given me a bit to think about. I’m a bit disappointed that HNO3 turns out not to be a strong enough oxidant to do the job; my reasoning was that if Cl can be produced by adding HNO3 to NaCl, then you should be able to get bromine from NaBr, but I had at the same time in mind the reluctance of bromine to oxidise (as exemplified by the relative instability of bromic acid as compared to the analogous acids of Cl and I, although I may be misunderstanding something here), so perhaps the result isn’t entirely surprising. If it isn’t the oxidation potential of HNO3 that is so much the problem, and precipitation/ relative insolubility of NaNO3 which is ( which surprises me, as from what I can gather the two have quite similar solubilities), then I may have to dissolve the NaBr before I add the HNO3, which means that then I’d have to add excess HNO3 to ramp- up the oxidising power of the solution. Adding other oxidants is I’m afraid not possible at the moment, excepting household bleach, and chlorine impurities are acceptable in small amounts, as I don’t do organic chemistry. My purposes in producing bromine are producing enough of the element to learn of its chemical and physical properties first hand, and to have some in stock so that I can do further experiments on if need be.

AJKOER - 14-5-2012 at 14:56

Commercial methods (see link below as an example from the Dead Sea plant) for the preparation of Bromine do not use strong acid as an aid in removing water from Bromine reactant mix (cost factor I would guess). Instead, there is a separation stage process which employs heating a Br2 water solution with steam and using the physical properties of Bromine (much heavier than water with a limited solubility) to facilitate separation (see page 86).

With respect to chemistry, the reference process attached just uses Cl2 (mentioned in my synthesis) on a bromide although I see more potential value in an HOCl method via a weak acid (preferably a low cost weak acid). My opinion is that such a method reduces the Cl2 presence (as well as BrCl formation, a function of moderate temperature, excess Chlorine and low pH), and introduces, instead, a small amount of free Cl2O depending on the concentration of the HOCl employed. However, to reduce initial dilution, one can simply employ Boric acid in place of Acetic and a high concentration NaOCl (or use Ca(OCl)2 access permitting). Barring the Chlorine route, the HOCl method, however, appears to suffer some yield loss due to bromate formation (see my edited thread).

A valid comment of the synthesis I presented is that it is more in line with a current commercial process for the extraction of Bromine than a laboratory synthesis where a different set of criteria (like percent yield, reagents' cost and reaction speed) are more prominent factors. However, for garage chemists, reagents' cost and access as well as yield, are likewise important factors.

Link:
http://www.weizmann.ac.il/sci-tea/Brombook/pdf/chapter3.pdf

[Edited on 15-5-2012 by AJKOER]

elementcollector1 - 14-5-2012 at 15:21

What of electrolysis: I have a makeshift membrane-separator, can I get Br2 and NaOH out of this like I would from an identical NaCl cell?

Edit: Electrolysis is working beautifully, and my makeshift filter paper membrane is holding up well (the two sides are slightly imbalanced by volume). So, free NaOH (filled catholyte with pure water) and free Br2 (is sinking to the bottom. Doesn't really seem to be soluble).

[Edited on 15-5-2012 by elementcollector1]

Edit 2: Membrane broke, cell overheated.

Now I have a yellow supernatant liquor, and a small amount of reddish bromine water. Should I start over, and use H2O2 and HCl to get the desired results?

[Edited on 15-5-2012 by elementcollector1]

BromicAcid - 14-5-2012 at 17:40

For crying out loud, use the search engine:

http://www.sciencemadness.org/talk/viewthread.php?tid=6953

This topic has been discussed ad nauseam. Not only are there several refined procedures but there are less refined procedures using most any oxidizing agent that you can get your hands on. Please, there is enough practical experience with this to bore anyone to death so armchair chemists need not apply.

elementcollector1 - 14-5-2012 at 20:20

I've been trying to make sodium bromate for the bromate+bromide+acid procedure, I get a yellow solution (the same color when I attempted to make sodium hydroxide/chlorate). Is this the correct color for the hypobromite? I don't think so, but I have not a clue as to what it really is. I also get bromine water, which sort of sinks to the bottom but not really a distinct layer as of yet. I have the 2-amp car battery running the electrolysis overnight, so come morning most to all of the bromine should have been converted into something.
Also, I have NO distillation setup - how do I dry and collect this stuff? Sulfuric acid seems like a good idea, but I'd have to concentrate this annoying pink 10% stuff I found at a pool store (can't find any better local sources).
I could pass Cl2 into NaBr (I have a lot of MnCl2 I need to make anyway), so tell me: Which works better, bromate/bromide/acid or bromide/chlorine?

AndersHoveland - 14-5-2012 at 21:54

NaBr + sodium bisulfate + dilute H2O2
makes Br2
The procedure is very simple, and works well if one does not need to isolate the bromine from solution.

[Edited on 15-5-2012 by AndersHoveland]

scatha - 15-5-2012 at 04:07

What about this reaction ?

4 NaHSO4 + MnO2 + 2 NaBr -> MnSO4 + 3 Na2SO4 + Br2 + 2 H20

All educts and products are solid, except the produced Br2 and the H2O - maybe this is helpful for your case ?

[Edited on 15-5-2012 by scatha]

[Edited on 15-5-2012 by scatha]

Vargouille - 15-5-2012 at 05:32

Scatha, that would only work in a solvent that isn't water, since MnSO4 and Na2SO4 are both fairly soluble in water. I would have to check on the solubilities of all of the materials involved to find a good solvent to precipitate out the MnSO4 and the Na2SO4, but if you enjoy a challenge, it might be possible.

woelen - 15-5-2012 at 06:00

The reaction of Scatha most likely can make Br2, but I'm quite sure some heating is required in order to get the powdered reagents to react. Solid-solid reactions are not the best ones to work with, so this is not a good idea for practical production of Br2. If some water is added, then it might work, provided the used MnO2 is not calcined and very inert.

But let's keep BromicAcid's remark in mind. Production of Br2 is discussed many times already and even in this thread we now hve seen a few useful and suitable methods. Now it is time to take out the glassware and chemicals and try things. Reporting back on your results is strongly encouraged.

longtermmadscientist - 15-5-2012 at 12:26

Quote: Originally posted by woelen

But let's keep BromicAcid's remark in mind. Production of Br2 is discussed many times already and even in this thread we now hve seen a few useful and suitable methods. Now it is time to take out the glassware and chemicals and try things. Reporting back on your results is strongly encouraged.

Will do, woelen, soon after my NaBr arrives. I thought I’d check out the feasibility of doing this before it arrived, as knowing me I’d probably have impulsively tried something I’d only just thought of, and using up a lot of my reagents. Thanks for your valuable insights from your own experiments.

I was unsure about whether to open a new thread on this one, and I did in fact search the forum for “bromine” and “HNO3/nitric acid", but nothing came up with the two tied closely, so I took the decision to open a new thread. If one of the mods would like to merge this topic with others on bromine production, then that’s fine, and it would possibly make things tidier.

[Edited on 15-5-2012 by longtermmadscientist]

[Edited on 15-5-2012 by longtermmadscientist]

Endimion17 - 15-5-2012 at 12:44

"bromine isolation" *chemistry nazi*

AJKOER - 17-5-2012 at 18:41

Quote: Originally posted by AJKOER

Now, ...over the possible loss of some of the Bromine to Bromate, one may be able to recover the loss by adding more NaBr and an excess of acid paralleling the reaction with Iodine and Iodate. In particular, per "Synthetic inorganic chemistry: a course of laboratory and classroom study ...", by Arthur Alphonzo Blanchard, Joseph Warren Phelan, page 232:

"Iodate and iodide ions alone have no action on each other, but with hydrogen ions present a mutual oxidation and reduction of the iodine takes place.

6H+ + 5 I- + IO3- -> 3H2O + 3 I2

No oxidation or reduction of the hydrogen occurs, but the hydrogen ion is used up, which explains why the presence of acid is necessary to make the reaction take place." where I would assume a similar reaction based on Bromine. Link:
http://books.google.com/books?ei=E1ewT5P1IOaR6gHf9fnWAQ&...

For those interested in the chemistry of Bromine, I have found precisely the same reaction referenced for Br2, however, clearly expressed in an equilibrium form. Source: "Inorganic Chemistry" by Egon Wiberg, A. F. Holleman, Nils Wiberg, page 450, equation (1):

BrO3(-) + 5 Br(-) + 6 H(+) <-----> 3 Br2 + 3 H2O

Link: http://books.google.com/books?id=Mtth5g59dEIC&pg=PA450&a...

Note, the following important comment by the author:
"Clearly, an increase of the proton concentration and decrease of pH causes the rapid protonation equilibrium (2a) to shift to the right; since this equilibrium precedes the rate determining step, the overall reaction rate increases", which echoes Woelen comments on acid strength.

So theoretically, with the appropriate addition of excess Bromide and acid, any Bromate formation may be reduced with the recovery of Bromine. Quite interestingly, add 1/5 of the equation above:

1/5 HBrO3(-) + Br(-) + H(+) <-----> 3/5 Br2 + 3/5 H2O

to my prior cited reaction:

NaOCl + HOAc + NaBr ---> NaOAc + NaCl + 1/5 HBrO3 + 2/5 Br2 + 2/5 H2O

produces the amazing simple and intuitive target net ionic equation:

OCl- + 2 Br- + 2 H+ ---> Cl- + Br2 + H2O

Thus, the outstanding question with the hypochlorite method (awaiting verification) is even with steps to reduce dilution (as concentrated HOBr is more unstable with its decomposition yielding Bromine, perhaps with the help of dry acid salts like Tartaric or Critic acid and also use of conc hypochlorite) and decreased pH (target range is a pH of about 2 to about 6.4 per previously cited Patent 5516501), and a bromide to hypochlorite molar ratio of at least two per the target equation above; will distillation then allow the recovery of elemental Bromine.

[Edited on 18-5-2012 by AJKOER]

elementcollector1 - 17-5-2012 at 20:11

I tried making bromate by 1-cell electrolysis of the bromide, but I don't know how it went. Is this the same as the iodine synthesis of tincture of iodine + H2O2 + HCl?

longtermmadscientist - 19-5-2012 at 05:08

OK, so I added a large excess of conc. HNO3 to solid NaBr dihydrate and the reaction started immediately- a deep red layer formed over the NaBr, and an even deeper red substance, probably Br2, formed amongst the crystals of NaBr dihydrate. However, the reaction seemed to stall. I think that the reason for this may be that bromine is denser than the NaBr dehydrate, and the latter forms a protective layer over the bromine. As the NaBr gradually reacts with the HNO¬3, and forms NaNO3, this effect continues, as NaNO3 is about as soluble as NaBr. Also, even if these compounds dissolve, there will still be a delay in the mixing of this salt layer and the HN03, and access of the latter to the NaBr. I may repeat the experiment but using a saturated solution of NaBr.

[Edited on 19-5-2012 by longtermmadscientist]

BromicAcid - 19-5-2012 at 07:16

Wait, were you trying to make nitrosyl bromide?

I would think that if you're trying to make bromine from NaBr / HNO3 obviously one gets reduced one gets oxidized.

2NaBr + 4HNO3 ----> Br2 + 2NaNO3 + 2NO2 + 2H2O

Then nitrogen dioxide can react with bromide salts:

2NO2 + NaBr ----> NaNO3 + NOBr

Just something to consider.

[Edited on 5/19/2012 by BromicAcid]

longtermmadscientist - 19-5-2012 at 12:26

Quote: Originally posted by BromicAcid

Wait, were you trying to make nitrosyl bromide?

I would think that if you're trying to make bromine from NaBr / HNO3 obviously one gets reduced one gets oxidized.

2NaBr + 4HNO3 ----> Br2 + 2NaNO3 + 2NO2 + 2H2O

Then nitrogen dioxide can react with bromide salts:

2NO2 + NaBr ----> NaNO3 + NOBr

Just something to consider.

[Edited on 5/19/2012 by BromicAcid]

Ah....that’s something I didn’t think of; thanks. I would hope that not much of this would form, as this is something else which could contaminate my bromine. I may have to use H2SO4, but I’m not enthusiastic about the yield. It seems that methods for producing bromine without involving distillation may well be unfeasible.

woelen - 20-5-2012 at 12:19

No, you do not need distillation. You can get a yield of e.g. 70% without distillation and the resulting bromine also can be quite pure. But you need the right chemicals for doing so.

HNO3 + NaBr is not the right combination. I already wrote about that before. The stalling of the reaction is even something which I predicted in one of my previous posts, due to limited solubility of NaNO3 in conc. HNO3. Your bromine most likely will be heavily contaminated with NOx or ONBr.

H2SO4 is even worse. Yields will not be better than a few percents, and for each mole of Br2 you also get a mole of SO2. This also seems like a bad contaminant and as soon as some water is present in the system, the water, SO2 and Br2 will react to H2SO4 and HBr.

longtermmadscientist - 21-5-2012 at 12:33

Right, this is something which is going to take more preparation than I thought. I tried the HNO3 method because this was an oxidant I had to hand and was curious as to why there was so little mention in the literature of producing bromine from bromides using it. Now I know from both theory and practice that this approach isn’t effective and why. If and when I manage to get any of the chemicals necessary for the more effective methods, I’ll be giving them a try, and will report back with my results. Thanks to everyone for their feedback and help.

AJKOER - 22-5-2012 at 12:18

Quote: Originally posted by AJKOER

OCl- + 2 Br- + 2 H+ ---> Cl- + Br2 + H2O

Thus, the outstanding question with the hypochlorite method (awaiting verification) is even with steps to reduce dilution (as concentrated HOBr is more unstable with its decomposition yielding Bromine, perhaps with the help of dry acid salts like Tartaric or Critic acid and also use of conc hypochlorite) and decreased pH (target range is a pH of about 2 to about 6.4 per previously cited Patent 5516501), and a bromide to hypochlorite molar ratio of at least two per the target equation above; will distillation then allow the recovery of elemental Bromine.

The prior mentioned Patent 5,516,501 issued on May 14, 1996 ( http://www.patentstorm.us/patents/5516501/description.html ) stated "Processes for preparing a relatively concentrated aqueous solution of about 700-3000 ppm hyprobromous acid are provided. Hypochlorous acid solutions are prepared by either reacting chlorine gas with water or sodium hypochlorite with an acid. The resulting hypochlorous acid is then reacted with an alkali metal or alkaline earth bromide in order to form the hypobromous acid. Critical parameters are pH, Br/Cl mole ratio, and chlorine concentration. Under optimum conditions, substantially 100% conversion of bromide to hypobromous acid can be attained."

The author also notes, "The bromide/chlorine mole ratio is also important. At pH values above about 3.0 it has been observed that as this mole ratio approaches an optimum value of 1.5 or greater, the yield of hypobromous acid approaches 100%."

A preceding patent (Patent 5795487, Link: http://www.docstoc.com/docs/49815324/Process-To-Manufacture-... ) to the one referenced above noted that:
"a. Mixing an aqueous solution of alkali or alkaline earth metal hypochlorite with a water soluble bromide ion source;

b. Allowing the bromide ion source and the alkali or alkaline earth metal hypochlorite to react to form a 0.5 to 30 percent by weight aqueous solution of unstabilized alkali or alkaline earth metal hypobromite; "

So one may consider changing the mixing order of the target equation above and react say, concentrated aqueous NaOCl with NaBr and then add a weak acid to liberate Hypobromous acid (as HOBr is a very weak acid), as a path to Bromine.

Reaction example, with Bleach, NaBr and Tartaric acid:

NaOCl + NaBr ---> NaCl + NaOBr

2 NaOBr+ C4H6O6 ---> Na2C4H4O6 + 2 HOBr

5 HOBr ---> HBrO3 + 2 Br2 + 2 H2O (more so for conc HOBr with mild heating or sunlight and low pH)

With respect to the stability of HOBr and its propensity to form Bromine, a source states:

"Direct sunlight will have a negative impact on the appearance of the product [referring to HOBr] by changing the color from clear to degrees of yellow or orange. The orange color is bromine (Br2), which is normally not a problem, but if allowed to remain in the sunlight more bromine will develop, which will have a distinct pungent odor. The solution is still useable in the "yellow-orange‟ state, but these conditions should be avoided by not allowing the precursor to be subjected to direct sunlight (UV) or excessive continuous heat to allow the temperature of the bulk liquid to exceed 90° F."

The author also notes: "Hypobromous acid degradation rate accelerates with increasing concentrations. The decay rate for a 200-300 ppm solution of available bromine would result in a half-life of about 10 days, whereas a 4000 ppm solution may only have a half life decay rate of only a few hours or less."

And also: "When hydrogen bromide and bleach are added in the correct proportions, the pH of 6.9-7.4 will always be the result, regardless of the hardness or alkalinity of the feed water. The decay of hypobromous acid always results in either hydrogen bromide or bromine (Br2) formation, or a mixture of both, because the decay of hypobromous acid is an acid-releasing process. At higher temperatures this decay rate increases proportionally with the rise in temperature. As the pH drops, the decay of hypobromous acid accelerates and it becomes an autocatalytic process. At concentrations less than 1,000 ppm this is a slow and relatively harmless side reaction."

So forming more concentrated HOBr solutions in the presence of sunlight or mild heat in a more acidic environment should help liberate Bromine. Link:
http://www.stabilizedbromine.com/pdf/Generation%20%28Blendin...

[Edited on 22-5-2012 by AJKOER]

Zan Divine - 29-5-2012 at 07:54

Quote: Originally posted by longtermmadscientist

Right, this is something which is going to take more preparation than I thought........ If and when I manage to get any of the chemicals necessary for the more effective methods, I’ll be giving them a try,......

One of my clients just ordered 300-400 g bromine. I got the H2SO4 for $10 from a hardware store. The KMnO4 came from e-Bay ($16 for 1 pound delivered). Leslie Pools sells KBr (4 Lbs for $33). For $59 and one afternoon I'll make way more than 300 - 400 g Br and have lots of starting materials left over.

Do any of the other methods discussed give this much interhalogen-free bromine for this low a cost with so simple a procedure? 6 to 7 g Br for $1. I tend to doubt it. H2SO4 is the second cheapest strong acid. KMnO4 is cheaper than appropriately concentrated H2O2 by a wide margin (I don't even consider any chlorine-containing oxidants as I don't want any BrCl).

Since the bromine spontaneously distills from the mixture due to the heat of reaction, collection of the crude product is as simple as adding a condenser column or still head.

So, while lots of routes are available, I don't see one better than this if your goal is to obtain Br2 and the chemistry is just a tool to get there.

Considering how simple the chemistry is, why complicate it? I'd think the things you can do with Br2 would be much more interesting than making it. When you're young, time seems infinite, but still, why waste it?

[Edited on 29-5-2012 by Zan Divine]

Endimion17 - 29-5-2012 at 08:14

Quote: Originally posted by Zan Divine

All factors considered, I favor a warm, stirred mixture of NaBr & KMnO4, treated dropwise with conc. H2SO4, which will quickly warm to a temp above the bp of Br2. The vapors are condensed, washed, dried over NaBr & distilled.

There are no interhalogens as with Cl2. This is operationally simple and it is considerably cheaper than H2O2, which is important for larger batches. Look at Wiki under bromine for the stoichiometry.

[Edited on 14-5-2012 by Zan Divine]

How well does the reaction proceed if the vessel is being cooled by an ice bath to prevent the fuming? I don't like the idea of immediate bromine vapors. It seems better to me if it's left to accumulate on the bottom, from where it can be easily drawn out, dried and distilled with a small amount of conc. sulphuric acid.

Also, wouldn't conc. hydrogen peroxide be cheaper, considering KMnO4 is hard to find in a non "analytical grade purity" state? Peroxide is easy to find concentrated, but not extremely pure, though sufficiently pure for this application.

[Edited on 29-5-2012 by Endimion17]

blogfast25 - 29-5-2012 at 10:29

Peroxide is cheaper. KMnO4 is great but 'overkill' for many reasons here.

elementcollector1 - 29-5-2012 at 12:38

Bleach works too, have a solution of bromine in water sitting in my garage as we speak. Making a distillation apparatus, though, is quite another thing.

Endimion17 - 29-5-2012 at 13:18

^^no, it sucks. We're talking about yields here.
Bromine water is for testing hydrocarbon saturation. Not for making bromine.

elementcollector1 - 29-5-2012 at 13:44

Some of us don't have access to KMnO4 or H2SO4, and thus we make do.

Zan Divine - 29-5-2012 at 19:17

I guess it's no surprise that cost is location dependent. Around here, New York, concentrated H2O2 is expensive and hard to get. Maybe because of acetone peroxide...

None of the British suppliers will ship here anymore.

Endimion17, as a lifelong synthetic chemist, I guess I tend to go straight to the heart of a synthesis with an eye toward minimizing manipulations and optimizing simplicity. This method requires no cooling bath, no heating bath, and absolutely no manipulations left to do after the reaction is done except drying & distilling. To me, there is synthetic elegance in making the reaction provide the driving force for the isolation and having the crude reaction product look like the below. You seem a bit leary about handling boiling bromine. If so, it's probably better if you don't because of toxicity concerns. If you want a cool reaction, peroxide is superior because of the greater ease in separating layers. I'm guessing you're in Britain so that's easy.

[Edited on 30-5-2012 by Zan Divine]

elementcollector1 - 29-5-2012 at 19:59

I've handled chlorine before without dying, so I can deal with bromine. Should I put anything into the distillation apparatus to deal with any salt that might have come over with the bromine?

Endimion17 - 30-5-2012 at 05:55

Quote: Originally posted by Zan Divine

I guess it's no surprise that cost is location dependent. Around here, New York, concentrated H2O2 is expensive and hard to get. Maybe because of acetone peroxide...

None of the British suppliers will ship here anymore.

That's a shame. Things like that piss me off.

Quote:

Endimion17, as a lifelong synthetic chemist, I guess I tend to go straight to the heart of a synthesis with an eye toward minimizing manipulations and optimizing simplicity. This method requires no cooling bath, no heating bath, and absolutely no manipulations left to do after the reaction is done except drying & distilling. To me, there is synthetic elegance in making the reaction provide the driving force for the isolation and having the crude reaction product look like the below. You seem a bit leary about handling boiling bromine. If so, it's probably better if you don't because of toxicity concerns. If you want a cool reaction, peroxide is superior because of the greater ease in separating layers. I'm guessing you're in Britain so that's easy.

I do appreciate the simplicity during the laboratory work. That's why it seems to me it's better to mix the reagents in an ice bath and just leave them alone in a weakly stoppered flask. Bromine essentially stays inside. It takes more than one hour before it's sufficiently progressed. I've tried it, and the only problem I see is that not many of us have any stirring equipment. Because it's a saturated bromide solution, and effectivelly a kind of piranha solution, all cooled down to ~0 °C, the reaction should be agitated because otherwise, layers occur, that are mixed by diffusion and convection by small oxygen bubbles, which is too slow.
However, decent mixers can be made using an electromotor and a shaped glass rod attached to it. Old school, but effective, and there aren't any teflon stirbars soiled with bromine.

Boiling bromine doesn't pose a problem to me because of the toxicity as I have experience with it. Lowering the yield does. The crude product has to be distilled, as you know. If one wants a reasonably pure bromine, doing it your way is essentially distilling it twice. Given the probable scenario that one doesn't have a semimicro destillation kit with a tight and tiny cooler, it means a significantly lowered yield.
Doing it slowly, in an ice bath, means the losses are with dissolving in the reaction mixture only, which are reasonably smaller.

It's all about the yields, because bromides are hard to find where I live (Croatia). AFAIK, there aren't any crude and cheap products, only analytical ones.

Zan Divine - 30-5-2012 at 15:13

Quote: Originally posted by elementcollector1

I've handled chlorine before without dying, so I can deal with bromine. Should I put anything into the distillation apparatus to deal with any salt that might have come over with the bromine?

Actually, bromine is somewhat more toxic/sneaky if memory serves, but your point is well taken, they are similar.

No salt should be entrained in the bromine vapors to speak of. It wouldn't matter anyway. The next step is to stir the crude bromine with some finely ground and thoroughly dried NaBr to dry it. Conc. H2SO4 can also be used. It just depends if you prefer a sep. funnel or a filter to remove your drying agent.
I use a funnel with a glass wool plug, outside.

There is a patent that describes holding Br2 just below boiling for some time to dry it, but I've never tried it.

If a tiny bit of salt gets through, no problem,it stays in the pot during distillation.

If you choose this method you can just start with HOT tap water (I can hear the purists shrieks but it doesn't matter one iota). Similarly, KMnO4 can be any purity, I've used 80 - 98% and the results are the same. The H2SO4 is just drain cleaner. Work with as concentrated solutions as you can get to maximize yields.

All that being said, everybody optimizes their syntheses to suit their equipment and you've made good points, Endimion17. Now that I know all the factors in play, you're probably ahead of the game with H2O2

[Edited on 30-5-2012 by Zan Divine]

elementcollector1 - 30-5-2012 at 20:17

I choose NaBr, as I have a steady supply.

How do I thoroughly dry this, a microwave? (Did it with MnO2 sludge to great effect)
Now, to find some glass tubing. I have no idea where to start looking.

barley81 - 31-5-2012 at 07:58

Elemental scientific has some. You could look for a neon sign shop or a glassblower in your area. They will probably have it. You could also get some on eBay..

Zan Divine - 31-5-2012 at 14:14

ec1, Yes, the microwave solution is perfect. Just heat the salt in 30 sec to 1 minute intervals until it no longer gets real hot. Then cook the hell out of it for several minutes (It's best not to leave it alone. Some microwave ovens have been known to react poorly to being driven with no real load to soak up the microwaves. They'll arc. If you turn it off right away, no harm done).

If you have a balance you can simply weigh it between heating periods until no more weight is lost. This is called "drying to constant weight".

E-bay would seem to be the best source of glassware of all sorts for many people. In the US, there aren't many ways more direct. You'll save over buying from distributors but for most glassware don't expect "a steal". The common things like round bottom flasks and condensers and especially heating mantles are still moderately costly. If you can use the smaller scale 14/35 (or 14/20) glassware instead of the old standby 24/40 you'll save money. E-bay has glass tubing (Stay away from the soft glass, it's a relic. Go for borosilicate or hard glass). I don't know how well e-Bay works for other countries.

[Edited on 31-5-2012 by Zan Divine]

elementcollector1 - 31-5-2012 at 14:44

I *have* a balance, but it kinda sucks. Only accurate to the nearest 50-gram interval. XD
As for the round-bottom flask, which is useful for distillation and making retorts, I was wondering if a lightbulb could withstand the heat. It's so thin I'd expect it to shatter upon touch after heating, but maybe...
How do you "shape" the hard glass? I know you can make holes by passing through a heated piece of copper pipe or wire, but how do you bend it without 'kinking' it?

watson.fawkes - 31-5-2012 at 17:23

Quote: Originally posted by elementcollector1

I know you can make holes by passing through a heated piece of copper pipe or wire, but how do you bend it without 'kinking' it?

Air pressure, usually from the lungs, hence the term glass-"blowing".

There are a number of old glassblowing manuals digitized. Search for them. I don't recall the titles offhand. I've read a number of them; they're all pretty much equally good for the basics.

Endimion17 - 1-6-2012 at 06:01

Quote: Originally posted by elementcollector1

How do you "shape" the hard glass? I know you can make holes by passing through a heated piece of copper pipe or wire, but how do you bend it without 'kinking' it?

You can't pass a heated wire through. No way. The heat dissipation is too great for the glass to reach sufficient temperature.
I recommend downloading one of several available flameworking (not glassblowing, that's just "making bottles") manuals on the Web.

If you want to obtain some skills, buy a handful of glass tubing of various sizes, one usual propane-butane blowtorch and one fine blowtorch with a hot, narrow flame like the ones used in kitchens.

When you get everything, start following the manual and soon you'll be able to join tubes, make small flasks, make holes through flasks, join tubes with those flasks, etc.
But don't fool yourself - you'll never be able to recreate a real flameworker's job, simply because you don't have the neccessary tools.

It's might seem like a futile practice, but believe me, elementary flameworking is essential for a home chemist as it can save a lot of money. I've saved quite a lot just by buying glass tubing and making my own stuff, some of them disposable. I didn't want to spend money on buying something that will have to endure lots of heating.

Forget the light bulbs. They're for entry level home chemist, for boiling water. That's glass with high percentage of sodium. It's soft, thin and fragile, with a too low softening point.
Just buy a handful of glass tubing.

elementcollector1 - 1-6-2012 at 14:33

Got it, thanks.

watson.fawkes - 1-6-2012 at 15:03

Quote: Originally posted by Endimion17

I recommend downloading one of several available flameworking (not glassblowing, that's just "making bottles") manuals on the Web.

Please tell that to the ASGS, the American Scientific Glassblowing Society, and the BSSG, the British Society of Scientific Glassblowers. Also to the Scientific Glassblowing Learning Center, one of the best web sites with instructional material.

If you're doing keyword searches, use "scientific glassblowing". You'll get more hits.

elementcollector1 - 15-6-2012 at 21:04

Been a while since I found this topic (necromancy powers activate!), but I just got most of my distillation apparatus today. I still have a few questions:
-From what I've heard, bromine tends to somehow get 'trapped' in glassware. If this becomes a problem, how do I get rid of it?
-How long should my condenser be for the most effective distillation? (Standard, non-coiled Liebig).
-How fast does the bromine eat through the rubber of the corks? (I couldn't find ground-glass.)

When I get the whole distillation setup done, I'll try to get pics and weight of the yield.

plante1999 - 16-6-2012 at 02:41

I recommend to use ground glass since rubber is not completely destroyed but some oil will be in your bromine.
Sodium thiosulphate sol. will reduce bromine to bromide which is water soluble.

elementcollector1 - 16-6-2012 at 10:44

Where do I get sodium thiosulfate?
So, is there a good length for my condenser to be, or will any length do?

woelen - 17-6-2012 at 22:53

Bromine distillation requires fairly good cooling of your condenser, otherwise too much will escape as vapor. Remember, bromine is very volatile and air-cooling most likely will be insufficient, except maybe on cold days. You certainly will collect some bromine, but having a loss of a few tens of percent is not pleasant at all, especially in a confined space.

If you do not have ground glass, then you can use rubber stoppers and that kind of things, but you have to wrap all these things in a fairly thick layer of white teflon tape, which is used in plumbing. This tape will absorb some bromine and become brown, but it does not fall apart and it does not give off grease or oil. Unprotected rubber or cork is a no go with bromine. It will lead to a lot of contamination of your bromine. It reacts with bromine, giving all kinds of (somewhat) volatile organics, which leach into your bromine.

[Edited on 18-6-12 by woelen]

Zan Divine - 22-6-2012 at 15:15

Quote: Originally posted by elementcollector1

Where do I get sodium thiosulfate?

I just bought a pound on e-bay for $10 delivered.

elementcollector1 - 12-8-2012 at 12:53

Sorry to revive this old thing, but I attempted distillation of bromine today with no success. Then, I had another idea: Since bromine was heavier than air, I could trap the bromine-water in a large flask, boil the flask, then 'pour' the bromine into something cold to condense it. Would this work?

Endimion17 - 12-8-2012 at 15:18

Quote: Originally posted by elementcollector1

Sorry to revive this old thing, but I attempted distillation of bromine today with no success. Then, I had another idea: Since bromine was heavier than air, I could trap the bromine-water in a large flask, boil the flask, then 'pour' the bromine into something cold to condense it. Would this work?

No. The smaller the distilling system and the more efficient the cooling system, the better the yield.

Bromine Systhesis

cal - 12-8-2012 at 17:28

http://www.youtube.com/watch?v=OB4MmPTOBxg, Elemental Bromine :cool:

Quote: Originally posted by longtermmadscientist

I've been thinking of a simple way of producing a few ml of elemental bromine using sodium bromide. Could conc. HNO3 be used on NaBr solid or solution? I've been researching up on this, and there's a lot of talk of H2SO4 but it seems it's not a strong enough oxidant to produce Br2 in good yield, and you get a lot of HBr.

My concern is that I will produce a lot of NO2, which, being the same colour as Br2, could lead to some confusion as to how much Br2 has been produced. Also, I'd like to know whether the bromine would contain more that a trace of NO2. Water impurity can be extracted by conc. H2SO4, but what about removing the NO2, if this turns out to be an issue?

[Edited on 13-8-2012 by cal]

elementcollector1 - 12-8-2012 at 20:49

I don't have much apparatus to bubble chlorine through water... and even if I did, my distillation setup doesn't work well. I'm thinking of getting ground glass stuff from united nuclear, but does anyone have cheaper sources for the entire kit?

MrHomeScientist - 13-8-2012 at 05:23

Dr. Bob, here on this forum, has a huge stock of glassware for sale. I've bought from him 3 or 4 times and he's given me tons of goodies. Get in touch with him for some great deals - last I heard he was running low on full distillation kits though! 24/40 is the size I like to use, but if you're focusing on small scale bromine production you might want to go smaller.

elementcollector1 - 13-8-2012 at 21:03

Working out a deal as we speak... Somewhat fittingly, he only has a smaller kit, 14/20. Worth it!

Pyro - 17-8-2012 at 19:51

it seems like everybody is taking the H2SO4+NaBr method, isn't it a LOT more effective to just replace the Br- with Cl- by simply bubbling Cl2 through it?

elementcollector1 - 17-8-2012 at 21:10

That's what I'm going to try next. The problem with the Cl2 method is that it seems to form interhalogens between chlorine and bromine (BrCl, for instance).

Pyro - 18-8-2012 at 03:11

but why does this guy do it like that: http://www.youtube.com/watch?v=OB4MmPTOBxg
he prefers this to the NaBr+H2SO4, i will just use NaBr instead of KBr for the chlorine method.
anyway, wouldn't almost all the Cl2 react with the NaBr before reacting with the Br2

elementcollector1 - 19-8-2012 at 14:25

That is correct, and he did mention "that if one smells the slightest hint of chlorine, the reaction is done". I would think that if you continued adding chlorine, the bromine would only then begin reacting (as the chlorine has little to nothing else to react with) to make interhalogens.
Any soluble bromide works as long as you use all ground-glass apparatus, a saturated solution, and a necessary amount of caution.

Pyro - 19-8-2012 at 14:45

to rid Br2 of BrCl just keep it in a beaker at about 10*C and the BrCl will evaporate first.

elementcollector1 - 19-8-2012 at 15:19

Good to know! Incidentally, I wonder if one could collect an interhalogen through a similar method to study its properties?

woelen - 19-8-2012 at 22:37

The method of bubbling chlorine through a solution of NaBr or KBr can be done, but for best results, you have to bubble just enough chlorine through the solution that all bromide is converted to bromine. Besides that, you still need a distillation setup to get all dissolved bromine out of the water.

If you don't have a means to do a distillation, then the BaBr/chlorine method is not easy at all. As long as you don't bubble enough Cl2 through the solution, hardly and Br2 separates from the liquid. Br2 dissolves in a solution of a bromide quite well, it forms a complex ion Br3(-), very similar to the I3(-) ion, known from iodine. This Br3(-) ion is more labile than I3(-) and hence, with a distillation setup you can break this ion and drive the bromine out of solution, but if you don't have a distillation setup and intend to pipette the Br2 from below the aqueous layer, then your losses will be huge.

Bubbling too much chlorine is equally bad. You get BrCl, but even worse, part of the bromine is lost as bromic acid. BrCl reacts with water, giving bromic acid, chloride ion and bromine.
Another issue is that when you bubble chlorine through a solution of NaBr not all chlorine reacts, some may make it to the surface and escape, giving Cl2-gas and BrCl-vapor above the aqueous layer. You need to disperse the Cl2-gas into very small bubbles.

Pyro - 20-8-2012 at 10:06

i got a distillation setup, and isn't using Cl2 still more effective than the H2SO4+NaBr method?
how are yeilds for that method?

vmelkon - 21-8-2012 at 05:44

Quote: Originally posted by Pyro

i got a distillation setup, and isn't using Cl2 still more effective than the H2SO4+NaBr method?
how are yeilds for that method?

According to myst32YT, it is more efficient to bubble Cl2 through NaBr solution.
I think it makes sense because the H2SO4+NaBr would probably waste some Bromine as HBr gas.

Both methods are going to produce contaminated bromine.

The ultimate method is electrolysis. Get PbBr2, melt it, electrolyze using a couple of graphite electrodes. This method was used to demonstrate element production in classes. Then it got replaced by ZnCl2.

Perhaps you can try ZnBr2?

elementcollector1 - 21-8-2012 at 11:05

True, but for us amateurs, it's already hard enough to distill bromine from a chemical mix, so while the bromine from the molten bromide would be very pure, how on earth would one get it out of there, keeping in mind that we don't have expensive equipment (at least, MOST of us don't)?

vmelkon - 21-8-2012 at 12:41

It doesn't require expensive equipment. You need a heat resistant ceramic dish. Perhaps some of us don't have one (I don't). I would use a metal can (the sawed off bottom of a propane can).

The melting point of ZnBr2 is only 394 °C. I would use propane torch.

Now for the important part. To capture the bromine gas, you need glass tubing and the graphite anode needs to be inside it. This is something that I have made myself in order to capture iodine vapors.

elementcollector1 - 24-8-2012 at 12:01

...Can we see your custom apparatus?
Also, one would need to lead the bromine away from the reaction into a suitable cooling vessel, as the bromine would definitely come out gaseous at 400 C.

vmelkon - 25-8-2012 at 08:20

lol
I would not call it an apparatus. It was a temporary thing.
Yes, the bromine will come out hot therefore you need to have a long glass tube onto which you can apply some ice water in order to get some good condensation. I would personally use salt water with a temperature of -15 °C. It would cause some of the bromine to solidify/half liquid.

elementcollector1 - 25-8-2012 at 14:25

In that case, why wouldn't the collection flask be in a bath of icy salt water? The bromine would partially freeze, and that would get rid of some of the dangers of bromine. (Slushie of death.... MWAHAHAHA)

plante1999 - 15-5-2013 at 15:28

I just found a way to oxidize pool 30% NaBr to bromine easily, fast and without much smell

I was trying the oxidation of NaBr using nitric acid, I add added enough sulphuric acid to NaBr sol. to convert it all to HBr, then I added some HNO3 70%, nothing happened, even after few minutes, so I added about 0.2g of NaNO2, vigorous effervescence occurred, and within second, a pool of bromine was formed on the bottom.

15 ml 30% Bromide
5 ml 70% HNO3
5 ml 98% H2SO4
0.2g NaNO2

I need to work to make the reaction perfect for preparative scale.

[Edited on 16-5-2013 by plante1999]

chemcam - 15-5-2013 at 17:34

I'd like to try that out but what do you mean 15ml 30% bromide? Is your bromide in solution already when you buy it or do you have a specific weight on hand?

plante1999 - 15-5-2013 at 17:38

It comes on solution already for pool, don't forget to mix well, and if the solution only become bright red, add more HNO3.

chemcam - 15-5-2013 at 22:45

Are you sure that is exactly what you did? I duplicated this but on a 5x smaller scale. I didnt need nitrite, After the first drop of HNO3 it started vigorously bubbling and I mean VIGOROUSLY just like plante said! Spewing bromine vapor everywhere, and the liquid almost boiling out, I had to vacate the garage. In the end I did have a small amount of bromine that sunk to the bottom but I had to have lost 90% during the eruption. Like you said though the prep needs a little work lol.

-EDIT-
This pic is after I added water to calm everything down, only a couple drops of actual bromine at the bottom.
I think with some tweaking it will be a fast way to get low yield Br2.

[Edited on 5-16-2013 by chemcam]

woelen - 15-5-2013 at 23:23

I have done a lot of personal research on this reaction and I have found that if you have very pure bromide with very pure nitric acid, then no reaction occurs. As soon as a little NO2 is added, then suddenly the nitric acid oxidizes the bromide to bromine. Adding NO2 can be done by adding a pinch of NaNO2 (which decomposes in acid), but adding another reductor also helps (e.g. some copper metal, or some Na2SO3). The reductor reacts with the nitric acid and this gives some NO+NO2 and then the reaction with bromide sets off.

This is not a good way to make bromine. The bromine will be highly contaminated with ONBr. ONBr is a very dark liquid and it mixes well with bromine. Gaseous ONBr is brown, more chocolate brown than red-brown, but a mix of ONBr and Br2 is very difficult to distinguish from pure bromine, especially when viewed just by eye. An interesting experiment though may be Endimion17's infrared imaging of bromine, which contains a lot of ONBr. That would be a nice way of testing the purity of bromine, produced in this way (provided that ONBr is not 'colorless' in IR).

Endimion17 - 16-5-2013 at 01:30

I just might be able to do that, microscale. That's about the largest scale I can do, given the quantities of bromide salts I currently have. Therefore, distillation is out of the question.
Write a short, simple procedure and I'll see what I can do. You can expect results next week.

woelen - 16-5-2013 at 02:05

Take a few ml of concentrated nitric acid and put this in a test tube. Make a picture of this with your camera. Just to check whether nitric acid is 'colorless' in infrared as well as in visible light.

Add a ml of concentrated solution of NaBr or KBr (you could also make a picture of this before you add it to the nitric acid).

Add a pinch of NaNO2 or KNO2 and let the reaction run. If you don't have NaNO2 or KNO2 you could try to add some Na2SO3 or any other suitable colorless reductor which gives NO2 with nitric acid. Try to keep as much of the bromine in the test tube, such that a decent blob of liquid remains at the bottom of the test tube.

Make a picture after the reaction has completed. I am looking forward to your results.

Bezaleel - 16-5-2013 at 03:53

From a very distant past I seem to remember that adding KI solution to a solution of CuSO4 gives a redox reaction, releasing iodine that turns the solution brown. Addition of CCl4 and shaking a bit, indeed gave a purple drop of iodine dissolved in the CCl4 at the bottom of the testtube.

Is this a known reaction, or was there an issue with the purity of the chemicals? If it works with iodine, might it work with bromine too?

plante1999 - 16-5-2013 at 04:16

Copper II iodide is unstable, and disproportionnate to copper I and iodine, the bromide does not do that. It seac acid and it seams my nitric acid and bromide is quite pure then. Glad to ear that, my reaction was slow, and after NaNO2 addition some bubling was observed, but nothing very fast, the solution turned brown, then clear and a pool of bromine was formed.

With 120ml of 30% bromide I had got about 5-7ml of bromine.

woelen - 16-5-2013 at 05:26

Copper(II) iodide does not really disproportionate, it simply does not exist.

What you describe is a simple redox reaction:

2Cu(2+) + 2I(-) --> 2Cu(+) + I2

In the presence of excess iodide ion, a precipitate is formed as well:

Cu(+) + I(-) --> CuI

So, if you add copper(II) and iodide to each other you get a precipitate of CuI (which is off-white) and you get iodine.

Copper(II) is not a sufficiently strong oxidizer to oxidize bromide ion, so nothing similar happens with bromide.

@plante1999: The word 'disproportionate' has the meaning that one compound shows a redox reaction with itself, where one part acts as oxidizer and the other part acts as reductor.

A well-known example is 2Cu(+) --> Cu(2+) + Cu (disproportionation from +1 to +2 and 0)
Another well-known example is 3BrO(-) --> 2Br(-) + BrO3(-) (disproportionation from +1 to -1 and +5)

The reverse also exists. That is called comproportionation. In that case two compounds with the same element in different oxidation states react to form a single compound with oxidation state somewhere between that of the original reactants. An example is the reaction of bromate and bromide to elemental bromine in acidic solution (comproportionation from +5 and -1 to 0).

Whether disproportionation or comproportionation occurs depends on the redox potential for transfer from oxidation state to another. A Frost-diagram nicely shows whether this occurs or not (please use Google to find more about that subject).

plante1999 - 16-5-2013 at 08:23

I know what disproportionate mean...

Chlorine in NaOH, etc...

Simply that I messed up a little what I meant,. I always taugth copper II iodide was simply way too instable, at least thats what old chemistry books says. Iodide in copper II iodide is reductor, and copper II the oxidizer, in this way it is a disproportion reaction.

chemcam - 16-5-2013 at 08:42

Quote: Originally posted by woelen

I have done a lot of personal research on this reaction and I have found that if you have very pure bromide with very pure nitric acid, then no reaction occurs. As soon as a little NO2 is added, then suddenly the nitric acid oxidizes the bromide to bromine. (cut)

My acids are ACS grade but I used pool grade 99% bromide. I do have ACS grade bromide though, but only 125g so I was saving it, I will use that later and record it.

Quote: Originally posted by woelen

Take a few ml of concentrated nitric acid and put this in a test tube. (cut)

Add a ml of concentrated solution of NaBr or KBr (you could also make a picture of this before you add it to the nitric acid). (cut)

I am going to perform this in a couple hours or so and I will record a video. I notice though that you don't mention H2SO4 in this prep. Was it overkill that I used HNO3 and H2SO4 in my last run, could it have cause all the spewing? I imagine it got fairly hot. I don't have infrared on my camera but i'll show the reagents I used to prove purity.

[Edited on 5-16-2013 by chemcam]

woelen - 16-5-2013 at 12:19

@plante1999: Disproportionation is specific to a single element in a single compound to act as oxidizer and at the same time as reductor. So, Cl2 reacting to give chloride and hypochlorite is indeed disproportionation, but CuI2 decomposing to CuI and I2 is not disproportionation, it is a normal redox reaction, in which copper(II) is the oxidizer and iodide is the reductor.

@chemcam: It is the IR-part which I find most interesting, but of course, it is always nice to have other people perform the reaction. The presence of conc. H2SO4 certainly will add to the reactivity of the mix. It makes the nitric acid (more) anhydrous and this makes the acid much more reactive and may even lead to decomposition of some HNO3 with formation of NO and/or NO2, which could be the reason for the violent reaction in your case.

plante1999 - 16-5-2013 at 12:59

Sorry for disproportion misconception, but I'm sure I read something about CuI2 in a old book. As for H2SO4, my idea was to save on nitric acid, as the bromide need to be turned to hydrobromic acid for the reaction, nitric acid do the trick too, but for me, it is way more costly.

[Edited on 16-5-2013 by plante1999]

Eddygp - 16-5-2013 at 13:03

What if you just leave the test tube in an ice bath before adding the NaNO2? This would keep the temperature under control.
NOBr undergoes photolysis or something like that. Maybe flashing a really strong light from different points might help getting as much NOBr as you can out of the Br2.

chemcam - 17-5-2013 at 09:12

You guys were right, when I used my high grade sodium bromide, NaNO2 was required to get the reaction going.

Endimion17 - 4-6-2013 at 12:24

woelen, I'm sorry for the delay. Here are the results of the NIR camcorder imaging.

No nitrites are required for the initialization of the reaction. After you mix the solution of bromide ions with nitric acid, it takes a couple of minutes for the reaction to start. At first it really doesn't look like anything is happening, but you have to wait.
It is not violent. The solution turns yellow, then brownish, and then turbid brown. Bromine starts raining down, precipitating from a "cloud". Business as usual.
It obviously depends on the temperature and the concentration of reactants.

The NIR images were taken few days after the reactions were complete. In VIS, both solutions look like Coke, with the exception of the blob of bromine at the bottom, which is pitch black.

Nitrites kind of speed things up. I think more heat was given off, too, and the colour change was faster.
Maybe they're consumed, maybe it's catalysis, I don't know and won't even bother. This is a messy, expensive and hazardous way to oxidize the Br- anion.

[Edited on 4-6-2013 by Endimion17]

woelen - 4-6-2013 at 12:47

Thanks for doing this experiment. Unfortunately it is not conclusive about impurities in the bromine. There are two possibilities:
1) The impurities are 'colorless' in IR light.
2) There are no impurities, other than water and acid.

Could you try one other experiment? The experiment is fast and very simple and produces a lot of ONBr, one of the supposed contaminants in the production of Br2 from NaBr and HNO3.

The experiment goes as follows:
- Prepare a very concentrated solution of NaBr in water.
- Add approximately an equal volume of 50% H2SO4 (not nitric acid). Some crystals of NaHSO4 may form, just decant the liquid from these crystals and keep the liquid.
- Add some solid NaNO2 to the liquid.

If you have HBr, then things are even more simple. Just add some solid NaNO2 to 40% HBr.

You get a very dark liquid (nearly black) and a chocolate brown gas. This gas is ONBr, not Br2. It is very dense (you can pour it out easily). The solution in water is practically black. If the liquid is diluted, then the ONBr hydrolyses and the liquid becomes colorless or yellow, due to the formation of a little amount of Br2 as well.

This is the gas and liquid which I really would like to see in IR imaging. If ONBr is 'colorless' in IR-light, then things still are not conclusive, but if ONBr is dark in IR light as well, then your other experiment gives interesting information.

I really wish I could do this kind of experiments myself. I dare not take apart my 400 euro digital camera, however, for trying out the filter trick.

[Edited on 4-6-13 by woelen]

bfesser - 2-11-2013 at 11:52

Picked up an old book yesterday at a thrift store. Feel like transcribing something. Enjoy:

Quote:

<div align="center">CHAPTER 18 The Chlorine Family <hr width="70" />BROMINE</div>    Preparation of bromine. The laboratory method and the industrial method for preparing bromine are as follows:     1. Laboratory method. Just as chlorine is set free by the oxidation of hydrochloric acid by manganese dioxide (p. 203), so bromine may be prepared by a similar reaction, substituting hydrobromic acid for hydrochloric : <table style="float: right"><tr><td>

</td></tr></table> 4(H+, Br-

+ Mn++++, 2 O-- → Mn++, 2 Br- + 2 H2O + Br2 ↑ Hydrobromic acid is unstable and on this account is not usually available in the laboratory ; so a mixture of sodium bromide and sulfuric acid is used instead. The equation for the complete reaction is like the one for chlorine (p. 203) : 2(Na+, Br-

+ 2(2 H+, SO4--

+ Mn++++, 2 O-- → 2 Na+, SO4-- + Mn++, SO4-- + 2 H2O + Br2 ↑     Laboratory apparatus. The materials are placed in a retort, A, arranged as shown in Fig. 162. The delivery end of the retort dips slightly into the water in flask B, which is partly immersed in ice water (C

. As the contents of the retort are headed, bromine distills over and is collected in the cold receiver. – Introduction to College Chemistry. McPherson, Henderson, Fernelius, & Quill. Ginn and Company, 1942/46. (<a href="http://catalog.hathitrust.org/Record/009227908" target="_blank">Hathi Trust</a> <img src="../scipics/_ext.png" />

Perhaps a little easier to read:

4 HBr + MnO2 → MnBr + 2 H2O + Br2 ↑

2 NaBr + 2 H2SO4 + MnO2 → Na2SO4 + MnSO4 + 2 H2O + Br2 ↑

[Edited on 3.11.13 by bfesser]

elementcollector1 - 3-11-2013 at 11:19

A few notes from my most recent attempts:
-If the solution of bromine still looks red or even orange after distillation has supposedly 'ended' (water was coming over instead of bromine), distill it again. This roughly doubled my yield.
-Do NOT use a Kjehldahl bulb - the bromine gas hung around in the bulb forever, being heavier than air, and never distilled over at all. I was only able to recover it by disconnecting the setup and quickly attaching a flask full of sulfuric acid to one end and stoppering the other.

The original solution is still pale orange, so I'm going to try one more time and see what happens. Probably not much.
Right now, from 25g of sodium bromide, electrolysis at 2A for about 13 hours, and far too much sulfuric acid, it looks like I have about 1/3 of a milliter of bromine (considering the density of bromine of about 3.1 g/mL, this is just over 1g). Now, for ampouling, I'm going to redistill from the sulfuric acid I kept it in. My only worry is that if my condenser takes up so much volume, the bromine might well just evaporate to fill the space and I'll lose all my yield.

woelen - 4-11-2013 at 04:27

The problem with using such small quantities is that your mechanical losses (such as loss through vapor, filling up a flask) may bring down total yield very much. Mechanical losses tend to be independent of the amount used and losing e.g. 1 gram from 2 grams or losing 1 gram from 10 grams is a lot of a difference.

If you use ordinary glassware for your distillation (e.g. 100 or 250 ml flasks, a cooler of 40 cm length, NS24 or NS29 joints), then using only 25 grams leads to near total loss. For this purpose I also purchased a micro distillation set with 25 and 50 ml flasks and NS14 joints.

grayson - 3-4-2017 at 00:32

Sorry to revive this old thread, but I feel it worth mentioning that KMnO4 + NaBr + H2SO4 should be done in that order.

I had the acid and the NaBr reacted with excess acid. I was aware KMnO4 and H2SO4 create Mn2O7, and somewhat energetically, so I added the KMnO4 with great caution.

It was... energetic.

A few grains make for a lot of sparks. Also bromine, but seriously do it the other way. Chilled with ice. Not salt ice, regular ice. Made that mistake at one point too. Dry ice is right out.

Attachment: sparks.m4v (4.2MB)
This file has been downloaded 781 times

JJay - 3-4-2017 at 00:51

Wow, that must have been a mess to clean up!

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