Sciencemadness Discussion Board

Boron tribromide

floraljoy - 13-11-2007 at 00:47

I have not seen this subject mentioned on this site, hopefully i searched correctley...

Boron tribromide, an interesting compound with many uses some may find it easy to buy while others may find it impossible. All of this aside, has anyone had any first or second hand experience producing it?

The little investigation that i have done on this subject has revealed small product. further research on my part needs to be done, admitadley.

amphorous Boron, wich may or may be accessible to many (different subject for different thread) can be reacted with bromine ( same hoohah as boron) at an elevated temperature to make Boron tribromide.

Or boron carbide with bromine at 300 C.

ok sounds easy enough. hah! So, the target compound alone reacts explosivley with water. its nasty stuff as is bromine high temperatures and glassware with potential explosives in it.

I have found few to no writeups on the subject and nothing as far as detailed procedure.

so! if anyone has a source for a procedure or advise (besides dont make the damn stuff, i never siad i was making it so shush!) please post away! thanks for the info

Sauron - 13-11-2007 at 01:09

BBr3's preparation is well documented in Brauer (available in forum library) and in Inorganic Syntheses as referenced from Brauer.

The preparation from boron and Br2 requires 600-750 C.

The usual laboratory preparation is from AlBr3 and BF3.

The yield is 70% But few amateur chemists are going to want to work with neat BF3 at home or, be able to purchase even a lecture bottle of BF3.

There is an alternative prep from KBF4 and AlBr3. The yield is lower.

AlBr3 is available, not sure about KBF4. That is potassium fluoroborate (tetrafluoroborate) and out to be available. Anyway it can be prepared from boric acid, 40% HF (always a joy) and 5 N KOH soln w/ methyl orange indicator to show end point. 90% yield, procedure included in attachment. This calls for a Pt bowl but I think telfon will do.

I do not see why we can't just neutralize fluoboric acid with potassium carbonate?

I am attaching the Brauer pages including the section on boron purification through BBr3 which describes the apparatus required in Fig.238 for reaction of B and Br3. Nice if you have a couple meters of 10 mm bore quartz tube and a pair of tube furnaces of appropriate dimensions.

Frankly only the alternate procedure (albeit lower yielding) from AlBr3 and KBF4 seems appropriate for even an advanced amateur lab.

KBF4 is $50/Kg from Alfa, BBr3 is commercially available but $1 a gram in 250 g packaging (Alfa).

[Edited on 13-11-2007 by Sauron]

[Edited on 13-11-2007 by Sauron]

Attachment: BBr3.pdf (241kB)
This file has been downloaded 2076 times

not_important - 13-11-2007 at 04:01

Quote:
Originally posted by Sauron...
AlBr3 is available, not sure about KBF4. That is potassium fluoroborate (tetrafluoroborate) and out to be available. Anyway it can be prepared from boric acid, 40% HF (always a joy) and 5 N KOH soln w/ methyl orange indicator to show end point. 90% yield, procedure included in attachment. This calls for a Pt bowl but I think telfon will do.

I do not see why we can't just neutralize fluoboric acid with potassium carbonate?
...


Both LDPE and HDPE are resistant to HF of up to at least 75%, at temperatures under 70 C or so. For at least lower concentrations such as 40% I think the issue is simple heat softening of the plastic rather than any reaction or solvent effect. Variant preparations used gutta-percha, effectively trans-1,4-polyisoprene; polyolefine plastics not containing plasticisers seem to have similar chemical resistance.

As you said, working with HF is not to be lightly undertaken, there's several threads concerning that.

I suspect that KOH was used instead of K2CO3 in order to avoid dealing with effervescent hydrofluoric acid. Plop, plop, fizz, fizz, oh what a disaster it is...


Looked but could find the reference, but I remember a BCl3 / BBr3 prep that was based on making an iron-boron alloy using the Goldschmidt reaction, then hot tube reacting that with the halogen, followed by distillation to separate the boron trihalide. Not as fancy of a setup as in Brauer, but still a bit demanding especially considering the moisture sensitivity.



[Edited on 13-11-2007 by not_important]

Sauron - 13-11-2007 at 04:15

Unless there's a reason why not, I'd be nclined to bypass the HF completely and instead neutralize tetrafluoroboric acid with KOH or K2CO3 soln, in the cold, I have not examined this from a cost perspective yet, but it may well be that the hazards of HF outweigh the cost factor.

Aqueous HF + boric acid = fluoboric acid soln anyway. I guess it will depend on how much BBr3 someone wants to make. It is nothing special to look at, water white when pure, heavy (d >2.5) and very moisture sensitive. I would not advise seeing what it smells like.

Anyway, the other preps which require either a tank of BF3 or tube furnaces with quarts tubes or both are clearly impractical for 99.99% of us. This procedure is therefore the only game in town.

Here is the Inorganic Syntheses Vol.III prep of BCl3 and BBr3 by both the BF3 method and the KBF3 method, with clearer details than Brauer and, comparative yields.



[Edited on 13-11-2007 by Sauron]

Attachment: Pages from IS-3.pdf (443kB)
This file has been downloaded 1663 times

garage chemist - 13-11-2007 at 07:08

I have made KBF4 myself some time ago, following a preparation from an ebook called "Inorganic preparations" or similar.

I first dissolved boric acid in 20% aqueous HF (I had 40% HF at hand, but decided to dilute it because the reaction with boric acid is exothermic) in a simple PE jar, let it stand for a while with frequent swirling (now a solution of HBF4), and added K2CO3 solution dropwise, again with swirling.
KBF4 is sparingly soluble, like KClO4, and precipitated in excellent yield. It was filtered and washed with cold water.

HF should be used in slight excess to avoid contamination of the KBF4 with KBF3OH (also sparingly soluble, HBF3OH slowly forms from HBF4 in neutral or basic solution). The theoretical amount of K2CO3 should be used in regard to the amount of boric acid used.

The KBF4 crystals have about the same index of refraction as water and therefore are almost invisible- it will look as if the solution has gelled when the KBF4 has precipitated.

One could use aqueous HF of even lower concentration (10% would surely work just as well), or even a solution of an alkali fluoride in HCl or dilute H2SO4 if no HF is available.
Purchasing HBF4 could also be possible, though I dont know about availability and price.

I have also seen KBF4 being available from chemical suppliers, and not very expensive.


Sauron, very interesting that BCl3 and BBr3 can be prepared from KBF4 and the aluminium halide! I have plenty of AlCl3 at hand and the KBF4 sample that I made myself. I could already make BCl3 with what I have, didnt know that until now!
I could make AlBr3 from the elements, although it is somewhat of a hassle.

The method with amorphous boron and chlorine or bromine in a quartz tube at 600°C sounds like a good alternative for those with no access to soluble fluorides. A tube furnace that only has to go to 700°C (hot enough for ketene production as well) is very easy to build, and you can use nichrome wire which is easier to work with (you can use waterglass as binder for embedding pastes) than Kanthal.

The problem I see is the preparation of B2O3 for the Mg reduction- the dehydration of boric acid is not trivial at all. If you dont have a seed crystal of crystalline (!) B2O3, you will not be able to get the water content (i.e. boric acid content) below a value of, I think, 5-12% (dont remember exact value), even if you heat to 1000°C!
The glassy B2O3-B(OH)3 melt that results from this is also extremely hard to crush, let alone turn into powder.
When I tried this in a small silver crucible I was completely unable to get the solidified melt out of the crucible in the first place, and had to abandon the experiment by dissolving it with water.
If you have a seed crystal, it is easy- you have to seed the melt at a certain temperature, and it rapidly gives off all water and you have crystalline pure B2O3. Brauer details this.
The crystallization of B2O3 from the glassy melt without the use of a seed crystal requires keeping it molten for weeks- might be worth doing when considering that you only have to do it once to get seed crystals.

Perhaps you can get away with using the B2O3-B(OH)3 glass in the Mg reduction if you increase the amount of Mg accordingly for H2 as additional byproduct (Mg readily reduces water at elevated temperature).

Brauers method for B2O3 that uses a drying pistol at 200°C with high vacuum and P2O5 as drying agent is too wasteful of valuable P2O5 for my taste.

An industrial method for B2O3 production involves heating anhydrous borax (sodium tetraborate) with concentrated H2SO4 at 800°C, phase separation of the B2O3 and Na2SO4 melts then occurs (B2O3 is the upper phase).
This gives amorphous B2O3 glass of about 96% purity.
The material of the vessel would be an issue- molten B2O3 attacks porcelain and glass, and hot concentrated H2SO4 attacks a lot of metals.


Gaseous BF3 is significantly easier to make than the other boron halides since its hydrolysis is reversible, contrary to BCl3 and BBr3 whose hydrolysis is irreversible.
BF3 also doesnt attack glass- no special apparatus required.

The thermal decomposition of KBF4 (with added B2O3 as catalyst) is one example (see Brauer for this). Probably the easiest method, but again, you need B2O3, and it should better be anhydrous in this application.
I have tried heating KBF4 alone and have not been able to get any useful amount of BF3 (fumes strongly in air) even at dull red incandescence. The B2O3 catalyst seems to be required.

BF3 can also be prepared from H2SO4, CaF2 and anhydrous borax. B2O3 and HF form and react in situ, an excess of H2SO4 binds the byproduct water. This mixture would have to be heated in a vessel resistant to HF in order to avoid SiF4 contamination in the BF3 gas. I have no details on this method, you may want to search the patent literature, it is an old industrial method.
Some HF could escape unreacted -an excess of borax would help against that- and would have to be removed from the gas stream.

Industry uses reaction of anhydrous borax with HF and H2SO4- for lab use, if HF or an alkali fluoride is available, it would be better to make HBF4 or KBF4 from it and go from there.

Perhaps the most appealing method for BF3 production in the laboratory would be the thermal decomposition of an aromatic diazonium tetrafluoroborate (Balz-Schiemann reaction).
Diazotise e.g. aniline in aqueous HCl with NaNO2 and add HBF4 solution to precipitate the sparingly soluble benzenediazonium tetrafluoroborate.
Filter, wash and thoroughly dry this salt, and heat it carefully in small batches to decomposition:
C6H5-N2*BF4 ---> C6H5-F + BF3 + N2
Fluorobenzene is obtained as byproduct, and BF3 in N2 after condensation of the fluorobenzene.
You can make other fluoroaromatics this way as well- there was a thread on preparation of 4-fluorobenzaldehyde some time ago, and I suggested Balz-Schiemann reaction with p-toluidine as the key step.
Brauer has a method for BF3 using this reaction, I think- you can also find examples of it on orgsyn.

BF3 diethyl etherate is commercially available and a convenient storage form of BF3, but I know of no method to get gaseous BF3 back from it (it distills unchanged).

[Edited on 13-11-2007 by garage chemist]

[Edited on 13-11-2007 by garage chemist]

not_important - 13-11-2007 at 07:25

The porous B2O3 in Brauer is also in Inorganic Syntheses #2, pages 22-23. The checkers found that a stream of dry air could be used in place of the drying pistol + vacuum. Molecular sieves can be used in place of P2O5/P4O10, the drying agent really isn't needed if you keep pumping.

I read a student dissertation in which they combined the two to make the B2O3 they needed - a slow stream of dry air until the oven had been slowly ramped up to 140C, then closing off the air intake and just pulling a decent vacuum until 185-195 C and holding it there for several hours. As with the original Brauer prep this gives a porous, slightly sintered oxide that is easy to crush.

garage chemist - 13-11-2007 at 07:34

What would be "decent vacuum"? I only have an aspirator, which has served me perfectly for all vacuum distillations until now.

Sauron - 13-11-2007 at 07:35

Yes, thanks, gc, I just reviewed Brauer on BF3 before reading your post and it does look like a BF3 generator to beed AlCl3 or AlBr3 would be practical, especially the H2SO4 method. KBF4 is only $48/Kg from Alfa (twice that from Acros, so shop around.) At that price it is hardly worth preparing one's own.

The yield from BF3 and AlCl3 or AlBr3 is about twice as good as from KBF4 and the Al halides so if you want to make a lot of it then it is probably worth setting up a BF3 generator. Also handy for preparing BF3 complexes.

What use is BBr3, I still do not know! But BF3 and BCl3 very handy indeed.

Sauron - 13-11-2007 at 07:45

There are some organic preparations that call for rather large amounts of BF3. For example the preparation of acetylacetone from acetone and acetic anhydride calls for 500 g Ac2) and 500 g BF3 neat in order to end up with a yield of little more than 160 g acetylacetone (2,4-pentanedione). As far as I am concerned that is a "buy it don't make it" situation in spades! And acetylacetone is cheap.

Fortunately more often than not BF3 and BCl3 are catalysts and such issues do not arise.

Turns out boron tribromide is more useful than I thought.

Attached, extracts from Paquette's books on all four boron halides.

[Edited on 15-11-2007 by Sauron]

Attachment: boronhalides.pdf (272kB)
This file has been downloaded 32556 times

LSD25 - 2-2-2008 at 05:27

Does somebody want to grab this:

http://jfs.sagepub.com/cgi/content/abstract/17/5/355

It is a study on the formation of boron trichloride from zinc borate and burning PVC

I was searching for 'Boron Carbide Chlorine' because I have seen some patents on the production of BCl3 from boron carbide with chlorine and was trying to determine whether the boron carbide could be heated (to 7-800C) with a chlorine donor. By the look of the study (above) it can.
Quote:
Journal of Fire Sciences, Vol. 17, No. 5, 355-361 (1999)
DOI: 10.1177/073490419901700502
© 1999 SAGE Publications
Flame Retardancy Behavior of Zinc Borate
Youming Yang

The Research Institute of Engineering Chemistry & Metallurgy The Central South University of Technology Changsha 410083, P.R. China

Xichang Shi

The Research Institute of Engineering Chemistry & Metallurgy The Central South University of Technology Changsha 410083, P.R. China

Ruirong Zhao

The Research Institute of Engineering Chemistry & Metallurgy The Central South University of Technology Changsha 410083, P.R. China

The effect of boron chloride formation from zinc borate in burn ing PVC was elucidated. By thermodynamic calculation, experiment, and infra red spectrum analysis, the conditions for boron halide formation when PVC was burned were addressed. It is found that when boron halide is produced, the B2O3 glass layer is destroyed and boron is volatilized, which is unfavorable to flame retardancy.

Key Words: zinc borate • fire retardant • polyvinyl chloride


Seems a whole lot easier to do at home than fucking with BF3:D Although the toxicity of burning PVC ain't trivial.

PS Seems like there is a whole lot of current interest in the preparation of porous carbons from the treatment of boron carbide with chlorine at elevated temperatures. I'll endeavour to find those articles.

aaparatuss - 28-12-2012 at 19:31

Are there any updates on the use of boron carbide in the reaction with bromine to produce the tribromide.

I saw one google book reference that stated temperatures of 800-1000 C are needed.

What type of device is needed to attempt the tribromide synthesis?

Thanks

yes, fumes hoods and danger. Im simply wondering what to do with all this boron carbide. ballistic panels or thermoelectric generators are a possibility, but i would like to ponder a tribromide attempt in the future.

Melgar - 29-5-2017 at 07:44

Quote: Originally posted by aaparatuss  
Are there any updates on the use of boron carbide in the reaction with bromine to produce the tribromide.

I saw one google book reference that stated temperatures of 800-1000 C are needed.

What type of device is needed to attempt the tribromide synthesis?

Thanks

yes, fumes hoods and danger. Im simply wondering what to do with all this boron carbide. ballistic panels or thermoelectric generators are a possibility, but i would like to ponder a tribromide attempt in the future.


Yep. I did it. Tiny amounts, but I did it. I took a piece of pyrex tubing, and twisted it so there was a thin tube tapering off, added boron carbide powder, added liquid bromine, then capped the opening with glue. Bromine will not create very much pressure when heated to its boiling temperature, and besides, it had plenty of space in the tube. So I essentially counted on the hydrostatic pressure of bromine, as well as bromine's inertia when it reached boiling point, to pressurize the tip of the tube enough that the contents would reach combustion temperature in at least the tip, when it was heated in a flame.

Indeed, when the tip of the tube was heated, sparks could be seen as the bromine boiled, much like you can see sparks when you throw a fleck of aluminum into liquid bromine.

I'm not sure if this is practical at all, and I have no way of knowing for sure that I have boron tribromide, but I did get boron carbide to react with bromine, which is the first step.

[Edited on 5/29/17 by Melgar]

IMG_20170529_112846.jpg - 254kB

nleslie321 - 19-2-2020 at 04:50

Did anyone get anywhere with the bromine and boron carbide at 800 C ?
I was going to try it in a 1200mm quartz tube thats sealed at one end like a really really long test tube. Then wrap the bottom of the tube in elements and wrap in heat blanket. This way i can seal the top and put like about 50 to 100ml of bromine in the bottom with boron carbide and heat the element up to around 800 - 1000 C.