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theAngryLittleBunny
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Making bromine (super cheap and easy, no destillation, quantitative yield)
Hey there, so I just discoverd something so amazing that I made an just to share this. So I saw how people make bromine by oxidation of bromide with
KMnO4, H2O2 or whatever and distill it, or through the proportionation of 1 mole bromate and 5 moles bromide, and all of this is a lot of effort. But
it turns out that you actually can oxidize bromide with nitrate! O.o So that's how I did it, I put 25g of NaBr and 12g of KNO3 in a beaker, and
dissolved this in about 100mL of water. Then I slowly poured in about 20mL of 95% H2SO4. On addition, it starts to fizz and makes red fumes. After I
added all the H2SO4, I put a round bottom flask filled with cold water onto it and let it react for like 10 minutes, and after that, a blob of bromine
collected on the bottom of the beaker! I'am just so hyped about this, because I didn't find this methode anywhere, but it's so easy, has it
quantitative yields, and the reagents are super cheap!
The nitrate also seems to oxidize 5 bromide ions, not just one, because there was no NO2 gas produced, like it would be when the nitrate make a one
electron oxidation, neither did I use enough KNO3 in order for this to work, I used a molar ration of 1:2 KNO3:NaBr.
And the nitrate had to oxidize all the bromide, because if it wouldn't, the bromine would have been in solution as the tribromide ion, and therefore
wouldn't have collected to a blob.
I hope someone finds this as exciting as I do, because not only is bromine just a super strange and cool element, but also a super useful reagent for
organic chemistry (I need it for bromoacetyl bromide). But I'am just wondering, why does no one seem to have known about this? I can't find anything
about this on youtube, and there are tons of youtube videos that show how to make bromine.
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gdflp
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I would read this thread. While this method does produce bromine, it's contaminated with other compounds, namely nitrosyl bromide. This is similar methods of
producing bromine involving Cl<sub>2</sub>, while they are fairly simple, they result in BrCl contamination which is why the
H<sub>2</sub>O<sub>2</sub> and permanganate methods are preferred.
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AJKOER
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Try this very available path that I surmised should theoretically work yielding bromine water and other contaminants.
To aqueous NaBr add a bit of iron rust, a piece of copper, some filtered lemon juice and treat with air from an air pump. Sunlight exposure should
also help.
A rough description of the chemistry (for details, see prior comments at http://www.sciencemadness.org/talk/viewthread.php?tid=70055#... ), in situ generation of reactive oxygen species, including H2O2. The latter feeds
a Fenton redox reaction, which cycles (with the presence of sunlight and citrate/ascorbate) producing hydroxyl radicals (.OH). The latter radical
turns HBr into .Br then Br2 (from .Br + .Br, also the radical anion .Br2- from .Br + Br-, see, for example, comments available at http://pubs.acs.org/doi/abs/10.1021/jp021498p ), and possible some bromine oxides also in the presence of BrO-, HOBr and/or O2 (see, for example,
https://www.jstor.org/stable/2416051?seq=1#page_scan_tab_con... ) and heating. Should also mention the stable tribromide complex, Br3- (from Br2 +
Br-, see, for example, http://scholarcommons.sc.edu/cgi/viewcontent.cgi?article=102... ).
Some nasty bromo-citrates and bromate formation possible also, so not a safe experiment. Performing outdoors with safety measures highly recommended!
Note, escaping bromine compounds can likely, via chain reaction in the upper atmosphere, do significant damage to the ozone layer (more potent than
chlorine, see http://people.oregonstate.edu/~muirp/stratozo.htm ).
Likely low yield in separating out the bromine from, for example, Br3- or even the transient radical anion .Br2-, however BOTH of which may be
converted to Br2 (or Br-) by .HO2, via (see Table 3 at http://albeniz.eng.uci.edu/dabdub/My_papers/2004_Hunt-et-al_... ) Br3- + .HO2 --> Br2- + Br- + H+ + O2, and .Br2- + .HO2 --> Br2 + HO2- (or
2 Br- + O2 + H+, see http://onlinelibrary.wiley.com/doi/10.1029/2003GL018572/full... ), which is another task, but at least, you started with low cost available
reagents.
Only do this experiment if you have prior experience with bromine and like it. While I enjoy working with Cl2, I now avoid Br2 .
[Edited on 7-3-2017 by AJKOER]
[Edited on 7-3-2017 by AJKOER]
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theAngryLittleBunny
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Quote: Originally posted by gdflp  | | I would read this thread. While this method does produce bromine, it's contaminated with other compounds, namely nitrosyl bromide. This is similar methods of
producing bromine involving Cl<sub>2</sub>, while they are fairly simple, they result in BrCl contamination which is why the
H<sub>2</sub>O<sub>2</sub> and permanganate methods are preferred. |
Nitrosyl bromide.....interesting point, didn't think about that. But this stuff hydrolyses in water: BrNO + H2O -> HBr + HNO2, so it should be
fairly easy to detect, since the HBr makes tribromide ions with the Br2, which should colour the solution quite dark. But if you use a large excess of
KNO3 and H2SO4, the HBr should immediately get oxidized to Br2 again. I mean this is only my speculation, but I think if you would wash the bromine
with a mixture of HNO3 + H2SO4, the nitrosyl bromide should be destroyed after a while, and I think you could do this fairly easily be using a large
excess of HNO3 + H2SO4.
| Quote: Originally posted by AJKOER | Try this very available path that I surmised should theoretically work yielding bromine water and other contaminants.
To aqueous NaBr add a bit of iron rust, a piece of copper, some filtered lemon juice and treat with air from an air pump. Sunlight exposure should
also help.
A rough description of the chemistry (for details, see prior comments at http://www.sciencemadness.org/talk/viewthread.php?tid=70055#... ), in situ generation of reactive oxygen species, including H2O2. The latter feeds
a Fenton redox reaction, which cycles (with the presence of sunlight and citrate/ascorbate) producing hydroxyl radicals (.OH). The latter radical
turns HBr into .Br then Br2 (from .Br + .Br, also the radical anion .Br2- from .Br + Br-, see, for example, comments available at http://pubs.acs.org/doi/abs/10.1021/jp021498p ), and possible some bromine oxides also in the presence of BrO-, HOBr and/or O2 (see, for example,
https://www.jstor.org/stable/2416051?seq=1#page_scan_tab_con... ) and heating. Should also mention the stable tribromide complex, Br3- (from Br2 +
Br-, see, for example, http://scholarcommons.sc.edu/cgi/viewcontent.cgi?article=102... ).
Some nasty bromo-citrates and bromate formation possible also, so not a safe experiment. Performing outdoors with safety measures highly recommended!
Note, escaping bromine compounds can likely, via chain reaction in the upper atmosphere, do significant damage to the ozone layer (more potent than
chlorine, see http://people.oregonstate.edu/~muirp/stratozo.htm ).
Likely low yield in separating out the bromine from, for example, Br3- or even the transient radical anion .Br2-, BOTH of which may be converted to
Br2 (or Br-) by .HO2, via (see Table 3 at http://albeniz.eng.uci.edu/dabdub/My_papers/2004_Hunt-et-al_... ) Br3- + .HO2 --> Br2- + Br- + H+ + O2, and .Br2- + .HO2 --> Br2 + HO2- (or
Br- + O2 + H+, see http://onlinelibrary.wiley.com/doi/10.1029/2003GL018572/full... ), which is another task, but at least, you started with low cost available
reagents.
Only do this experiment if you have prior experience with bromine and like it. While I enjoy working with Cl2, I now avoid Br2.
[Edited on 7-3-2017 by AJKOER]
[Edited on 7-3-2017 by AJKOER] |
Wow, this stuff is a little bit beyond me honestly O.o So I can't say too much to it (I'am only 19 by the way). But it sounds interesting, but I've
never really worked with or included UV light in chemistry. I think I could try this, but I don't think it's a viable methode of making bromine
though, would still be interesting to know if it would work.
And don't worry, I worked a lot with bromine, I think it a really lovely element actually ^-^, much more then chlorine, simply because it's a liquid
and therefore much easier to handle. The only big disadventage of bromine compared to chlorine is that the atom is twice as heavy, therefore you have
to use twice as much for the same reaction compared to chlorine.
[Edited on 7-3-2017 by theAngryLittleBunny]
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DraconicAcid
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That's okay- the stuff that AJKOER suggests is a bit beyond most of us, especially those of us with a solid understanding of chemistry (both
theoretical and hands-on).
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
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AJKOER
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DraconicAcid:
Correction, you must mean your momentarily lapse of understanding in chemistry, as I point to my use of the capitalized word 'BOTH' with references
citing, what I believe, is the innovative utilization of .HO2 radical as a possible low cost aid in the recovery of bromine from the problematic
formation of both the Br3- complex and the .Br2- radical anion.
Please check out the chemistry, which is displayed in Table 3.
Interestingly, such halogen complexes apparently occur in various preparation paths for other halogens as well.
This is, of course, requires experimentation with and knowledge of possible convenient/low cost method of introducing the hydroperoxyl radical in
likely an alternate, but related, preparation path.
----------------------------------------------
[Edit] Here is a possible embodiment using UV and/or solar light or pulse radiation, an acid, NaBr, H2O2, to which is added N2O (requires solar light)
or another photo active agent producing hydroxyl radicals (like TiO2 which uses UV as also does H2O2).
For the acid, I would first experiment with aqueous CO2 (carbonic acid) noting yields, and then move up to stronger acids.
To mitigate the decomposition of H2O2 by NaBr before the formation of .OH radicals, especially convenient for the N2O pulse radiation treatment in a
microwave, I would suggest dropping the NaBr into the system as a solid mass just before pulse irradiation.
Thheoretically, the creation of hydroxyl radical from H2O2 in the presence of UV light is straight forward:
H2O2 + UV = .OH + .OH (see http://pubs.acs.org/doi/abs/10.1021/jp100204z )
but this embodiment follows an enhanced UV/H2O2/TiO2 advance oxidation process (AOP) or H2O2/N2O/irradiation by solar light or microwave pulse
radiation, both resulting in increased hydroxyl radical generation (the TiO2 AOP also creates superoxide anions).
The hydroperoxyl radical can be created via:
.OH + H2O2 = .HO2 + H2O (see above and also http://aip.scitation.org/doi/abs/10.1063/1.2943315?journalCo... )
An embodiment employing an excess of .HO2 formation (from employing more H2O2 and TiO2 in UV light, or with H2O2/N2O in solar light or pulsed
radiation in a microwave, all adding to hydroxyl radical formation), hopefully increases yield of bromine by reversing bromine complex formation.
Expected reactions of interest, per the previously cited reaction from Table 3:
Br3- + .HO2 --> .Br2- + Br- + H+ + O2
.Br2- + .HO2 --> Br2 + HO2-
This last reaction is particularly compelling as in the presence of sufficient acid (supply of H+), we have regenerated some of the H2O2 (or .HO2
absence H+):
H+ + HO2- = H2O2
HO2- + .OH = .HO2 + OH-
[Edited on 8-3-2017 by AJKOER]
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unionised
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Lemon juice contains vitamin C which reduces bromine back to bromide.
The same is true of glucose and other things.
So, as is often the case, Ajkoer has come up with something very complex; and wrong.
Eventually, you would oxidise all of those things- including the citric acid. You might even get some bromine.
But it's never going to work well.
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AJKOER
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Good point Unonized, the citrate/ascorbate are indeed biological reductants. However, there do not need to be added in large amounts, just in amounts
corresponding to the 'bits' (which is in respect to the much larger amount of bromide) of the iron oxide and copper. Ideally, one should run the
Fe2O3/Cu/Citric acid/O2 system first for a while to accumulate ions and then massively dilute by addition of the aqueous NaBr and continue the
aeration.
Citrate/ascorbates are employed as complexing agents to extend the pH range of the iron Fenton and copper Fenton-like reactions to near neutral range.
They further help in recycling Fe(lll) to Fe(ll). That function can also be accomplished by light and in the presence of hydroxyl radicals.
However, please note that the existing small amounts of citrate/ascorbate ions are also consumed in the redox cycling reaction as hydroxyl radicals
are created. The latter provides fast reaction paths to bromine.
Further, please note from the Table 3 reference, that the reaction rates for complex formation are very fast:
A6: Br2 + Br- --> Br3- k = 1.5xE09
which would largely remove any bromine (as Br3-) before the remaining citrate could likely attack the bromine.
So, you may have a point, but it is, in my opinion, not likely large, as I understand the mechanics of the reaction system and my related direct
experience.
The primary interesting point of my suggested preparation was to eliminate the need of H2O2.
[Edit] Note, I have also provided an embodiment above that does actually uses H2O2 enhanced by TiO2 in UV, or H2O2/N2O/irradiation by solar light or
pulse radiation, both absent citrates.
[Edited on 8-3-2017 by AJKOER]
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JJay
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The first time I made bromine, many years ago while I was taking general chemistry in college, in my dorm, I used hydrogen peroxide, hydrochloric
acid, and sodium bromide. I thought the yield was awful and had no idea how pure it was. Distilling it was out of the question; I had a distillation
apparatus, but I would have no sooner distilled nerve gas. I ended up neutralizing the bromine I made and buying some in tiny 25 gram vials for
$1/gram.
[Edited on 7-3-2017 by JJay]
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symboom
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That is an interesting reaction though
Nitric acid formed it situ
Then protinating the potassium nitrate to nitric acid
Nitric acid reacting with potassium bromide forming
nitrogen dioxide gas and liquid bromine
Alot of these reactions become beneficial when the side products such as nitrogen dioxide can be used in another reaction.
Sodium bromide and (oxone) potassium monopersulfate forms bromine liquid
Just mix the two solids and add water
Oxone is avalible as non chlorine shock for pools
Works with potassium iodide too formine iodine solid
And sodium chloride forming chlorine gas
Quick and easy :-)
[Edited on 8-3-2017 by symboom]
[Edited on 8-3-2017 by symboom]
[Edited on 8-3-2017 by symboom]
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theAngryLittleBunny
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| Quote: Originally posted by JJay | The first time I made bromine, many years ago while I was taking general chemistry in college, in my dorm, I used hydrogen peroxide, hydrochloric
acid, and sodium bromide. I thought the yield was awful and had no idea how pure it was. Distilling it was out of the question; I had a distillation
apparatus, but I would have no sooner distilled nerve gas. I ended up neutralizing the bromine I made and buying some in tiny 25 gram vials for
$1/gram.
[Edited on 7-3-2017 by JJay] |
The yield was probably shit because most of the bromine dissolved because of the chloride ions that were present. I tried this where I made a solution
of NaCl, and put bromine into it, and it dissolved. I guess it forms a chlorodibromide ion (Br2Cl-), but I didn't really find anything aboug this.
| Quote: Originally posted by symboom | That is an interesting reaction though
Nitric acid formed it situ
Then protinating the potassium nitrate to nitric acid
Nitric acid reacting with potassium bromide forming
nitrogen dioxide gas and liquid bromine
Alot of these reactions become benifityal when the side products such as nitrogen dioxide can be used in another reaction.
Sodium bromide and (oxone) potassium monopersulfate forms bromine liquid
Just mix the two solids and add water
Oxone is avalible as non chlorine shock for pools
Works with potassium iodide too formine iodine solid
And sodium chloride forming chlorine gas
Quick and easy :-)
[Edited on 8-3-2017 by symboom]
[Edited on 8-3-2017 by symboom] |
I've never seen that stuff being sold here in Austria, here, pretty much only trichloroisocyanuric acid or sodium dichloroisocyanurate is sold. But if
KHSO5 works, H2O2 + H2SO4 should work as well.
And by the way, even if you would use HNO3 instead of KNO3, you would still need H2SO4, since you need to protonate EVERYTHING, also the bromide ions.
[Edited on 8-3-2017 by theAngryLittleBunny]
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byko3y
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The reason why sole hydrogen peroxide does not work for bromine production is because bromide-bromine pair catalytically decomposes hydrogen peroxide,
just as well as iodide-iodine or Fe(III)-Fe(II) does.
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theAngryLittleBunny
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| Quote: Originally posted by gdflp | | I would read this thread. While this method does produce bromine, it's contaminated with other compounds, namely nitrosyl bromide. This is similar methods of
producing bromine involving Cl<sub>2</sub>, while they are fairly simple, they result in BrCl contamination which is why the
H<sub>2</sub>O<sub>2</sub> and permanganate methods are preferred. |
Okay, so I tried to make bromoacetly bromide now, and I have to say, gdflp, this person knows his shit. So I put red phosphorus and acetic acid in a
molar ration of 1:9 in a beaker, and then added a 4.5 molar ratio of the bromine I made with the nitrate dropwise. The idea was that the first third
of the Br2 would react with the P to form PBr3, the PBr3 then would react with the acetic acid to form acetyl bromide while also catalysing the so
called Hell-Volhard-Zelinsky halogenation, where the alpha carbon of a barboxylic acid is halogenated, which would use up the other two thirds of the
bromine. When I dropped it in, red/brown fumes appeard, and I kinda thought "hmm, that kinda smells like nitric acid.....wait....nitric acid? O.o"
Well, turns out, these fumes are NO2 gas! So there must be a nitrogen compound in the bromine, which most likely is nitrosyl bromide, I mean it's the
only stabile nitrogen bromine compound I can think of, and there is also nitrosyl chloride in aqua regia, and since I just made the bromine equivalent
of that, so...BrNO sounds reasonable.
I actually left bromine in the water for very long, in the hopes the BrNO would be destroyed through that, but it obviously wasn't ._.
I decided to not continue, since I planned to distill off the bromoacetyl bromide, but I don't wanna blow myself up because it might have formed some
weird explosives like acetyl nitrate, you never know.
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Eddygp
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Nitrosyl bromide itself would be interesting as a reagent to be fair.
there may be bugs in gfind
[ˌɛdidʒiˈpiː] IPA pronunciation for my Username |
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unionised
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| Quote: Originally posted by AJKOER | Good point Unonized, the citrate/ascorbate are indeed biological reductants. However, there do not need to be added in large amounts, just in amounts
corresponding to the 'bits' (which is in respect to the much larger amount of bromide) of the iron oxide and copper. Ideally, one should run the
Fe2O3/Cu/Citric acid/O2 system first for a while to accumulate ions and then massively dilute by addition of the aqueous NaBr and continue the
aeration.
Citrate/ascorbates are employed as complexing agents to extend the pH range of the iron Fenton and copper Fenton-like reactions to near neutral range.
They further help in recycling Fe(lll) to Fe(ll). That function can also be accomplished by light and in the presence of hydroxyl radicals.
However, please note that the existing small amounts of citrate/ascorbate ions are also consumed in the redox cycling reaction as hydroxyl radicals
are created. The latter provides fast reaction paths to bromine.
Further, please note from the Table 3 reference, that the reaction rates for complex formation are very fast:
A6: Br2 + Br- --> Br3- k = 1.5xE09
which would largely remove any bromine (as Br3-) before the remaining citrate could likely attack the bromine.
So, you may have a point, but it is, in my opinion, not likely large, as I understand the mechanics of the reaction system and my related direct
experience.
The primary interesting point of my suggested preparation was to eliminate the need of H2O2.
[Edit] Note, I have also provided an embodiment above that does actually uses H2O2 enhanced by TiO2 in UV, or H2O2/N2O/irradiation by solar light or
pulse radiation, both absent citrates.
[Edited on 8-3-2017 by AJKOER] |
You are still talking bollocks.
The reaction Br2 + Br- --> Br3- is fast and has a high eqm constant.
But it's reversible.
In effect you are saying that you can't titrate iodine in the presence of iodide but everyone knows that's nonsense.
Also if you add oxygen to a solution of potassium bromide you will produce some bromine.
But you don't remove the potassium.
And, if you look at the ions and atoms that are present, you must make potassium hydroxide.
And that will build up until it is at a high enough concentration to react with the bromine you are trying to make- .
You need to add an acid.
However the only one you have suggested is citric- that's not much of an acid, but it might do the job- if it wasn't for all the ascorbate, glucose
etc that you are proposing to add along with it.
This is another of your pointless suggestions where, when someone points out that you are talking nonsense you just add more nonsense on top.
Why not stop?
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AJKOER
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| Quote: Originally posted by byko3y | | The reason why sole hydrogen peroxide does not work for bromine production is because bromide-bromine pair catalytically decomposes hydrogen peroxide,
just as well as iodide-iodine or Fe(III)-Fe(II) does. |
Actually, keeping the amount of H2O2 small in contact with the solution containing acidified bromide, and any bromine released therefrom, can be made
to work. See, for example, U.S Patent 4,029,732, link: https://www.google.com/patents/US4029732 ,"Preparation of bromine", where apparently Br2 is, to quote, "prepared by a method which comprises
contacting hydrogen peroxide with an aqueous solution containing bromide ion and rapidly removing the bromine as it is formed."
Also, may suggested use of photolysis of bromide involving TiO2 has been investigated. See http://bard.cm.utexas.edu/resources/Bard-Reprint/363.pdf . The paper cites the use of platinized TiO2 and aeration with oxygen. The author cites
the formation of surface intermediates H2O2 and .OH radical, and an increase in bromine formation as the pH was lowered (more acidic conditions).
My original method also attempted to incorporate Fenton and Fenton-type reactions to boost the photolysis products. Some support for this from an
article, "Formation of Organobrominated Compounds in the Presence of Bromide under Simulated Atmospheric Aerosol Conditions", by Davide Vione Dr., et
al., 2008, that employs light, Fe(lll) and likely air, and also, H2O2 in the dark (Fenton), at low pH. To quote from the abstract:
"Photobromination of phenol takes place upon UV/Vis irradiation of FeIII and bromide under acidic conditions, and most likely involves the brominating
agent Br2−.. Bromination is also observed in the presence of nitrate and bromide under UV irradiation, most likely involving Br2−. formed upon
oxidation of bromide by .OH. Moreover, quantitative bromination of phenol is observed in the dark in the presence of hydrogen peroxide and bromide.
This process is strongly favored under acidic conditions, but a residual, pH-independent bromination pathway is also present. The rates and yields of
bromination (up to 100 %) are considerably higher than those reported for chlorination under comparable conditions, suggesting that the higher
activity of bromine species could compensate for the lower concentration of bromide ions in aerosol compared to chlorides. The reported processes are
potential sources of reactive bromine species (Br2−., HBrO) and aromatic bromo derivatives in atmospheric aerosols, in particular after the
acidification process linked with aerosol aging."
Link: http://onlinelibrary.wiley.com/doi/10.1002/cssc.200700031/ab...
Note, the photolysis of aqueous nitrate is a source of hydroxyl radicals and is interestingly selected here, in the presence of bromide, over the
known superior photocatalytic properties of nitrite.
My further look at Table 3, listing bromine reactions, (link: http://albeniz.eng.uci.edu/dabdub/My_papers/2004_Hunt-et-al_... ), suggests that methods forming, under acidic condition (pH < 4.8), in situ
H2O2 and associated .HO2 by liberating lots of hydroxyl radicals, are likely a good path. The latter .OH can form the .Br monoatomic radical (and
water) from HBr, which can feed further reactions like:
A44 .Br + .Br --> Br2
A55 .Br + H2O2 --> Br- + .HO2 + H+
A41 .Br + Br- --> .Br2-
A58 .Br2- + HO2 --> Br2 + HO2- (see "Hydroperoxyl radical (HO2•) oxidizes dibromide radical anion (•Br2−) to bromine (Br2) in aqueous
solution: Implications for the formation of Br2 in the marine boundary layer" by Matthew, M., at http://onlinelibrary.wiley.com/doi/10.1029/2003GL018572/full )
A45 .Br + .Br2- --> Br2 + Br-
A59 .Br3- + HO2 --> .Br2- + Br- + O2 + H+
to list a few reactions.
Given the apparent importance of the hydroperoxyl radical, if one is adding air/O2 to the system, putting some H2 into the gas mix may be beneficial
in creating yet another path to the radical via:
H2 + .OH --> H2O + .H
H2 + .Br --> HBr + .H
.H + O2 --> .HO2 (if pH < 4.8)
However, there are undesirable reactions (some listed below), so this would require testing on the quantity of hydrogen to be added (if any):
.H + .OH --> H2O
.H + .Br ---> HBr = H+ + Br-
.H + .HO2 --> H2O2 (and possibly: H2O2 +.OH --> H2O + .HO2 , which recycles hydoperoxyl but consumes hydroxyl radicals)
Note, Table 3, per standard convention with radical reactions, cites forward reaction rates. There are instances where the reverse reaction are also
listed. The reactions I listed above are forward reactions where any cited corresponding reverse reaction are slower (at least, at the tested
temperature, pH,..)
Lastly, there has been expressed some concern over my extension of the pH range of the Fenton system by citrate, in the presence of bromide, being
essentially the introduction of an organic. However, in natural systems, there are key roles played by, for example, some organic like aromatic
ketones. See "Photoinduced oxidation by sea salt halides by aromatic ketones" at http://www.atmos-chem-phys.net/9/4229/2009/acp-9-4229-2009.p... .
[Edited on 9-3-2017 by AJKOER]
[Edited on 9-3-2017 by AJKOER]
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Panache
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| Quote: Originally posted by symboom | That is an interesting reaction though
Nitric acid formed it situ
Then protinating the potassium nitrate to nitric acid
Nitric acid reacting with potassium bromide forming
nitrogen dioxide gas and liquid bromine
Alot of these reactions become beneficial when the side products such as nitrogen dioxide can be used in another reaction.
Sodium bromide and (oxone) potassium monopersulfate forms bromine liquid
Just mix the two solids and add water
Oxone is avalible as non chlorine shock for pools
Works with potassium iodide too formine iodine solid
And sodium chloride forming chlorine gas
Quick and easy :-)
[Edited on 8-3-2017 by symboom]
[Edited on 8-3-2017 by symboom]
[Edited on 8-3-2017 by symboom] |
Are you serious about this? Is there more to it? Stoiciometry I Can work out but have you actually (or anyone else wiling to comment) performed the
conversion on all three halides? Any additional notes? Heating? Cooling, how much water is added? From what I recall of oxone oxidations they are
notoriously solvent (water, acetone) heavy. Does it splurge out the entire yield of halide immediately or slowly?
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woelen
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Nitrosyl bromide (and also nitrosyl
chloride) are much more stable in presence of water than most peaple think.
If you have NaBr and H2SO4 or HBr, and you have NaNO2, then try adding a pinch of solid NaNO2 to NaBr dissolved in 40% H2SO4, or solid NaNO2 to
40...50% HBr. In both cases you see that the liquid turns very dark chocolate-brown, nearly black and you get a very dense dark brown gas above the
liquid. The gas is more brown than NO2, it is more like chocolate.
I ampouled some of the very dark liquid, made by adding NaNO2 to appr. 40% HBr and have this ampoule around for half a year or so, and it still has a
very dark liquid in it, with a dark brown gas above it. The color differs a lot from the color of Br2, which is quite red, when dissolved in water.
The ONBr has no red hue.
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AJKOER
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| Quote: Originally posted by symboom | .......
......
Sodium bromide and (oxone) potassium monopersulfate forms bromine liquid
Just mix the two solids and add water
Oxone is avalible as non chlorine shock for pools
Works with potassium iodide too formine iodine solid
And sodium chloride forming chlorine gas
Quick and easy :-)
[Edited on 8-3-2017 by symboom]
..... |
A path liberating Cl2 from NaCl, Br2 from NaBr, I2 from NaI,...has been ascribed to the in situ formation of the sulfate radical anion. The process
actual follows from the formation again of the monoatomic halide radical (namely, .Cl, .Br, .I,...) in a reversible reaction. Per one source:
.SO4- + Cl- ⇆ SO4(2-) + .Cl
Link: http://www.sciencedirect.com/science/article/pii/13590197919... ). However, one may not need to spend $$ (see http://www.sigmaaldrich.com/catalog/product/sial/228036?lang... with 1,000 grams going for $65.80, but only 50.5% is KHSO5, so it is actually $130
for 1 kg) to buy oxone, as there are patents forming said active radical by passing air into waters (to be treated and in the current context,
containing bromide) in the presence of solar light, sulfite salt and transition metals (including cobalt in even trace amounts, manganese in larger
amounts,...), which may be successful here in liberating bromine from bromide.
Note, the radical chemistry (starting from O2, SO2, transition metals and sunlight), is even more complex than anything presented so far in this
thread, so unless someone asks, I will not attempt a presentation. Here is a link to a prior thread of mine with some good sample reference links for
those interested: http://www.sciencemadness.org/talk/viewthread.php?tid=71477.
Now, one still needs to spend $ on the preparation of a soluble sulfite from either the element sulfur (to create SO2,...) or purchase sulfite
directly being sold online for wine making and such, but I suggest, this could be a powerful method. Here is a quote on Calcium sulfite at $30 for
1,000 grams (link: http://www.molbase.com/en/search.html?search_keyword=calcium... ).
[Edit] For a more general discussion of the sulfate radical, see page 6 at https://www3.nd.edu/~ndrlrcdc/Compilations/Ino/Ino.pdf by Neta and Huie.
[Edited on 9-3-2017 by AJKOER]
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AvBaeyer
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Panache,
I make iodide "on demand" using oxone and potassium iodide. The iodine precipitates immediately as the oxone is added to the iodide solution. The
reaction is somewhat exothermic.
I also use oxone to generate bromine "on the fly" using sodium bromide. I never isolate the bromine but use a two-phase system to brominate double
bonds. I am currently working on a write up of an example this method. I will post this in the near future.
AvB
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clearly_not_atara
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Nitrosyl bromide sounds about as interesting as bromine if not moreso... it's a replacement for N2O3 or HNO2 right? I bet you can do some cool stuff
with it. I wonder if you can optimize the conditions to maximize NOBr production.
Oh, and bromine water will react with citric acid... it reacts with alpha-hydroxy acids in general, I remember studying this reaction as a way of
making a compound that we don't talk about from a certain perfume ingredient. The product is acetonedicarboxylic acid... which is also cooler than
bromine.
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CuReUS
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You can make Br2 by heating NaBr with just conc H2SO4 and scrubbing the gases produced to remove the
SO2.There is no need to add the nitrate since conc H2SO4 is itself a strong oxidiser. | Quote: Originally posted by JJay | | The first time I made bromine, I used hydrogen peroxide, hydrochloric acid, and sodium bromide. |
won't the peroxide oxidise HCl also ?
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JJay
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Yeah, but the chlorine produced would then oxidize the bromide.
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theAngryLittleBunny
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| Quote: Originally posted by theAngryLittleBunny |
I decided to not continue, since I planned to distill off the bromoacetyl bromide, but I don't wanna blow myself up because it might have formed some
weird explosives like acetyl nitrate, you never know. |
Oh, after thinking about that, this actually makes no sense, because nitrosyl bromide would most likely just brominate the acetic acid to acetyl
bromine:
CH3CO2H + BrNO -> CH3COBr + HNO2
And the HNO2 then decomposes, releasing NO and NO2 gas.
| Quote: Originally posted by clearly_not_atara | Nitrosyl bromide sounds about as interesting as bromine if not moreso... it's a replacement for N2O3 or HNO2 right? I bet you can do some cool stuff
with it. I wonder if you can optimize the conditions to maximize NOBr production.
|
I don't know what you could optimize, it's just the bromine equivalent of aqua regia:
HNO3 + 3HBr -> BrNO + Br2 + 2H2O
That's a really wild guess, but maybe you could use nitrite instead of nitrate, and the reaction would be like that:
HNO2 + HBr -> BrNO + H2O
Probably not, but who knows, I mean HNO2 can also be called nitrosyl hydroxide, and HBr being a really strong acid might be able to protonate it to
make nitrosyl bromide. But why would you use that stuff instead of HNO2? which is made from NaNO2, a perfectly stabile salt. Nitrosyl bromide is quite
unstable and boils at 14°C, and it's 73% by weight bromine, which means you have to use a lot if it compared to NaNO2.
| Quote: Originally posted by CuReUS |
You can make Br2 by heating NaBr with just conc H2SO4 and scrubbing the gases produced to remove the
SO2.There is no need to add the nitrate since conc H2SO4 is itself a strong oxidiser. |
I wouldn't call H2SO4 a "strong" oxidizer, it's pretty weak actually, but yes, I knew that it could oxidize bromide, but would have to use
concentrated H2SO4, I would have to distill the bromine, and the worst of all, it produces SO2 gas, and you can't imagine how much I hate SO2, it's
seriously the grossest gas ever >_< at least for me. Seriously, I would rather smell H2S then SO2, I even hate working with sulfite and
bisulfite.
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clearly_not_atara
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I can imagine dissolving sodium percarbonate in NaBr/H3PO4 as a particularly simple method. Sodium percarbonate is aka OxyClean.
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