| Pages:
1
2 |
walruslover69
Hazard to Others
Posts: 235
Registered: 21-12-2017
Member Is Offline
Mood: No Mood
|
|
Synthesis of pure Bromine with no byproducts
I plan on carrying out this synthesis this weekend but wanted to post it to make sure I haven't missed anything and hear your guy's thoughts.
250g of NaBr is acidified with 150ml of concentrated H2SO4 (I plan on diluting the acid to prevent oxidation reactions).
HBr will be distilled over.
The HBr will then be reacted with 65g of Fe2O3 to produce FeBr3.
The flask will then be boiled to produce Bromine by the reaction
2FeBr3-->2FeBr2+Br2.
Since there will be no bromide ions in the receiving flask the solubility of the bromine will be only marginal (~.33g/L). The bromine will be
separated from the water in a separatory funnel and dried in a yet unknown way.
The final yield should be 32 grams of bromine.
Sodium bromide and Iron oxide are extremely cheap at ~$10 a pound making this a very cost effective way of producing bromine.
All of the steps should produce near quantitative yield and I do not see an opportunity for other byproducts or side reactions that would contaminate
the bromine.
I have seen sulfuric acid being used by other people to dry their bromine. are bromine and sulfuric acid completely insoluble in each other? I
couldn't find any data but that seems a little fishy to me.
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
Based on this work, https://agupubs.onlinelibrary.wiley.com/doi/full/10.1029/200... , in general to liberate bromine from bromide, we want hydroxyl radicals (OH•)
and therefrom perhydroxyl radical (HO2• ) in acidic conditions, producing the following reaction system interacting with bromide and H2O2:
Br- + OH• --> BrOH•-
BrOH•- --> Br- + OH• (the reversible reaction of above)
BrOH•- + H+ --> Br• + H2O
OH• + H2O2 --> H2O + HO2•
Br• + Br• --> Br2
Br• + Br- --> Br2•-
Br2•- + Br2•- --> 2 Br- + Br2
Br2•- + HO2• --> Br2 + HO2-
Br2•- + HO2• --> 2 Br- + O2 + H+ (competing reaction to above)
H+ + HO2- = H2O2
Now, since you have FeBr3, add H2O2 to form the perhydroxyl radical:
Fe(3+) + H2O2 --> Fe(2+) + H+ + HO2•
The hydroxyl radical then flows from the Fenton reaction with H2O2 acting on ferrous, which allows us to proceed to the first equation above:
Fe(2+) + H2O2 + H+ --> Fe(3+) + OH• + H2O
I suspect to the extent that light promotes a photo-Fenton reaction, the results may improve.
----------------------------------------
Now, less effective but cheaper, one could substitute O2 (from rapid boiling) for H2O2 as:
Fe(2+) + O2 = Fe(3+) + O2•- (a reversible reaction)
H+ + O2•- = HO2• (pH < 4.88, reaction moves to the right)
HO2• + HO2• = H2O2 + O2
and the presence of H2O2 now resembles the system above.
-------------------------------
Note, in place of the transition metal iron (Fe), try also copper (Cu) or, I suspect, due to the influence of favorable redox couple equilibrium
reaction:
Fe(3+) + Cu(+) = Fe(2+) + Cu(2+)
both FeBr3 and CuBr2 together with acid and H2O2 may actually perform better liberating bromine from the bromide.
[Edited on 10-7-2018 by AJKOER]
|
|
|
CobaltChloride
Hazard to Others
Posts: 239
Registered: 3-3-2018
Location: Romania
Member Is Offline
|
|
Why make your life harder and lower yields by using such a weak oxidant? Why wouldn't you want to use something like hydrogen peroxide or pool
chlorinator to oxidize bromide to bromine?
|
|
|
walruslover69
Hazard to Others
Posts: 235
Registered: 21-12-2017
Member Is Offline
Mood: No Mood
|
|
There is no oxidant involved in the reaction. FeBr3 thermally decomposes to FeBr2 +Br2.
The problem with using hydrogen peroxide is that bromine acts as a catalyst to decompose it.
In theory pool chlorinator works to oxidize bromide, I believe Extractions&ire has a great video on it. But you have the problem of over
chlorinating and creating bromo mono chloride which I don't know of a way to effectively separate.
the FeBr2 could also be reoxidized in the presence of HBr to regenerate the FeBr3, but I havn't worked out what an effective oxidant yet since I
haven't been able to find any good info on it. If I can regenerate the FeBr3 then It would only be using FeBr2/FeBr3 as a catalyst in the direct
conversion
2HBr-->Br2.
|
|
|
MrHomeScientist
International Hazard
Posts: 1806
Registered: 24-10-2010
Location: Flerovium
Member Is Offline
Mood: No Mood
|
|
Seems to me that your disproportionation of FeBr3 would waste potential bromine, since 2/3 of the atoms are still bound up with iron!
In my video on making bromine ( https://www.youtube.com/watch?v=NKjyM2AkZSY ), I followed woelen's procedure to make it electrolytically. This worked out great, and I don't
believe there were any side products. Certainly not any interhalogens, since there weren't any others around. Shameless plug of my video, but I'm very
proud of that one. 
I dried mine by shaking with 98% sulfuric acid. I don't know what the solubility would be, if any. I suppose you could keep shaking your product with
fresh acid, and if the weight of the bromine stops changing after a few additions then solubility would be negligible.
Edit: quit changing the tags for subscripts! One day it's [] and the next it's <> 
[Edited on 7-10-2018 by MrHomeScientist]
|
|
|
clearly_not_atara
International Hazard
Posts: 2818
Registered: 3-11-2013
Member Is Offline
Mood: Big
|
|
FeBr3 cannot be dehydrated by heating, last I checked. I don't think this synthesis will work the way you expect it to. More likely heating FeBr3*NH2O
results in the release of HBr and leaves some iron oxybromide hydrate in the flask.
The same reaction with CuBr2 does work, though, IIRC.
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
Quote: Originally posted by CobaltChloride  | | Why make your life harder and lower yields by using such a weak oxidant? Why wouldn't you want to use something like hydrogen peroxide or pool
chlorinator to oxidize bromide to bromine? |
Actually, this is a very good question.
One actually needs something like the HO2• to remove the pesty and persisent Br2•- radical!
However, now I think about it, one might be able to substitute a more powerful organic radical (RO2•) for HO2•, but only after taking into
consideration safety concerns. While concentrated H2O2 is dangerous, per Wikipedia (https://en.wikipedia.org/wiki/Diacetyl_peroxide ) generally on organic peroxides to quote:
"Organic peroxides are typically explosive since they contain both the oxidizer, the O-O bond, and reducing agents, the C-C and C-H bonds."
In the particular case of diacetyl peroxide with the formula (CH3CO2)2, to quote the same Wikipedia source:
"Since the pure material poses an explosion hazard, it is often used as a solution, e.g., in dimethyl phthalate as a solvent.[1]"
which may(?) be a suitable stronger substitute here for H2O2. Test for safety first given possible impurities in the reaction mix.
[Edited on 11-7-2018 by AJKOER]
|
|
|
CobaltChloride
Hazard to Others
Posts: 239
Registered: 3-3-2018
Location: Romania
Member Is Offline
|
|
| Quote: Originally posted by walruslover69 | There is no oxidant involved in the reaction. FeBr3 thermally decomposes to FeBr2 +Br2.
The problem with using hydrogen peroxide is that bromine acts as a catalyst to decompose it.
In theory pool chlorinator works to oxidize bromide, I believe Extractions&ire has a great video on it. But you have the problem of over
chlorinating and creating bromo mono chloride which I don't know of a way to effectively separate.
|
The weak oxizing agent I was reffering to here was Fe (III). In that decomposition reaction Fe (III) oxidizes Br- to Br2, itself being reduced to Fe
(II). You're right that using any hypochlorite based oxizing agent here could introduce a minor impurity of bromine chlorides, but what would the
problem be if bromine catalyzes the decomposition of H2O2? You could add it slowly to the reaction mixture like here https://www.youtube.com/watch?v=HuHceKvSHk0&t
|
|
|
walruslover69
Hazard to Others
Posts: 235
Registered: 21-12-2017
Member Is Offline
Mood: No Mood
|
|
| Quote: Originally posted by clearly_not_atara | FeBr3 cannot be dehydrated by heating, last I checked. I don't think this synthesis will work the way you expect it to. More likely heating FeBr3*NH2O
results in the release of HBr and leaves some iron oxybromide hydrate in the flask.
The same reaction with CuBr2 does work, though, IIRC. |
I have found at least one source that references "Solutions of
iron(III) bromide decompose to iron(II) bromide and bromine on boiling."
https://sci-hub.tw/https://doi.org/10.1002/0471238961.091815...
What same reactions were you referring to with CuBr2? From what I have found it seems to be thermally stable.
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
On the issue of the reaction between Br2 and H2O2, more likely I suspect the issue is between the reaction:
Br2 + H2O = HBrO + HBr
forming HOBr for which there is a fast reaction between the latter and H2O2:
HOBr + H2O2 --> HBr + H2O + O2 (or even some Singlet oxygen)
However, if the bromine water is over 59 C, the Br2 boils off, reducing the possible HOBr formation and H2O2 consumption.
-----------------------------------------------------------------------------
On the question of copper in place or in conjunction with iron, the action of cuprous with H2O2 is called a Fenton-like reaction, which in acidic
conditions proceeds similarly to that of ferrous creating the required hydroxyl radical:
Cu(+) + H2O2 + H+ --> Cu(2+) + OH• + H2O
Note, the cuprous must be soluble, as in a complex with bromide, for example. Also, in neutral to basic conditions, no hydroxyl radical, but Cu(3+),
which is not likely congruous with the path outlined above (see https://www.researchgate.net/publication/235921389_Fenton-li... ).
-----------------------------------------
In a reaction system without H2O2, I claim that oxygen is required. If one disagrees add a layer of oil over the solution and carefully heat, leaving
only dissolved oxygen in the solution, expect a minor product yield.
[Edited on 11-7-2018 by AJKOER]
|
|
|
clearly_not_atara
International Hazard
Posts: 2818
Registered: 3-11-2013
Member Is Offline
Mood: Big
|
|
Interesting. I'm not sure how to dry the produced bromine but it sounds feasible. Wrt CuBr2 see e.g.: http://www.journal.csj.jp/doi/10.1246/bcsj.43.3468
|
|
|
walruslover69
Hazard to Others
Posts: 235
Registered: 21-12-2017
Member Is Offline
Mood: No Mood
|
|
I think it could probably be dried with something like magnesium sulfate or copper sulfate.
CobaltChloride- the H2O2 method would work in a similar manor, I overestimated the amount of H2O2 that would be decomposed. I just don't have an
equalizing addition funnel or concentrated H2O2 at the moment.
|
|
|
CobaltChloride
Hazard to Others
Posts: 239
Registered: 3-3-2018
Location: Romania
Member Is Offline
|
|
Ok then. Do you have access to sodium/potassium/ammonium persulfate (peroxydisulfate) from electronics stores? Or do you have acces to
sodium/potassium peroxymonosulfate (Oxone) from pool disinfectant? Those are the cleanest and most powerful oxidants available to the amateur
chemist. Oxone wouldn't even require any acid addition.
[Edited on 11-7-2018 by CobaltChloride]
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
| Quote: Originally posted by CobaltChloride | Ok then. Do you have access to sodium/potassium/ammonium persulfate (peroxydisulfate) from electronics stores? Or do you have access to
sodium/potassium peroxymonosulfate Oxone) from pool disinfectant? Those are the cleanest and most powerful oxidants available to the amateur
chemist. Oxone wouldn't even require any acid addition.
[Edited on 11-7-2018 by CobaltChloride] |
One would try to activate peroxymonosulfate (or PMS) with mineral rich tap water, or even try carbonated mineral water.
Apparently, however, one of the best way to activate peroxydisulfate (PDS) and peroxymonosulfate (PMS) is siderite, (Fe,Mg,Ca,Mn,Zn,Co)CO3, where Fe,
Mn and cobalt are the key metal ions (apparently, a trace of the toxic cobalt ions in the presence of HSO5- is capable of producing the powerful
sulfate radical anion, see "COBALT/PEROXYMONOSULFATE AND RELATED OXIDIZING REAGENTS FOR WATER TREATMENT" a thesis by Georgios P. Anipsitakis, https://etd.ohiolink.edu/rws_etd/document/get/ucin1130533674... ). The ensuing advanced oxidation process can then produce both the usual sulfate
radicals (see https://www.sciencedirect.com/science/article/pii/S004313540... ) and the even more powerful hydroxyl radical (see "Activation of Persulfates
Using Siderite as a Source of Ferrous Ions: Sulfate Radical Production, Stoichiometric Efficiency, and Implications", at https://pubs.acs.org/doi/abs/10.1021/acssuschemeng.7b03948 ).
The siderite functions as a source of Fe2+ (think fenton with KHSO5 replacing H2O2) with the sulfate radicals as the major active radical. In place of
siderite, also consider FeS2 (Fool's Gold, see for example, https://pubs.acs.org/doi/abs/10.1021/ie100740d to which I would also pump in air/O2 to expand possible reactive oxygen species creation. Further,
I would suspect adding the Bravoite variation of Fool's Gold, which contains toxic nickel and cobalt, would be favorable).
Example of activation reactions with cobalt ions:
Co2+ + HSO5- --> Co3+ + SO4•- + OH- (2.1 from thesis)
Co3+ + HSO5- --> Co2+ + SO5•- + H+ (1.11 from thesis)
Note, the action of the sulfate radical anion, for example, on bromide (or other halogens, in reverse order of their reactivity, so the action on
iodide is faster than with bromide...) proceeds as follows:
SO4•- + Br- --> SO4(2-) + •Br
Note, the sulfate radical anion is weaker than the hydroxyl radical. Hypochlorous acid activated by say cuprous creates the more powerful hydroxyl
radical. In some cases, PMS is not recommended for swimming pools as it does not completely kill off problem microbes.
I would think that the following reactions are likely (but no source):
Br2•- + SO4•- --> Br2 + SO4(2-)
Br2•- + SO5•- + H+ --> Br2 + HSO5-
So, PMS may work for Br2 extraction but is certainly more costly with less access than dilute H2O2.
[Edited on 11-7-2018 by AJKOER]
|
|
|
walruslover69
Hazard to Others
Posts: 235
Registered: 21-12-2017
Member Is Offline
Mood: No Mood
|
|
I can buy ~50% potassium PMS at my local walmart for ~$10 a pound so it is cheap and readily available. I am curious how well sulphuric acid would act
as a oxidizer by itself if heated to 100-200C. There might be some problem with sulfur dioxide converting some of the bromine in the collection vessel
to HBr but that could easy be separated with distillation or washing with water.
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
My examination of the online products available at Walmart indicated one item, Leisure Time Spa Care for $27.99 for 2.2 pounds of 32% PMS. Found a
picture on another website, see https://spaandpoolstore.com/leisure-time-renew-2-2-lbs-granu... . Price does not include shipping. Note, the inert ingredient is a pH buffer,
which may impair effectiveness of the redox reaction with various transition metals (see http://superfund.berkeley.edu/pdf/425.pdf ).
--------------------------------------
For those wanting to test a new approach, which is based on related work of mine, try adding an acidified copper acetate/FeBr3 solution to dry sodium
percarbonate. In essence a peroxymonocarbonate transition metals activated path to carbonate radicals to release bromine. This method also does form
the perhydroxyl radical also via:
H2O2 + CO3•‒ → HO2• + HCO3-
to address the Br2•- radical.
[Edited on 11-7-2018 by AJKOER]
|
|
|
JJay
International Hazard
Posts: 3440
Registered: 15-10-2015
Member Is Offline
|
|
If sulfuric acid is significantly soluble in bromine (and it's an issue), why not simply distill the bromine after it is dried?
|
|
|
clearly_not_atara
International Hazard
Posts: 2818
Registered: 3-11-2013
Member Is Offline
Mood: Big
|
|
So I guess my real point here is: if you're making bromine in an aqueous phase, your procedure does not have any real advantage over the use of H2O2
to oxidize bromide salts, which is the standard way to make bromine. In acidic solution, the only product is bromine. Doesn't matter whether your
oxidizing agent is Fe3+, persulfates, Cu2+, or whatever, an aqueous process is an aqueous process and the product is bromine water.
Perhaps you could expand on why you're interested in "pure" bromine, so that a better procedure can be determined?
| Quote: | | I think it could probably be dried with something like magnesium sulfate or copper sulfate. |
This is good if you're dealing with a gas stream which is 90% bromine, 10% water. But if it's 90% water, 10% bromine, you might want to rethink that.
The reference gave no mention of the rate of bromine production in refluxing FeBr3/H2O.
[Edited on 11-7-2018 by clearly_not_atara]
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
| Quote: | | Quote: Originally posted by clearly_not_atara |
................
In acidic solution, the only product is bromine. Doesn't matter whether your oxidizing agent is Fe3+, persulfates, Cu2+, or whatever, an aqueous
process is an aqueous process and the product is bromine water.
|
........
[Edited on 11-7-2018 by clearly_not_atara] |
In my opinion, given the rate of this reaction (source: see attached file):
Br• + Br− --> Br2•- ( k= 1.2×10^10 M-1 s-1 )
where the Br• is created from a hydroxyl radical interacting with bromide ion in acidic conditions:
Br- + OH• --> BrOH•-
BrOH•- + H+ --> Br• + H2O
Br2•- + Br2•- --> 2 Br- + Br2
Br2 + Br- --> Br3- ( k = 9.6×10^8 M-1 s-1 )
and per the last reaction above, what is referred to above as 'bromine water' is, in effect in my opinion, perhaps more correctly a mixture of bromine
complexes and bromine water.
In my first reference in this thread, '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 Brendan M. Matthew, et al, at https://agupubs.onlinelibrary.wiley.com/doi/full/10.1029/200... , to quote:
"One proposed pathway for the formation of Br2 is the oxidation of Br− by hydroxyl radical (•OH) [Mozurkewich, 1995]. As shown in Figure 1,
•Br2− is a key intermediate formed from this reaction. The subsequent fate of •Br2− (to produce either Br2 or Br−) likely plays a
significant role in reactive halogen release. "
Also:
"Although Mozurkewich [1995] has examined Br2 release from sea‐salt particles using reaction R1, no other published models allow HO2• to oxidize
•Br2− to Br2 with a rate constant on the order of 10^9 M−1 s−1 (as reported in the first three papers in Table 1). Thus these current models
might be underestimating the release of Br2 and other reactive halogens."
My take on the above is that there is an important difference between free bromine and complexed bromine with respect to implications for chemical
reactivity.
Attachment: molecules-22-01684-s001 (3).pdf (287kB)
This file has been downloaded 380 times
[Edited on 11-7-2018 by AJKOER]
|
|
|
walruslover69
Hazard to Others
Posts: 235
Registered: 21-12-2017
Member Is Offline
Mood: No Mood
|
|
| Quote: Originally posted by clearly_not_atara |
Perhaps you could expand on why you're interested in "pure" bromine, so that a better procedure can be determined?
| Quote: | | I think it could probably be dried with something like magnesium sulfate or copper sulfate. |
This is good if you're dealing with a gas stream which is 90% bromine, 10% water. But if it's 90% water, 10% bromine, you might want to rethink that.
The reference gave no mention of the rate of bromine production in refluxing FeBr3/H2O.
[Edited on 11-7-2018 by clearly_not_atara] |
I plan on just separating the bromine/water by seperatory funnel and then drying the bromine with magnesium/copper sulfate.
FeBr3 decomposes to bromine at 200C so I imagine that the bromine coming off from boiling the solution is comparable to steam distillation. If it is
very slow and only small amounts come over as the water boils then once most of the water has boiled off and the FeBr3 remains it will quickly heat up
and decompose fairly quickly.
I have a few plans for Bromine. The first is that I would just like a small sample to have in my collection. Since I am synthesizing bromine I thought
why not try and explore some other technique that hasn't been attempted in the amateur community. Decomposition of a bromide salt to bromine just
struck me as interesting.
I would also like to use it to make aluminum bromide and some bromoform that I could use for other projects in the future.
|
|
|
woelen
Super Administrator
Posts: 8065
Registered: 20-8-2005
Location: Netherlands
Member Is Offline
Mood: interested
|
|
Making bromine is covered on sciencemadness many many times before. the method, propesed here with Fe2O3 does not seem like a viable route to me. If
you are unlucky, then you even cannot dissolve your Fe2O3 in the aqueous HBr you want to make. Fe2O3 can be remarkably inert. Try adding some to
dilute H2SO4. Does it dissolve to give a ferric sulfate solution. If not, then it also will not dissolve in aqueous HBr.
Search the forums for threads on making Br2. One option is to use electrolysis and the easy to get NaHSO4 (H2SO4 of course also works). One can also
use all kinds of oxidizers. One oxidizer works better than the other, but of course, availability may also be a determining/limiting factor.
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
May I suggest some interesting experiments starting with Fool's Gold. Per Wikipedia (https://en.wikipedia.org/wiki/Iron(II)_sulfate) :
"Ferrous sulfate is also prepared commercially by oxidation of pyrite:
2 FeS2 + 7 O2 + 2 H2O → 2 FeSO4 + 2 H2SO4 "
I would add some sea salt and use a microwave to jump start the reaction.
Next:
FeSO4 (aq) + 2 NaBr = Na2SO4 (aq) + FeBr2 (aq)
and try to freeze out the Na2SO4 hydrate leaving aqueous ferrous bromide.
This is followed by an electrochemical reaction with aqueous ferrous, pumped in oxygen and much added H+ (from say adding NaHSO4 followed again by
freezing out the Na2SO4) to hopefully limit the formation of a basic bromide salt:
4 Fe(2+) + O2 + 2 H+ → 4 Fe(3+) + 2 OH- (see, for example, https://www.chegg.com/homework-help/questions-and-answers/co... )
Unfortunately, the product may still include some Fe(OH)Br2 (or a lot which would be very problematic for successful production of Bromine).
Then, the thermal decomposition of dry FeBr3, which per a reference (see https://chemiday.com/en/reaction/3-1-0-5163 ) to quote:
"The thermal decomposition of iron(III) bromide to produce iron(II) bromide and bromine. This reaction takes place at a temperature of over 139°C."
2FeBr3 + Heat → 2FeBr2 + Br2 (at 139°C)
The FeBr2 could be theoretically recycled by the electrochemical reaction with oxygen and acid discussed above.
[Edited on 13-7-2018 by AJKOER]
|
|
|
nezza
Hazard to Others
Posts: 324
Registered: 17-4-2011
Location: UK
Member Is Offline
Mood: phosphorescent
|
|
If you are going to distil the bromine off anyway why not use sulphuric acid and permanganate. I have used that in the past as a method of preparing
bromine.
If you're not part of the solution, you're part of the precipitate.
|
|
|
clearly_not_atara
International Hazard
Posts: 2818
Registered: 3-11-2013
Member Is Offline
Mood: Big
|
|
| Quote: | | If it is very slow and only small amounts come over as the water boils then once most of the water has boiled off and the FeBr3 remains it will
quickly heat up and decompose fairly quickly. |
Again, ferric halides can't be dehydrated by heating. Dehydrating FeCl3 hydrates results in HCl release and the formation of iron oxychlorides. With
FeBr3 you'll get something similar.
I don't know what the rate of decomposition of FeBr3 is, but I expect a messy process. If bromine is retained in the boiling flask it will reoxidize
FeBr2 as the solution cools; if it is distilled out you might also get a lot of HBr. The presence of bromide ions increases the aqueous solubility of
Br2 by the reaction:
Br2 + Br- >> Br3-
I don't know the best way to separate bromine from bromine water, but I suspect that adding enough sodium acetate (no chlorides!) to lower the
freezing point below -7 C and then cooling to freeze the bromine will do better than a separatory funnel. I guess yield is not a major concern to you,
though.
Also if you want bromine in a collection I hope you can ampoule it, otherwise it's a hazard.
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
| Quote: Originally posted by clearly_not_atara | | Quote: | | If it is very slow and only small amounts come over as the water boils then once most of the water has boiled off and the FeBr3 remains it will
quickly heat up and decompose fairly quickly. |
Again, ferric halides can't be dehydrated by heating. Dehydrating FeCl3 hydrates results in HCl release and the formation of iron oxychlorides. With
FeBr3 you'll get something similar.
I don't know what the rate of decomposition of FeBr3 is, but I expect a messy process. If bromine is retained in the boiling flask it will reoxidize
FeBr2 as the solution cools; if it is distilled out you might also get a lot of HBr.
.............
|
Some comments from Atomistry.com (based on possibly dated extracts from chemical journals) on FeBr3 (link: http://iron.atomistry.com/ferric_bromide.html ):
" Ferric bromide, FeBr3, results when iron is heated in excess of bromine vapour, when ferrous bromide is heated from 170° to 200° C. with twice its
weight of bromine, and when sulphur monobromide, S2Br2, is passed over ferric oxide at 450° to 650° C. It is a dark red, crystalline but
deliquescent substance, yielding a red solution in water. At boiling-point the solution dissociates to ferrous bromide and free bromine. On heating
away from air ferric bromide partly sublimes and partly dissociates. When dry the salt is reduced by nitric oxide with the formation of ferrous and
nitrosyl bromides.
The aqueous solution upon concentration in the ordinary way decomposes with the precipitation of insoluble basic bromides.
The hexahydrate, FeBr3.6H2O, separates as dark green needles when the dark brown solution obtained by the action of bromine under water is slightly
evaporated and concentrated in a desiccator over sulphuric acid. It is soluble in alcohol and ether, and melts at 27° C. without decomposition."
So, per the above comments, apparently boiling aqueous FeBr3 may result in the creation of insoluble basic bromides, likely per a path I discussed
above.
------------------------------
Per Wikipedia on FeBr2 (link: https://en.wikipedia.org/wiki/Iron(II)_bromide) to quote:
"FeBr2 is synthesized using a methanol solution of concentrated hydrobromic acid and iron powder. It adds the methanol solvate [Fe(MeOH)6]Br2 together
with hydrogen gas. Heating the methanol complex in a vacuum gives pure FeBr2.[2] Iron(II) bromide cannot be formed by the reaction of iron and
bromine, because that reaction would produce ferric bromide.[citation needed]
....
FeBr2 reacts with bromide and bromine to form the intensely colored, mixed-valence species [FeBr3Br9]−."
From above, my suggested pure aqueous path for FeBr2 and FeBr3 appears more suspect (however, adding alcohol to the water mix may prove to be of
benefit).
[Edited on 13-7-2018 by AJKOER]
|
|
|
| Pages:
1
2 |
![[*]](images/xpblue/default_icon.gif){kind=icon}
