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Molecular Manipulations
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Quote: Originally posted by j_sum1  | Good point.
Amazing how easy it is to overlook the obvious.
(But then if you are doing that, you could always do electrolysis of NaBr!!) (Overlooked by both of us  And everyone who read what I wrote without catching the glitch. ) |
Literally overlooked by you. MrHomescientist linked a video by him on it earlier in this thread. I personally dislike most
electrolysis isolations, but that's just me. They take a long time, are usually open to air and thus can get contaminated, plus my MMO gets degraded
by bromine, and I hate using carbon rods.
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deltaH
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One might be able to form bromine by melting and then distilling a mixture of ferric nitrate nonahydrate (melting point 47.2 °C) with a bromide salt.
The hypothetical reaction for the sodium salt, for example, that I'm thinking of would correspond to:
10NaBr(s) + 4Fe(NO3)3.9H2O(l) + heat => 4FeOOH(s) + 32H2O(l) + 10NaNO3(aq) + 5Br2(g) + 2NO(g)
There could be some contamination with nitrosyl bromide as there is an equilibrium between bromine and nitric oxide.
Again, heed the warning of the EXTREME toxicity and corrosiveness of some of the
hypothetical products! Chemical pneumonitis is easily induced from these gases even in very small amounts.
[Edited on 23-2-2015 by deltaH]
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blogfast25
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I doubt if that would work. The low temperature works against it, I think.
When the ferric nitrate nonahydrate 'melts' it essentially dissolves into its own crystal water. Nitrate ions alone can't oxidise bromide (better
check that against a reduction potentials table!).
Edit:
NO<sub>3</sub><sup>-</sup> + 4 H<sup>+</sup> + 3 e ===> NO + 2 H2O
Ered = + 0.95 V
Br<sup>-</sup> ===> 1/2 Br<sub>2</sub> + e
Eox = - 1.07 V
So E = Ered + Eox = - 0.12 V < 0
Which suggests this won't work. Oxidation with nitrates at higher temperatures might work by removing volatile Br vapours.
[Edited on 23-2-2015 by blogfast25]
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deltaH
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| Quote: Originally posted by blogfast25 |
I doubt if that would work. The low temperature works against it, I think.
When the ferric nitrate nonahydrate 'melts' it essentially dissolves into its own crystal water. Nitrate ions alone can't oxidise bromide (better
check that against a reduction potentials table!). |
Surprisingly, I had done a reaction where a chloride salt was heated with molten Fe(NO3)3.9H2O. This produced a very noxious gas with a strong
chlorine odour, so I suspected formation of at least some chlorine or nitrosyl chloride (which can decompose to chlorine anyhow in an equilibrium with
NO) and or NOx.
If that was the case, then the bromide version should proceed more readily.
But again, the product might be contaminated by the other gases [or not].
| Quote: | Edit:
NO3- + 4 H+ + 3 e ===> NO + 2 H2O
Ered = + 0.95 V
Br- ===> 1/2 Br2 + e
Eox = - 1.07 V
So E = Ered + Eox = - 0.12 V < 0
Which suggests this won't work. Oxidation with nitrates at higher temperatures might work by removing volatile Br vapours.
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Those reduction potentials are very close blogfast.
I don't know what would happen if one used nitric acid and a bromide salt as is, the reason for stating the iron nitrate version is because of that
experimental observation I had (chlorine case).
With aqua regia, there is also a complex equilibrium between nitric acid, water, hydrochloric acid, nitrosyl chloride, NOx and chlorine that is
suggestive (chlorine case).
Finally, this route of using ferric nitrate sticks to the spirit of the thread, i.e. avoiding the use of liquid acids. However, nitric acid can be
considered to be generated in situ by the easy decomposition of ferric nitrate.
[Edited on 23-2-2015 by deltaH]
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blogfast25
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E = 0 V does mean K approx. 1. Maybe a bit of sodium bisulphate would help too.
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deltaH
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Why?
Thinking of the fused sodium bromide + sodium nitrate + sodium bisulfate variant?
Could work, at high temperatures, nitrosyl bromide is probably not favoured thermodynamically, so generated NOx should pass straight through the
condenser in part (aside from what is soluble in bromine).
However, generating HBr instead is a competing parallel reaction that could be a big problem with the addition of "dry acid".
Would you be in a position to safely try out the ferric nitrate version I have described by any chance?
[Edited on 23-2-2015 by deltaH]
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Molecular Manipulations
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How about a dry method? Again theoretical:
2 NaBr + Ca(OCl)2 + Mg --> Br2 + 2 NaCl + CaO + MgO.
I haven't yet checked out the Gibbs Free Energy for this, everything is forming a very stable ionic compound, except bromine. Could this work?
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blogfast25
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No, I was still thinking of your ferric nitrate nonahydrate idea.
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blogfast25
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| Quote: Originally posted by Molecular Manipulations | How about a dry method? Again theoretical:
2 NaBr + Ca(OCl)2 + Mg --> Br2 + 2 NaCl + CaO + MgO.
I haven't yet checked out the Gibbs Free Energy for this, everything is forming a very stable ionic compound, except bromine. Could this work?
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ΔG almost certainly quite negative. But possibly a can of worms in terms of byproducts, I think, not sure...
Interesting thought though.  Might want to add a low melting flux to promote
contact.
Or why not leave out the Mg altogether? Ca(ClO)2 is a strong oxidiser as such.
NaBr + Ca(ClO)2 === > NaCl + CaO + 1/2 Br2
('I may be a dreamer but I'm not the only one...'  )
[Edited on 24-2-2015 by blogfast25]
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Molecular Manipulations
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Because if you look closely you'll notice that can't be balanced, there's two oxygens and two chlorines in calcium hypochlorite, no matter how you try
it, either oxygen or chlorine must evolve. Chlorine is in the +1 oxidation state (I know, oxidation states don't really exist), hence the necessary
reducing agent.
[Edited on 24-2-2015 by Molecular Manipulations]
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blogfast25
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| Quote: Originally posted by Molecular Manipulations | Because if you look closely you'll notice that can't be balanced, there's two oxygens and two chlorines in calcium hypochlorite, no matter how you try
it, either oxygen or chlorine must evolve. Chlorine is in the +1 oxidation state (I know, oxidation states don't really exist), hence the necessary
reducing agent.
[Edited on 24-2-2015 by Molecular Manipulations] |
Ooopsie. That Brandy and Coke is taking its toll.
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deltaH
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I don't think it's a good idea to add magnesium to an oxidant if you desire to make bromine.
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blogfast25
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Flash powder alert!
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Molecular Manipulations
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I was thinking the same because of magnesium oxide contamination. Is this why, or some other reason? Glass wool could be use at the neck of the flask
to stop some of it, but it would probably need to be distilled twice.
Of course it could just not work, or go way too fast.
Or the magnesium oxidation could be fast, but it would take a while for chlorine to oxidize bromide, I think the sodium bromide would need to be
fused, and even still it will take a while.
[Edited on 24-2-2015 by Molecular Manipulations]
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blogfast25
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How about replacing that 'nasty' Mg with something a bit tamer like Zn?
Mg with chlorates is explosive, not sure about with molten hypochlorites.
I don't get your point about MgO: magnesium powder is pretty stable and forgiving.
[Edited on 24-2-2015 by blogfast25]
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Metacelsus
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Instead of adding magnesium, what about adding an acid?
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blogfast25
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Equation?
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Molecular Manipulations
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It's just a matter of avoiding contamination, magnesia is very stable, I just don't want it in my bromine.
Zinc seems like a good idea, it would melt easily and not react to fast. Also whether zinc or magnesium it won't be powdered, probably chips, so
flash powder isn't going to happen.
As for Cheese's acid, first, this entire point is not to use acid, if acid was to be considered I'd just use the regular acid +
bromide + oxidizer.
Besides, calcium oxide will react with acids...
[EDIT] Shouldn't this be is R and A Acquisition I've already waisted so many posts here that I'm now most active in 'beginnings'...
[shudder]
[Edited on 24-2-2015 by Molecular Manipulations]
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We are all here on earth to help others; what on earth the others are here for I don't know. -W. H. Auden
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blogfast25
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But how would it get in there? You distil off the bromine, so the worst that can happen is mechanical entrainment and that can happen with the other
reaction products too.
Sodium bisulphate is a salt.
[Edited on 24-2-2015 by blogfast25]
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deltaH
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What's wrong with ferric nitrate and bromide salt route?
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Molecular Manipulations
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Yeah, I was referring to mechanical 'entrainment'. I had the image of flash powder in the back of my mind when I wrote that, it makes a lot of fine
magnesium oxide dust (among other things). I guess if the reaction is slow this might not be a problem.
You're quite right, sodium bisulfate is indeed a salt, and an acidic one at that. Is there a point to that, or are just saying words now?
Tin tetrachloride is a covalent liquid...
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blogfast25
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The emoticon should have been a bit of a giveaway (it was a joke, i.o.w.)
Post with 'you can't use XYZ' never have much sway with me. NaHSO4 is a case in point: so OTC no one can reasonably object to recommending it in an
amateur synthesis.
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blogfast25
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We've already discussed that, no?
| Quote: Originally posted by deltaH | | Quote: Originally posted by blogfast25 |
I doubt if that would work. The low temperature works against it, I think.
When the ferric nitrate nonahydrate 'melts' it essentially dissolves into its own crystal water. Nitrate ions alone can't oxidise bromide (better
check that against a reduction potentials table!). |
Surprisingly, I had done a reaction where a chloride salt was heated with molten Fe(NO3)3.9H2O. This produced a very noxious gas with a strong
chlorine odour, so I suspected formation of at least some chlorine or nitrosyl chloride (which can decompose to chlorine anyhow in an equilibrium with
NO) and or NOx.
If that was the case, then the bromide version should proceed more readily.
But again, the product might be contaminated by the other gases [or not].
| Quote: | Edit:
NO3- + 4 H+ + 3 e ===> NO + 2 H2O
Ered = + 0.95 V
Br- ===> 1/2 Br2 + e
Eox = - 1.07 V
So E = Ered + Eox = - 0.12 V < 0
Which suggests this won't work. Oxidation with nitrates at higher temperatures might work by removing volatile Br vapours.
|
Those reduction potentials are very close blogfast.
I don't know what would happen if one used nitric acid and a bromide salt as is, the reason for stating the iron nitrate version is because of that
experimental observation I had (chlorine case).
With aqua regia, there is also a complex equilibrium between nitric acid, water, hydrochloric acid, nitrosyl chloride, NOx and chlorine that is
suggestive (chlorine case).
Finally, this route of using ferric nitrate sticks to the spirit of the thread, i.e. avoiding the use of liquid acids. However, nitric acid can be
considered to be generated in situ by the easy decomposition of ferric nitrate.
[Edited on 23-2-2015 by deltaH] |
Without some source of H<sub>3</sub>O<sup>+</sup> that won't work too well, if you rely on the 'classic' oxidising power of
nitrates.
Also, why ferric nitrate hydrate? Why not Al nitrate hydrate?
[Edited on 24-2-2015 by blogfast25]
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deltaH
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| Quote: |
Without some source of H<sub>3</sub>O<sup>+</sup> that won't work too well, if you rely on the 'classic' oxidising power of
nitrates.
Also, why ferric nitrate hydrate? Why not Al nitrate hydrate? |
Ferric nitrate solutions are highly acidic. Also ferric nitrate is not very stable, tending to decompose to ferric oxide/hydroxide and so
generating nitric acid, even on gentle heating. Plus, you're overlooking that I've had encouraging results with the chloride version.
As for the aluminium nitrate, in principle that could also work, but ferric nitrate is easier to obtain IMHO and not as aggressive as aluminium
nitrate.
[Edited on 24-2-2015 by deltaH]
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blogfast25
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| Quote: Originally posted by deltaH | Ferric nitrate solutions are highly acidic. Also ferric nitrate is not very stable, tending to decompose to ferric oxide/hydroxide and so
generating nitric acid, even on gentle heating. Plus, you're overlooking that I've had encouraging results with the chloride version.
As for the aluminium nitrate, in principle that could also work, but ferric nitrate is easier to obtain IMHO and not as aggressive as aluminium
nitrate.
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Al nitrate hydrate melts easily too, as well as being quite acidic too. All highly charged cations hydrolyse.
Test, test, test!
[Edited on 24-2-2015 by blogfast25]
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