Yes, except I cannot purchase aluminium nitrate but I do have ferric nitrate and I've already tested it with a chloride. Don't have a bromide salt
though, nor am I kitted out for safely conducting such a test at the moment.blogfast25 - 24-2-2015 at 09:34
Excuses, excuses... woelen - 24-2-2015 at 11:30
Hypochlorite, mixed with Mg, using as oxidizer does not sound good at all to me. Hypochlorites like Ca(OCl)2 decompose at low temperatures, giving the
chloride and oxygen gas. Mixes of these with Mg could be quite energetic. I tried such mixes with red P and with aluminium and they are very sensitive
and dangerous. I expect the mix with Mg to be very sensitive as well.
What I could imagine better is using Ca(OCl)2, mixed with some acid and a bromide, maybe NaHSO4 as acid. With the acid, you can make the equation
balanced.
In practice, however, I hardly think that any of the methods, presented at the last two pages in this thread, are useful for production purposes.
Maybe interesting from an academic point of view, but not really practical.blogfast25 - 24-2-2015 at 12:42
What I could imagine better is using Ca(OCl)2, mixed with some acid and a bromide, maybe NaHSO4 as acid. With the acid, you can make the equation
balanced.
That was mentioned earlier, what's the equation here? Wouldn't an acid just protonate the bromide? I can't think of how an acid can literally replace
magnesium in the reaction I posted.
The best I can come up with is:
I don't think this will work, and really all it does is complicates my equation by adding water to the the products list and still requires a reducing
agent to slurp up the oxygen.
Also using sodium bisulfate would have an even worse effect by requiring sodium hydroxide as another product, which I believe would react with
bromine, further complicating everything.
So am I missing something?
[Edited on 4-3-2015 by Molecular Manipulations]woelen - 4-3-2015 at 01:28
This reaction really works. Dissolve some NaHSO4 and NaBr in water. Make a slurry of Ca(OCl)2 in a separate container and mix the liquids. You get a
cloudy liquid with CaSO4 and copious amounts of bromine vapor and possibly also some liquid bromine.
Of course, if you have H2SO4, then you could do the following:
This reaction also works well in practice.Molecular Manipulations - 4-3-2015 at 06:27
Well I clearly didn't spend enough time trying to balance that one.
I must have been tired because now it seems obvious that it would work, as it could easily be analogous to using calcium hypochlorite to oxidize
chloride to chlorine, which would be reduced again by bromide, although it might not occur with those steps. And you can always use more acid to react
with the sodium hydroxide - duh!
Thanks for clearing that up man.
Still as long as we're back to using acid, there's better oxidizing agents. blogfast25 - 4-3-2015 at 06:44
Well I clearly didn't spend enough time trying to balance that one.
The first one in woelen's last post is balanced. It opens up the possibility of dry distilling Br vapour from a mixture of dry solids.
[Edited on 4-3-2015 by blogfast25]Molecular Manipulations - 4-3-2015 at 07:03
Yeah, but water is also a product.
[EDIT] Well now that I think about it, unless you really overshoot the temperature it should remain dry via the formation of sodium sulfate
decahydrate, which decomposes completely at 100°C.
[Edited on 4-3-2015 by Molecular Manipulations]woelen - 4-3-2015 at 10:31
I did the reaction in concentrated aqueous solution. If you use dry material I think that it works as well, but probably you do have quite a few
losses, because the solid-solid interaction only occurs at the boundary of particles. I do not know how well the reaction proceeds as a dry solid one,
but I am quite sure that mixing these three chemicals (previously finely and separately powdered by crushing them) will lead to formation of Br2.
Please try at a small scale first, reactions with hypochlorites can be quite exothermic and dangerous.
