gardenvariety - 19-1-2010 at 10:44
Hi, I am attempting to use Oxone as an oxidizer for aromatic bromination. I've found one supply of OTC Oxone, in a pool supply "non-chlorine
oxidizer" that is labelled as 47% Oxone (in the triple salt) and the remainder is sea salt. I assume that using this directly would generate
chlorinated aromatics, so I want to remove the sea salt. Has anyone attempted this?
My first guess is to dissolve out the Oxone in an alcohol, leaving the bulk of the sea salt behind, but I can't find a reported solubility for Oxone
in anything besides water.
Thoughts?
bbartlog - 19-1-2010 at 11:26
I'd expect Oxone to oxidize alcohol. I don't know what solvent would be useful here. You could add H2SO4 and then draw off HCl vapor using some
combination of heat and vaccum until all of the chlorides were converted to sulfates (which should not interfere with your oxidation); but I'm not
convinced that it's necessary - do NaCl and MgCl2 really chlorinate aromatics? Seems unfavorable somehow, oxidizer or not.
gardenvariety - 19-1-2010 at 12:08
From my research, the same mechanism that drives KBr to brominate aromatics drives chlorides to chlorinate aromatics. Which is a problem, as aryl
chlorides are largely inert to alkoxide displacement.
bbartlog - 19-1-2010 at 13:16
Hm, yes - I was thinking that you would need to add H2SO4 to liberate the halogen but since oxone is an acid salt I suppose it would work as well,
forming MgSO4 and Na2SO4 while leaving the chlorine free to chlorinate. This makes me wonder about the stability of the commercial product you
describe, though. Wouldn't it give off either chlorine or HCl via something like
2KHSO5 + 2NaCl -> O2 + K2SO4 + Na2SO4 + 2HCl ?
I mean maybe it does do that when you add water, I'm unfamiliar with the product, but if it isn't giving off HCl and/or Cl2 in aqueous solution I'm
not sure it would be able to chlorinate, either.
Nicodem - 19-1-2010 at 13:48
Cl2 has a higher oxidation potential than Br2, so no relevant chlorination will occur using the product mixture to be the oxidant for KBr. You can not
purify it anyway, because as soon you will dissolve it in water it will slowly start to release chlorine (unless bromide or any other more easy to
oxidise ion is present). If you can not find pure Oxone, then why don't you use some other oxidant like H2O2 for example?
gardenvariety - 19-1-2010 at 14:16
I need to go back to my pchem books and bone up on oxidation potentials. Do oxidations proceed stepwise from lowest to highest through available
substrates, or do the various substrates get oxidized in parallel, with percentages according to potentials and mole qty of each? If the latter,
what's the math to compute that balance? For that matter, what's the math to compute how fast each reaction proceeds?
This OTC formulation might be bunk for my purposes. There's another pool chem out there that is Oxone and MgCO3, which is more or less insoluble in
water. If not that, then yes, time to look for pure.
I also found this in the European Journal of Org Chem, 2002:
"Preparation of Purified KHSO5·H2O and nBu4NHSO5 from Oxone by Simple and Efficient Methods"
...which describes a simple yet temp-and-pH sensitive process for preferentially dissolving out the KHSO5 from the triple salt using water and
methanol, then crystallizing it in the freezer. This might not work for the above formulation, but is worth a shot.
It also notes that the nBu4NHSO5 they form from that will do alkene cleavage into aldehydes rather than carboxcylic acids (as does oxone), which is a
nice extra tool in the kit.
gardenvariety - 19-1-2010 at 18:22
Ha! If one goes across the street to the competing hardware store, one finds (in the shop where, unlike its competitor, which sells NaBr), 85% oxone
with the remainder as NaCO3. One day this country will ban molecules, but until then...