JefferyH
Hazard to Self
Posts: 97
Registered: 7-5-2014
Member Is Offline
Mood: No Mood
|
|
Would Persulfate adequately oxidize bromide?
I'm searching for a good substitute for oxidizing halides, to replace Oxone. Oxone has such horrible solubility, so I am searching for something new!
Much of what I have found online about Persulfate has to do the usage of Persulfate in swimming pools. The consensus among the pool-community is that
persulfate is not a good oxidizer of bromide, but only because persulfate does not decompose unless it is activated by heat. When heated, one
persulfate decomposes into two sulfate radicals.
1) Will these sulfate radicals react with water and auto-neutralize, reducing good yields?
2) Each sulfate radical can only oxidize one electron, whereas oxone oxidizes two. Would these sulfate radicals adequately take both electrons from
bromide in a two-step fashion?
I've read that persulfate is very exothermic, but I have only heard this in the context of it being used to oxidize alcohols. Does anyone experience
using it?
I want to weigh the advantages and disadvantages of Oxone vs Persulfate. If persulfate is uncontrollably exothermic in general then I might just stick
with oxone.
|
|
|
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
There are multiple threads here on oxidising bromides, so search for them.
What's the desired end product: bromine? Bromates?
|
|
|
woelen
Super Administrator
Posts: 7977
Registered: 20-8-2005
Location: Netherlands
Member Is Offline
Mood: interested
|
|
You must be more specific on what you mean with "persulfate". There are multiple ions, which could pass as "persulfate" ion:
- S2O8(2-), peroxodisulfate. This is available in pure form as the sodium salt, potassium salt (which is only sparingly soluble) and the ammonium salt
(which only has limited shelf life).
- SO5(2-), monoperoxosulfate. This is one of the ions in "oxone", which is a triple salt of KHSO4, K2SO4 and K2SO5. Oxone is much more soluble than
potassium peroxodisulfate, and approximately as soluble as Na2S2O8 and (NH4)2S2O8.
Both oxone and peroxodisulfates are capable of oxidizing bromide to bromine. Peroxodisulfate, however, is sluggish at room temperature and for a
decent reaction speed you either need to heat this to 60 C or so, or you need to add a catalyst, being Ag(+) ion. Oxone reacts quickly. If you add
solid oxone to hydrochloric acid, then it bubbles, releasing chlorine gas. So, I expect oxidation of bromide to occur even faster.
|
|
|
JefferyH
Hazard to Self
Posts: 97
Registered: 7-5-2014
Member Is Offline
Mood: No Mood
|
|
Quote: Originally posted by woelen  | You must be more specific on what you mean with "persulfate". There are multiple ions, which could pass as "persulfate" ion:
- S2O8(2-), peroxodisulfate. This is available in pure form as the sodium salt, potassium salt (which is only sparingly soluble) and the ammonium salt
(which only has limited shelf life).
- SO5(2-), monoperoxosulfate. This is one of the ions in "oxone", which is a triple salt of KHSO4, K2SO4 and K2SO5. Oxone is much more soluble than
potassium peroxodisulfate, and approximately as soluble as Na2S2O8 and (NH4)2S2O8.
Both oxone and peroxodisulfates are capable of oxidizing bromide to bromine. Peroxodisulfate, however, is sluggish at room temperature and for a
decent reaction speed you either need to heat this to 60 C or so, or you need to add a catalyst, being Ag(+) ion. Oxone reacts quickly. If you add
solid oxone to hydrochloric acid, then it bubbles, releasing chlorine gas. So, I expect oxidation of bromide to occur even faster.
|
Ammonium Persulfate has a limited shelf life? To what extent? It is one of the most commonly sold etchants in quantities of up to 25 kg + for resell.
What is it that limits the shelf life if it is sold in such quantities?
My only gripe with Oxone, which is only available as the triple salt, is that its solubility is about 5-10x worse than an oxidizer such as ammonium
persulfate. Oxone requiring about 1.5 - 2L of water per active mole of substrate, whereas ammonium persulfate may only need 200-300mL of water per moL
of substrate. Both are fairly cheap. I just dislike working in large flasks for such small reactions
| Quote: Originally posted by blogfast25 | There are multiple threads here on oxidising bromides, so search for them.
What's the desired end product: bromine? Bromates? |
I've read them a lot of them. Not many have to do with oxidation of bromide by (ammonium/sodium/potassium) persulfate. I am trying to generate the
same product as the oxone oxidation [Br+]. It seems persulfate can oxidize bromine, but now I am wondering if persulfate will effect the substrate
negatively. I recall reading in one that that persulfate can hydroxylate aromatic rings.
|
|
|
woelen
Super Administrator
Posts: 7977
Registered: 20-8-2005
Location: Netherlands
Member Is Offline
Mood: interested
|
|
Ammonium persulfate decomposes. Two years ago or something like that, I purchased 500 grams each of Na2S2O8, K2S2O8 and (NH4)2S2O8 from a UK-based
eBay seller (pure-domestic-chemistry, or minerals-water, or something like that). The chemicals were sold in similar bottles like this:
http://www.ebay.nl/itm/500g-Ammonium-Persulfate-persulphate-...
The K-salt and Na-salt still are as if I purchased them yesterday, the ammonium persulfate has deteriorated so much that it is nearly useless. It is a
wet mass, which nearly lost all its oxidizing power and just has become very wet ammonium bisulfate. I decided to not buy the ammonium salt again, it
is a waste of money, unless you use it up within a year or so. The material is so bad, that I discarded it.
I agree that oxone is not a very concentrated oxidizer. It is mostly sulfate and potassium ions and only contains a little amount of "oxidizing
oxygen". But the same, to a lesser extent, is true for the peroxodisulfates.
|
|
|
|
![[*]](images/xpblue/default_icon.gif){kind=icon}
