teodor
National Hazard
Posts: 903
Registered: 28-6-2019
Location: Netherlands
Member Is Offline
|
|
Some experiments with ammonium persulfate (peroxydisulfate)
As I already said here and there, I am moving my lab and I don't know when I will complete the process. Obviously, moving things like a fume hood is
time consuming and I don't know when I can finish it. But I am eager to start using my new place, so I decided to start some relatively easy
water-based experiments there.
Usually I do a report only when experiment is completed (so many interesting experiments are still not published), but in this case I would like to
have an additional motivation for making the new lab running (and you can imagine all that setup of sinks, water supply, light, shelves, tables etc),
so I am starting this thread with some piece of information which I believe can spark your interest, so we can have some discussion here and may be we
will be able to do some experiments together, and during this time I will rebuild my lab step-by step, and it is much more fun than just set up
everything in the lab without doing actual research.
So, saying that I would now move your attention to some underrated compound of ammonia, namely ammonium persulfate (NH4)2S2O8. SM Wiki mentions 2
projects you can do with it: PCB etching and making a low explosive. But I am quite sure you can do much more than that. To start, I propose some old
article about properties of the persulfate in presence of a silver ion. Depending on conditions it can produce different oxidations product of
ammonia, starting from nitrogen and going as far as nitric acid. Oxidation of ammonia to nitric acid in a water solution is not what you have very
often with other oxidisers, so I think this compound really can have quite interesting properties worth to investigate.
So, let's start with the article:
[Edited on 28-11-2024 by teodor]
Attachment: c449b37950e25156a369977890593c0c_241128_202550.pdf (302kB)
This file has been downloaded 57 times
[Edited on 29-11-2024 by teodor]
|
|
|
Sir_Gawain
Hazard to Others
Posts: 437
Registered: 12-10-2022
Location: [REDACTED]
Member Is Offline
Mood: Stable
|
|
Ammonium persulfate seems to be relatively difficult to buy; is there an amateur-accessible way of making it?
“Alchemy is trying to turn things yellow; chemistry is trying to avoid things turning yellow.” -Tom deP.
|
|
|
Keras
National Hazard
Posts: 912
Registered: 20-8-2018
Location: (48, 2)
Member Is Offline
|
|
You can buy some here assuming you’re in the EU. Not restricted, not particularly difficult to source.
|
|
|
teodor
National Hazard
Posts: 903
Registered: 28-6-2019
Location: Netherlands
Member Is Offline
|
|
I think there are 3 ways, depending of what you have: time, platinum or H2O2.
The classical one is electrolysis of (NH4)2SO4 solution, but in all publications I am aware of the platinum electrode was used.
Mixing H2SO4 or its salt with H2O2 should also work but we need to do some more investigations here.
And the third way is to do experiments with different electrode materials to find a substitution for Pt in a classical electrochemical process.
I would stay with electrolisys because access to H2O2 is rather limited in some places, and (NH4)2S2O8 could be also used to make H2O2 itself.
[Edited on 29-11-2024 by teodor]
|
|
|
Bedlasky
International Hazard
Posts: 1240
Registered: 15-4-2019
Location: Period 5, group 6
Member Is Offline
Mood: Volatile
|
|
Sodium persulfate is used in etching electronics, I think it should't be a problem to get persulfate in most places. Not to mention that Na and K
persulfates are very stable (unlike ammonium which decompose with time).
[Edited on 29-11-2024 by Bedlasky]
|
|
|
RU_KLO
Hazard to Others
Posts: 216
Registered: 12-10-2022
Location: Argentina
Member Is Offline
|
|
One thing with ammonium persufate, its that it starts working when temperature reaches 50C or more.
At least thats my experience.
Go SAFE, because stupidity and bad Luck exist.
|
|
|
Keras
National Hazard
Posts: 912
Registered: 20-8-2018
Location: (48, 2)
Member Is Offline
|
|
I’m currently trying to oxidise naphthalene with ammonium persulphate at 60/70 °C. Let see if it works.
|
|
|
teodor
National Hazard
Posts: 903
Registered: 28-6-2019
Location: Netherlands
Member Is Offline
|
|
Quote: Originally posted by Keras  | | I’m currently trying to oxidise naphthalene with ammonium persulphate at 60/70 °C. Let see if it works. |
I didn't read about naphthalene, but phenols and glycols could be oxidised. AgNO3 is often required, depending of the reactions but some unable to
proceed without it.
|
|
|
Keras
National Hazard
Posts: 912
Registered: 20-8-2018
Location: (48, 2)
Member Is Offline
|
|
| Quote: Originally posted by teodor | | Quote: Originally posted by Keras | | I’m currently trying to oxidise naphthalene with ammonium persulphate at 60/70 °C. Let see if it works. |
I didn't read about naphthalene, but phenols and glycols could be oxidised. AgNO3 is often required, depending of the reactions but some unable to
proceed without it. |
Oh, okay for AgNO₃. TBH, I’m not sure of what I got. There was a distinctive yellow layer floating above the aqueous phase, and I would like to
ascribe it to glyoxal, which is apparently a side-product of naphthalene oxidation by potassium permanganate, so why not by other oxidising agents? I
couldn't properly isolate anything, but that might be due to my being clumsy at the workup. I’m currently testing Oxone™ instead of ammonium
persulphate, which is not really ideal because it is much less soluble. I had the same yellow layer, and now I have a bunch of crystals floating.
I’ll try to see if these are unreacted naphthalene or proper phthalic acid.
|
|
|
teodor
National Hazard
Posts: 903
Registered: 28-6-2019
Location: Netherlands
Member Is Offline
|
|
| Quote: Originally posted by Keras |
Oh, okay for AgNO₃. TBH, I’m not sure of what I got. There was a distinctive yellow layer floating above the aqueous phase, and I would like to
ascribe it to glyoxal, which is apparently a side-product of naphthalene oxidation by potassium permanganate, so why not by other oxidising agents? I
couldn't properly isolate anything, but that might be due to my being clumsy at the workup. I’m currently testing Oxone™ instead of ammonium
persulphate, which is not really ideal because it is much less soluble. I had the same yellow layer, and now I have a bunch of crystals floating.
I’ll try to see if these are unreacted naphthalene or proper phthalic acid. |
What do you try to get, phthalic acid, alpha-naphthol or alpha-naphthoquinone? What do you use as a solvent?
Well, I have no experience with naphthalene oxidation, but I can share some information I have:
1. Chromic acid ( it is often used with glacial acetic acid) -> mostly phthalic acid, alpha-naphthoquinone < 16%
2. Ceric sulphate in H2O + H2SO4 + acetic acide -> alpha-naphthol
3. Vanadium(V), acetic acid is required, have no information about yield.
4. Lead tetraacetate + acetic acid -> 1-cetoxynaphthalene (26% yield)
So, it looks like acetic acid is required with all 4 oxidisers, the main compound and yield is dependent on the oxidiser and the conditions.
I can't contribute to it more. Usually when I try to use a new oxidiser I personally follow this path: repeat some well-documentet procedure, measure
and analyze result and then try my procedure with variations, comparing my result and behaviour with what I have observed with the known procedure.
|
|
|
charley1957
Hazard to Others
Posts: 164
Registered: 18-2-2012
Location: Texas
Member Is Offline
Mood: Beginning to cool off
|
|
Ammonium persulfate is easily had from Amazon. Prices are all over the map, but I saw a 1K bottle for about $35.
You can’t claim you drank all day if you didn’t start early in the morning.
|
|
|
teodor
National Hazard
Posts: 903
Registered: 28-6-2019
Location: Netherlands
Member Is Offline
|
|
There is some controversy about it's shelf life. I have a plastic bottle for several years without any visible changes. So, it could be something with
its purity when it starts to decompose (transitional metals impurities?). I plan to do analysis, so I'll report how much persulfate is in my aged
sample.
|
|
|
woelen
Super Administrator
Posts: 8016
Registered: 20-8-2005
Location: Netherlands
Member Is Offline
Mood: interested
|
|
I have bad experience with storage of ammonium peroxodisulfate. I once purchased half a kilo of it from an electronics shop (PCB etchant). I stored it
in a well sealed container, but one year after that, the material had become a very hygroscopic mass, which had clumped together. But most annoyingly,
it hardly had any oxidizing power anymore, most of the persulfate had decomposed. On addition to water, I just got a strongly acidic solution of
ammonium bisulfate with only very little peroxodisulfate left.
Lateron, I again purchased (NH4)2S2O8, but again, this batch also became useless in two year's time. So, I decided not to buy this chemical again, I
now do not have any of it.
I do have Na2S2O8 and K2S2O8, and both of these store exceptionally well. Even after 10 years of storage, these still are as if I purchased them
recently.
I also purchased oxone (a triple salt of KHSO5, K2SO4 and KHSO4) and this also stores quite well. It, however, has properties, very different from
peroxodisulfate. It is a weaker oxidizer, but if it oxidizes, then it is much faster. The two also are reacting with each other. They also react with
hydrogen peroxide. All three oxidizers have different (and incompatible) properties. I never fully investigated the precise differences, it is one of
the things I consider doing in the near future.
|
|
|
teodor
National Hazard
Posts: 903
Registered: 28-6-2019
Location: Netherlands
Member Is Offline
|
|
I remember you telling that, woelen.
My sample (which is still looking well) is from deoplosmiddelspecialist.nl (now duchchem.nl). Taking in account that decomposition occures only with
ammonium salt I would suspect that ammonia oxidation is the reason of its instability but it is not always triggered (storage conditions, purity,
crystall type?).
The different possible types of crystall structure were reported by Hugh Masrhall and there is one which adsorbes solvent is known to be very
unstable.
Well, personally I'd like to concentrate on investigation of ammonia oxidations paths depending on the conditions. I think if we will understand that
well we can understand the reasons of its decompositions better. If you have a decomposed samples, is there a chance to investigate them for NO3-
ions? Because it can mean the decomposition of a cation (another option is NH4+ -> N2 but it's impossible to detect, also I think the NO3-/N2 paths
could happen in some proportions).
As for oxone, do you think would it be possible to separate H2SO5 as an acid solution or its salt from the mix?
[Edited on 2-12-2024 by teodor]
[Edited on 2-12-2024 by teodor]
|
|
|
Keras
National Hazard
Posts: 912
Registered: 20-8-2018
Location: (48, 2)
Member Is Offline
|
|
My ammonium persulphate is more than a couple of years old and albeit it is stored in an outhouse which can be very dank during winter, it hasn't
clumped (probably the bottle in which it is stored was correctly sealed). I still have to test its oxidising properties, though.
[Edited on 2-12-2024 by Keras]
|
|
|
teodor
National Hazard
Posts: 903
Registered: 28-6-2019
Location: Netherlands
Member Is Offline
|
|
| Quote: Originally posted by Keras | My ammonium persulphate is more than a couple of years old and albeit it is stored in an outhouse which can be very dank during winter, it hasn't
clumped (probably the bottle in which it is stored was correctly sealed). I still have to test its oxidising properties, though.
[Edited on 2-12-2024 by Keras] |
You mean you plan to use it in some organic reaction with unknown result like naphthalene oxidation  It's not a real test for oxidation properties because we have no idea how it should work in this case.
So, it is possible only to test it with some well-known oxidation. Like oxidation of oxalic acid with presence of catalitic amount of Ag+ ions. It
will bubble CO2. You can even titrate the rest of oxalic acid with permanganate solution to find how mush persulfate ion was there. (This is actually
what I plan to do with my sample).
|
|
|
teodor
National Hazard
Posts: 903
Registered: 28-6-2019
Location: Netherlands
Member Is Offline
|
|
It's also well known that warm peroxydisulfate solution can dissolve different metals, for example, Cu. You need either nitric acid for that or
persulfate, and the persulfate is more affordable nowadays.
I would like to experiment with different metals to see what could be dissolved.
And dissolving Cu (a known expected result) could be also used to measure of persulfate sample quality (but not quantitative comparint to H2C2O4
method I believe)
@woelen: How do you think, can we start with the metal dissolving experiments as a first line of experiments to compare S2O8(2-), SO5(2-) and H2O2
oxidation properties?
[Edited on 2-12-2024 by teodor]
|
|
|
RU_KLO
Hazard to Others
Posts: 216
Registered: 12-10-2022
Location: Argentina
Member Is Offline
|
|
| Quote: Originally posted by teodor | It's also well known that warm peroxydisulfate solution can dissolve different metals, for example, Cu. You need either nitric acid for that or
persulfate, and the persulfate is more affordable nowadays.
I would like to experiment with different metals to see what could be dissolved.
And dissolving Cu (a known expected result) could be also used to measure of persulfate sample quality (but not quantitative comparint to H2C2O4
method I believe)
@woelen: How do you think, can we start with the metal dissolving experiments as a first line of experiments to compare S2O8(2-), SO5(2-) and H2O2
oxidation properties?
[Edited on 2-12-2024 by teodor] |
That will be a good idea (It was lingering in my mind, to make a list of inorganic oxidizers and reducers).
Also the conditions should be also be stated (for example, H2O2 is oxidizer in alkaline media, but reductor on acidic media).
Go SAFE, because stupidity and bad Luck exist.
|
|
|
teodor
National Hazard
Posts: 903
Registered: 28-6-2019
Location: Netherlands
Member Is Offline
|
|
| Quote: Originally posted by RU_KLO |
Also the conditions should be also be stated (for example, H2O2 is oxidizer in alkaline media, but reductor on acidic media).
|
Good point.
H2O2 is H-O-O-H and S2O8(2-) is -O-S(O2)-O-O-S(O2)-O- . So, the same -O-O- group. Could it be converted to a reducer the same way?
|
|
|
DraconicAcid
International Hazard
Posts: 4339
Registered: 1-2-2013
Location: The tiniest college campus ever....
Member Is Offline
Mood: Semi-victorious.
|
|
I doubt it. When hydrogen peroxide acts as a reducing agent, it gives off oxygen gas (and will only do this in the presence of a strong oxidizing
agent, regardless of the acidity of the conditions). For peroxydisulphate to act as a reducing agent, it would have to give off neutral S2O8, which
isn't going to be very stable.
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
|
|
|
teodor
National Hazard
Posts: 903
Registered: 28-6-2019
Location: Netherlands
Member Is Offline
|
|
I see. Also I think giving off one oxygen atom in the case of H2O2 means breaking a (weak) O-O bond and letting one proton to go. In the contrary
-S(O2)-O- group is not as mobile as a single proton, so it stays with the middle oxygen atom. Together this is actually SO4(2-) group which is known
to be stable. After -O-O- breaks SO4(2-) badly requires 1 proton to become an HSO4- ion, but it can take it only from water. But I think actually
water can break the middle bond giving off 2 protons and an atomic oxigen. Which is why (as I suppose) the persulfate is a more strong oxidiser than
H2O2.
[Edited on 3-12-2024 by teodor]
Oh. It's not SO4(2-), it is SO4(-1). Eager for H atom with an electron. Leaving oxygen from water in the perfect state of a single arom.
I doubt I can guess what is really happening, but it is my first try.
Well, we can probably check how water increasing the oxidation properties of persulfate.
[Edited on 3-12-2024 by teodor]
Of course I was not right with the single oxigen atom. At the present moment I have the information that oxidation potential and kinetics is depending
on the "activation" mechanism of persulfate. There are several ways, for example:
S2O82- -(heat)-> 2SO4- (the sulfate radical, reduction potential +2.6)
2S2O82- + 2H2O -(basic media)-> SO4- + 3SO42- + 4H+ + O2- (the superoxide radical, reduction
potential -2.4 - what??)
S2O82- + Fe2+ -> SO4- + SO42- + Fe3+ (usually with EDTA to increase the radicale production
by Fe chelation)
SO4- + H2O -> SO42- + H+ + OH (the hydroxyl radical, the reduction potential +2.8 )
For comparison: Persulfate anion itself - +2.1, Ozone - +2.1, H2O2 -> +1.77, Permanganate: +1.7. I believe all those data is for aqueous media.
So, we can really tune it for different oxidation strength.
[Edited on 3-12-2024 by teodor]
|
|
|
Keras
National Hazard
Posts: 912
Registered: 20-8-2018
Location: (48, 2)
Member Is Offline
|
|
Also the sulphate radical is not SO₄⁻, but SO₄• 
|
|
|
![[*]](images/xpblue/default_icon.gif){kind=icon}
