Petit Homme
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Dissolving Selenium in Water with H2O2
I tried to dissolve a (black, vitreous) piece of selenium in water. I put it in 30% H2O2 under stirring assuming the highly toxic Selenic acid would
form. It doesn't dissolve however.
I don't have the exact solubility data for H2SeO4 other than that it's supposed to be extremely soluble. There is way enough water to dissolve the
SeO2/H2SeO3 that would form however.
What is happening?
Would it have dissolved if adding H2O2 drop-wise over time to the piece of selenium in the same amount of water?
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j_sum1
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Why would you expect selenium to be water soluble?
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DraconicAcid
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I don't think they expected selenium to be water-soluble; I think they expected the selenium to react with the hydrogen peroxide to give a soluble
product.
I don't think the reaction of selenium with hydrogen peroxide is going to be facile at room temperature. It doesn't strike me as a rapid reductant.
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.
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Fery
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at first step you need HNO3 dissolve Se, then you can let the intermediate react with H2O2
Se + ... HNO3 -> H2SeO3 + ... NOx
H2SeO3 + H2O2 - > H2SeO4 + H2O
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Petit Homme
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Quote: Originally posted by DraconicAcid  |
I don't think they expected selenium to be water-soluble; I think they expected the selenium to react with the hydrogen peroxide to give a soluble
product.
I don't think the reaction of selenium with hydrogen peroxide is going to be facile at room temperature. It doesn't strike me as a rapid reductant.
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That is correct. Thank you very much for contributing.
I left the solution under stirring within the of 50-70C range (I didn't want to go higher than that for two reasons: the first is that I want to avoid
as much as possible getting any selenium compound in the air or outside the vessel, which is loosely covered but solution could spill onto the
hotplate; the second reason is that I want to avoid H2O2 decomposition as much as possible).
Initially it seemed to me the elemental selenium was reacting (as per the color of the surface, but it might have been an illusion from looking at it
through the liquid). Later on it seemed to have stopped reacting (a darker color of the surface, which may also have been an illusion). It almost
seemed that adding more reactant or more solvent would make it dissolve. I figured I must have underestimated H2O2 decomposition so I added H2O2 which
seemed to help but it later seemed that if it did help it then was the H2O part of it and not H2O2 being a limiting reagent, as I ended up with a
solution in which (still now) tiny bubbles of O2 can be seen sparkling (as reasonably concentrated H2O2 looks like with any metal inside).
I can't add any more water now as my vessel is reasonably full within the limits of the volume I want to have in there. But if the solvent isn't the
problem (which according to the solubility data, it shouldn't be), I could carefully let the solution concentrate without boiling, which should
increase the reaction speed. Do you think this could cut it?
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Petit Homme
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| Quote: Originally posted by Fery | at first step you need HNO3 dissolve Se, then you can let the intermediate react with H2O2
Se + ... HNO3 -> H2SeO3 + ... NOx
H2SeO3 + H2O2 - > H2SeO4 + H2O |
Thank your for contributing! 
I don't have any HNO3 and it's not immediately available where I'm at. I know I could make some but if I could avoid and just use H2O2 as an oxidizer
that would be great.
My goal is to end up with H2SeO3 not H2SeO4. I figured H2SeO4 would form in a H2O2 solution but I initially quantified the H2O2 to be the limiting
reagent, thinking Se would reduce H2SeO4 in the formation of additional SeO2 (then reacting with water to form H2SeO3) as long as Se is present in the
reaction medium. Please tell me if that assumption is correct, or whether it is wrong.
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bnull
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You may have a hard time making selenious acid. According to Brauer, it is reduced even by dust.
You need selenium dioxide. There are two methods in p.428 of Brauer's Handbook of Preparative Inorganic Chemistry. The first requires burning selenium
in a mixture of oxygen and nitrogen dioxide (by passing oxygen through fuming nitric acid), which doesn't seem a good idea for an amateur. The second
employs concentrated nitric acid to oxidize selenium and produces nitrogen dioxide. Either way your solution turns out contaminated with elemental
selenium because, again, the acid is reduced even by dust. Same situation of sodium selenite.
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teodor
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You have a quick eye, Petit Homme. By the action of hot water selenium changes its allotrope form and becomes darker.
HNO3, H2SO4 (concentrated) or Cl2, those can oxidise Se (according to Treadwell), but I didn't try it. All 3 are quite essential chemicals in any
laboratory. May be hypochlorite can oxidise it also? Warning: I never did experiments with selenium, please check / google the possible reactions
before trying.
Warning 2: water soluble selenium compounds have similar toxicity to cyanides. Based on this I would recommend perform more experiments with basic
chemicals (including HNO3) before attempting do something with selenium.
P.S. Sodium sulfide solution can dissolve Se according to https://woelen.homescience.net/science/chem/exps/se_chemistr...
as well as peroxydisulfate can dissolve some selenium allotropes at high temperature.
[Edited on 2-9-2025 by teodor]
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bnull
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You're not trying to make Meker/Mohnke/Mencke/Mecke reagent, are you? It is the only use of selenious acid that I can think of and one reason for you
to try to purify selenium from 5N to 6.5N. If that is indeed the case, the presence of transition metals won't be a problem, especially when they're
present at 10 ppm.
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Petit Homme
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Well, according to teodor and bnull's contributions it appears Fery's method (using HNO3) then may be the best way to go after all. So I'll have to
make some. I've wanted to avoid that for many reasons, but there's also many uses of HNO3 I've postponed because of this. Making some HNO3 to use will
speed up a lot of things.
The good side is that (I believe) it's easier to make superclean HNO3 than superclean H2O2 that won't introduce metal impurities in the solution per
se (though I'll probably need to leave the glassware soaked in HNO3 inside before using the glassware with HNO3).
For Cl2 I don't see how I could use that as "safely" as I would HNO3 (which to me is already dangerous enough to use that I've avoided it where
possible).
H2SO4 would be on par with HNO3 safety-wise (if having to go through distillation, though probably much safer to use once clean), however I don't
think I could get H2SO4 to be anywhere as clean as HNO3 as easily (trace metal basis). Also the end product I want is selenous acid (H2SeO3), not
selenic acid (H2SeO4). And while H2SO4 indeed does oxidate Se, I think it oxidates it all the way to Selenic acid, while HNO3 may oxidate it only to
selenium dioxide, turning to selenous acid in water (this is my current understanding, please correct me if I am wrong). Also it evolves SO2 which,
according to wikipedia, has a fair solubility in water (much greater than NO), and which I'd have to get rid of somehow. Worse yet is the remaining
H2SO4, which I won't get rid of anywhere as easily as I may cleanly get rid of remaining HNO3.
Traces of elemental selenium is not a problem for me, especially if its tiny particles, and if it's a minor portion of the total selenium in solution.
I'm concerned about resulting with extremely low metal traces after recrystalization (if its the most simple and effective way to achieve this
starting from 5N Se), about not ending up with any H2SeO4 in solution, about not getting selenium in the air or all over the place.
The reason for doing this is twofold: (1) a means to purify Se from metal traces as far as reasonably possible with fairly simple means, (2) ending up
with a selenium compound that is soluble and non-toxic to human/animal health in minute amounts (see RDA range).
As per HNO3:
Can I use it (fairly concentrated) onto Se in an aqueous solution such as to directly form H2SeO3?
Is it correct that it will only oxidate Se to SeO2 (forming H2SeO3 in aqueous solution) and not further to H2SeO4?
Will it evolve NO or NO2 in the process? Or both? If correct NO2 reacts with water to form back HNO3 and NO. NO has fairly low solubility in water
compared to SO2, and the solubility decreases with temperatures, though I expect some to remain of it, as for HNO3. How do I get rid of that?
Evaporating to dryness? Or just filtering off crystalized Se salt from solution after reducing the water volume by sub-boiling evaporation? (I'll have
to lose Se salt any ways at each crystallization round if that route is the most simple and effective way to rid metal traces starting from 5N Se).
[Edited on 3-9-2025 by Petit Homme]
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bnull
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| Quote: | As per HNO3:
Can I use it (fairly concentrated) onto Se in an aqueous solution such as to directly form H2SeO3?
Is it correct that it will only oxidate Se to SeO2 (forming H2SeO3 in aqueous solution) and not further to H2SeO4?
Will it evolve NO or NO2 in the process? Or both? If correct NO2 reacts with water to form back HNO3 and NO. NO has fairly low solubility in water
compared to SO2, and the solubility decreases with temperatures, though I expect some to remain of it, as for HNO3. How do I get rid of that?
Evaporating to dryness? Or just filtering off crystalized Se salt from solution after reducing the water volume by sub-boiling evaporation? (I'll have
to lose Se salt any ways at each crystallization round if that route is the most simple and effective way to rid metal traces starting from 5N Se).
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See the first volume of Inorganic Syntheses, chapter 6. But be warned that no home process can purify selenium the way you want. That's the good, old,
plain boring truth.
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teodor
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Indeed, making high purity matherial is time consuming and unreliable without practice and repetitions. I believe the most important is to check the
purity of your product because usually you can contaminate it even without noticing. I think it is wise to assume your result could be as much pure as
much your testing method is sensitive.
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Petit Homme
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I'm writing back here to report that – unlike what the quoted literature seems (at least by omission) to suggest – what I was initially going for
(that is: Dissolving Selenium in Water with H2O2, as the title of this thread reads) does work, within a long but very much
practicable timeframe.
After posting here I left the aforesaid beaker as it was, unattended, covered, at room temperature. About a month later, looking into the beaker, I am
surprised to see that the piece of elemental selenium is gone, leaving a completely transparent (still sparkling) solution of water, H2O2, and (I
assume must be) dissolved H2SeO4, as well as a tiny very thin remnant skeleton (think the appearance of a lose silk mesh), of what I assume may have
been impurities in the Se piece (which may or may not have been as pure as the 5N grade it was sold as), since I fail to see it disappear.
In any case: Oxidizing elemental selenium with a solution of H2O2, alone, works. It just takes a very long time and (at least in this case) was (has
to be?) done with a generous excess of H2O2.
My specific purpose being beyond the scope of this post – not just to form a selenium compound that dissolves in water but to form one that is a
fair candidate to purify the selenium from select metal impurities via recrystallization in deionized water, and to end up with an SeO2/H2SeO3 aqueous
solution – I initially wanted to use stoichiometric amounts of H2O2 over the course of the reaction/dissolution (which I wrongly assumed would occur
much faster that it ended up doing), complementing towards the end of the reaction/dissolution to compensate for unavoidable H2O2 decomposition. The
purpose behind this desire to use limiting amounts of H2O2 was the idea that the unreacted selenium would reduce any H2SeO4 that would form such that
I may end up with mainly H2SeO3. The reason for this being twofold: First, I have some (partial) solubility data for SeO2/H2SeO3 which (in theory)
makes it seem it could be a form of selenium used to purify it from select metals via recrystallization in deionized water, while I lack specific
numbers on H2SeO4 water solubility vs temperature. Second, I wanted to avoid having to deal with H2SeO4 altogether, as far as possible, since it is a
degree of magnitude more toxic (and thus greater of a hazard to handle) than H2SeO3 already is and since I'd like to end up with a very pure (trace
metal basis) SeO2/H2SeO3 solution anyways.
As the case went, I very likely have a beaker full of H2SeO4 (in addition to remaining H2O2 and water). I now wonder how / if I can slowly reduce this
to H2SeO3 while introducing as little undesirable metal impurities as possible. It's fine if I get a little bit of reduced elemental selenium dust
along the way, as long as the bulk is H2SeO3, that I introduce as little trace metals impurities as possible, and that I may recrystallize the H2SeO3
afterwards. Does anyone have ideas about the best way to achieve this? if it is at all achievable?
I have small amounts of very pure ascorbic acid available, for instance, if it would be a good idea to use it. Would it work if I added a very dilute
solution of that, drop-wise, under good stirring of the H2SeO4 + H2O2 water solution? To add as little ascorbic acid as possible, maybe I could leave
the beaker at 75C for a while to thermally decompose as much H2O2 as possible beforehand?
All constructive ideas / contributions to achieve the desired purpose are welcome!
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bnull
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One thought that keeps pestering me is, why do you need such degree of purity for selenium? I mean, it's none of my business of course, yet I can't
think of any use for "virginal" selenium. Perhaps I'm getting dumber with age.
How pure is your ascorbic acid? How do you know there are no dissolved impurities in the selenium solution?
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Petit Homme
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| Quote: Originally posted by bnull | One thought that keeps pestering me is, why do you need such degree of purity for selenium? I mean, it's none of my business of course, yet I can't
think of any use for "virginal" selenium. Perhaps I'm getting dumber with age.
How pure is your ascorbic acid? How do you know there are no dissolved impurities in the selenium solution? |
Of course there are dissolved impurities in the selenium solution (which is very much expected and highly desirable for the purification route I am
going for).
The ascorbic acid has <5ppm of total trace metals, which is cleaner than the selenium solution on a total trace metal basis. But I want to use as
little ascorbic acid as possible anyways because individual trace metal composition in the ascorbic acid is different from that of the Se, and if it
indeed truly was 5N Se to begin with (which it might not have been), then individual trace metal levels in the ascorbic acid for some metals would
still be higher than their levels in the Se solution, so the ascorbic acid still will "contaminate" the solution on the basis of some specific
individual trace metals. Note that I do expect the H2O2 to have heavily contaminated the Selenium solution already, as it must have contained some
unknown stabilizer (Sn, or phosphoric acid with "heavy" amounts of alkaline elements, including traces of some alkaline elements – other than
K/Na/Mg which I am fine with – that are undesirable to me).
In any case. Further writing on this line would be wandering off the question:
How can I slowly reduce the H2SeO4 to H2SeO3 while introducing as little undesirable metal impurities as possible? Would it work to add a very dilute
solution of the ascorbic acid, drop-wise, under good stirring, to the H2SeO4 + H2O2 water solution? Would selenium dust formed in the process be
re-oxidized to H2SeO3 by H2SeO4 until H2SeO4 is mostly exhausted?
[Edited on 9-10-2025 by Petit Homme]
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