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Author: Subject: Stability of diluted chloric acid solutions
chaosday
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[*] posted on 8-1-2022 at 06:45
Stability of diluted chloric acid solutions


I recently made some dilute chloric acid, but I have nothing to use it for at the moment. So now I wonder if diluted solutions of chloric acid are shelf-stable. Does someone here know the answer to this question?
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[*] posted on 17-2-2022 at 13:38


of what ive heard it decomposes above 5*C, vague memory of mine recalls it being used in chemical operations up to 30% concentration- but they decided to leave it to the past since it had a tendency to violently detonate for reasons they probably didnt bother to figure out

if one could carefully decompose it, maybe with UV light, chloric acid would maybe be capable of turning into perchloric acid? which is also very dangerous and can form explosive salts with metals- theres some horror stories of people working in ventilation equipment suddenly realizing that metal salts formed from vapours of perchloric acid- and the vent itself has some properties they dont come across oftenly

as for decomposition above 5*C it should be Cl2 and maybe ClO2 i think? if youre really unlucky Cl2O




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[*] posted on 18-2-2022 at 00:03


Solutions of chloric acid are fairly stable at low concentrations at room temperature. No need to worry about detonations at concentrations below 25%. I have made such solutions myself and kept them around for a while, without problem. Just to be sure, keep it in a glass bottle, with its screw cap loosely attached. Assure that no dust can get into the bottle, but if any gas is produced, that it can escape. Chloric acid does not fume like HCl or HNO3, so it is not that bad on storage.

Any reducing matter causes instant reaction, so avoid that, use clean glass with a good plastic cap.

With concentrations above appr. 25% I would be more careful. Dilute those if you want to store them for a longer time.

[Edited on 18-2-22 by woelen]




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[*] posted on 18-2-2022 at 02:25


i don't know about dilute chloric acid, but i have 1 liter concentrated chloric acid which is made in 1981. it's very good. onced i precipitated caesium chlorate.



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[*] posted on 18-2-2022 at 06:29


so chloric acid may infact be an industrial secret?

@woelen
if chloric acid is used for oxidations, would it also chlorinate, and in case of just purely oxidations how strong would it be compared to for instance permanganate?

@vano can you test an approximate of how concentrated it is? chlorate is typically tested for by adding HCl to a sample and its fairly sensitive test




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[*] posted on 18-2-2022 at 07:57


i dont't remember exactly, if i remember corectly percentage was somewhere 65-70. if i will measure percentage again then i will write here. also my 1 liter bottle was unopend. it is in a transparent glass bottle, but it was in a dark room during decades.



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[*] posted on 19-2-2022 at 02:33


@vano: Are you sure it is chloric acid? That percentage you write about is common for perchloric acid. I think that chloric acid is not stable above 40% concentration.

@Antiswat: Permanganate is a stronger oxidant, but chloric acid can be obtained in higher concentrations in aqueous solution, and so the latter can be more violent. Chloric acid is not a "clean" oxidizer. It can oxidize in many different ways, leading to formation of many different species, e.g. ClO2, Cl2 and Cl(-). Especially if excess oxidizer is used, the reactions of chloric acid tend to be messy. In organic chemistry I don't think it is very useful, because it not only oxidizes, but also indirectly chlorinates (if Cl2 is formed). In certain inorganic oxidations, it can be quite useful, e.g. for dissolving gold in a mix of HClO3 and HCl. In that situation, intermediate formation of some Cl2 is not a problem.




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[*] posted on 19-2-2022 at 04:25


haha, yes woelen it's perchloric acid, sorry a small misunderstanding.



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[*] posted on 19-2-2022 at 05:37


very interesting, i looked around at YT yesterday for HClO3 and its mostly just people mixing 98% H2SO4 with chlorate to prove that its too dangerous to deal with
but absolutely nobody making a stable solution with it
i believe to use it for oxidations one would have to dump in something that very readily absorbs the chlorine? doable- from my currently ignorant perspective, maybe im just naïve. lets find out.

adding a base wouldnt do- because, its an acid.




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[*] posted on 19-2-2022 at 11:28


Pure HClO3 cannot exist, so if you add conc. H2SO4 to solid KClO3, then indeed you get instant decomposition and possible explosion. That is an insanely dangerous experiment. What can be done safely, however, is adding dilute H2SO4 to a saturated solution of Ba(ClO3)2, so that BaSO4 precipitates and HClO3 remains in solution. Such a solution even can be concentrated further, by letting water evaporate from it. I would not go further than 20% though.



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[*] posted on 19-2-2022 at 12:42


Typical dismutation reactions of overly concentrated HClO3 go like
3HClO3=HClO4+2ClO2+H2O.

This is actually how perchloric acid was discovered back in 19th century.

Which hydrates does chloric acid form? Because unless the acid precipitates as hydrates, freezing also concentrates the acid.
Under low temperatures, ClO2 is actually a liquid at temperatures as high as +11, and unpleasantly explosive when separated in bulk.
What happens when strong (30...40%) chloric acid is rapidly frozen, like being poured on dry ice? Does it dismute, or freeze without dismutation?
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[*] posted on 21-2-2022 at 01:12


Chloiric acid is a strong acid, and the stability of this acid depends on how well it is ionized in solution.

HClO3 is not stable and quickly decomposes to HClO4, ClO2 and water, and there also will be side reactions with formation of Cl2. If the acid is sufficiently dilute, then (nearly) all of it will be split in hydrated ions H(+) and ClO3(-). These chlorate ions are perfectly stable. This explains why concentrated solutions of HClO3 decompose, while dilute solutions are stable. From concentrations of appr. 30% and up, a considerable part of the acid is not split in ions anymore and molecules of HClO3 exist, which decompose. At concentrations below 20% practically all of the acid is split into ions.

For the same reason, HClO4 is stable up to appr. 70%. HClO4 is an even stronger acid than HClO3 and even at 70% concentration, this is fully split into hydrated H(+) and ClO4(-), making the solution stable. But the anhydrous acid is extremely dangerous and is unstable. It explodes on storage, within one month of its preparation.




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[*] posted on 21-2-2022 at 16:27


"It explodes on storage, within one month of its preparation." is this, kind of like demonstrated with pure ozone? i heard of one prankster that loaded about 100mL of concentrated HAc and, well probably not anhydrous HClO4 into an analysis machine at a university
the whole thing came apart

i have tried ascorbic acid with chlorate and perchlorate in solution, warm, and it reduces both
now i got to think- maybe these ions could be strong enough to without acidification oxidize some chemicals even? ascorbic acid and chlorate can selfignite at about 60*C or so- maybe it was at 80*C where ascorbic acid starts to decompose, making a composition with this is very difficult to do in same manner as "golden powder" where the 2 chemicals are dissolved in water and evaporated, because it ignites at such very low temperature
maybe chlorates/perchlorates would be able to act oxidizing in caustic environment, i believe thats the basis for "stoplight" "chemical chameleon" KMnO4 NaOH sugar experiment




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[*] posted on 22-2-2022 at 01:33


Perchlorate ion does not oxidize in aqueous solution, it is remarkably inert. At high temperatures (e.g. in fireworks), it is a strong oxidizer. Chlorate ion is strongly oxidizing in aqueous solution, but only at low pH.



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[*] posted on 23-2-2022 at 12:02


Quote: Originally posted by Antiswat  
"It explodes on storage, within one month of its preparation." is this, kind of like demonstrated with pure ozone?

The original paper of Roscoe on discovery of pure perchloric acid, in 1862, in Proceedings of Royal Society.

Perchlorate is a pretty inert oxidant. There are few simple reducers that manage to reduce perchlorate in cold dilute aqueous solution - these include Ti3+ and a few other metals, but even those react only slowly when not complexed. There do exist perchlorate reductase enzymes which manage to reduce perchlorate in mild conditions.
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[*] posted on 24-2-2022 at 09:44


well might be time for me to shoot this Roscoe guy a letter about ascorbic acid, i managed to reduce KClO4 in solution with it - this would maybe also go to show that ascorbic acid is quite powerful reducing agent, i managed to also reduce benzoic acid with it, not very efficiently, but its quite something.



~25 drops = 1mL @dH2O viscocity - STP
Truth is ever growing - but without context theres barely any such.

https://en.wikipedia.org/wiki/Solubility_table
http://www.trimen.pl/witek/calculators/stezenia.html
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