budullewraagh - 26-11-2004 at 14:22
i've always viewed chlorates and nitrates as being similar to one another, the chlorates being less stable and stronger oxidizers than the
nitrates but all in all still similar.
i know that nitric acid can be made using a nitrate salt and sulfuric acid, as the sulfate anion is more negative than the nitrate anion.
my question is can i make chloric acid using the same method? as well, what sort of tubing should i use to bring the HClO3 gas (if produced) to the
second beaker for to dissolve it? im thinking of using PTFE, but are there others?
thanks in advance
S.C. Wack - 26-11-2004 at 18:53
First, you have to realize that chloric acid is obtainable in solution only. And not real conc. soln, either.
Since it cannot be crystallized or distilled, you have to do things a certain way - you need insoluble byproducts. Barium sulfate or much less
preferably sodium oxalate would fit that description.
So one would dissolve Ba chlorate in water, chill, and add an equal amount of dil. H2SO4 slowly enough to not raise the temp.
BromicAcid - 26-11-2004 at 19:22
Even under incredible vacuum chloric acid will decompose under attempts to distill it, it can be concentrated once in solution under a heat lamp if
kept free of dust and other air born contaminates. I think the most concentrated a solution you can get is somewhere around 40 or 50%. The book
'Preparative Inorganic Chemistry' gives a preparation for bromic acid by using the barium bromate / sulfuric acid methodology, the
preparation of barium bromate being simple, but the prep in the book for barium chlorate being quite complicated, requiring NH4ClO3 as an
intermediate. Perhaps you could take a look at it, I think it's up on the FTP.
I've never worked with chloric acid, but bromic acid is surprisingly powerful and hard to predict, once it gets started oxidizing something it
starts to heat up and the ions put into solution from the dissolution of whatever metal and the heat cause the bromic acid to decompose faster and
faster releasing bromine fumes, I'm sure with chloric acid, chlorine oxides may well be released though a vigorous oxidation... sounds fun 
BTW: Mixing a chlorate with sulfuric acid usually results in the chloric acid thus formed decomposing into chlorine oxides which can explode readily.
[Edited on 11/27/2004 by BromicAcid]
S.C. Wack - 26-11-2004 at 20:16
I suppose I really should search my computer before posting - Brauer uses conc. H2SO4, and does it hot. I thought that the acid was less stable than
that. I'll look up the JACS and Z. anorg. refs to find out why they use such technique, no idea what IVA is.
budullewraagh - 26-11-2004 at 20:43
ah, one last thing. i know that sulfuric acid and nitrate salts yield aqueous NO2+. do sulfuric and chlorate salts yeild aqueous ClO2+? if so,
that's crazy stuff
JohnWW - 27-11-2004 at 00:59
An interesting possibility, ClO2+. Also ClF4+. I read somewhere that someone has made ClF6+, isoelectronic with SF6.
S.C. Wack - 28-11-2004 at 15:42
Looked up the refs while at the library like I said I would, but have little to say. I now know that the library is missing 11 years of Z. anorg., and
the JACS ref. was on iodic acid and anhydride. Their making of chloric acid (pretty much what is in Brauer) was on the side. That was what they mixed
with I2 to get their iodic acid. They gave no reason why they used boiling chlorate and hot H2SO4, but did say that they didn't think that it
hurt at all. It may be that their method was colored by their previous unsatisfactory experience making iodic acid from iodate/sulfuric which is slow.
Maybe they wanted chlorate concentration almost as high as possible.
Or maybe my previous advice is too sissy.
And like Gmelin's and the refs that it points to, everything else that I found was in German.
EDIT: I suppose I should mention other methods, since the thread is here. These are mentioned in both Gmelin's and Mellor.
From hydrofluorosilicic acid and Na, Ba, or K chlorate. Na chlorate and oxalic acid. K chlorate, Al sulfate, and H2SO4 - with EtOH precipitation of
the alum. (hypo)chlorous acid, ClO2, etc. of probably not much preparative value.
[Edited on 29-11-2004 by S.C. Wack]
kclo4 - 2-1-2005 at 13:40
http://www.sciencemadness.org/talk/viewthread.php?tid=12 reed this Perchloric acid preparation
evilgecko - 3-1-2005 at 11:56
IIRC, chloric acid can randomly detonate at any high concentration
Fluorite - 14-11-2020 at 08:02
Can chloric & perchloric replace nitric acid? Like oxidizing inositol to rhodizonic acid? Or sugar to oxalic acid? Can these acids dissolve gold?
Platinum silver copper?
ArbuzToWoda - 14-11-2020 at 08:15
Do your research. Metals: https://scholarworks.uni.edu/cgi/viewcontent.cgi?article=703...
Inositol: https://sci-hub.do/https://onlinelibrary.wiley.com/doi/abs/1...
Why am I even spoonfeeding you like that? Not sure.
Bedlasky - 14-11-2020 at 08:28
Perchloric acid doesn't have oxidizing properties, unless it have very high concentration.
Fluorite - 14-11-2020 at 13:06
Ahah Thank you ArbuzToWoda
Dr.Bob - 14-11-2020 at 17:05
Perchlorates can be powerful oxidants, including the acid, which is well known for causing explosions when mixed with fuels or certain chemicals.
Chlorates have an even more well known propensity to detonate with many other compounds, making the acid should only be done in a very controlled
manner behind a shield or in a hood. It does not play well with others. Chlorates are very useful, and the salts have many uses, but even then,
any real amount of them needs to handled in an area away from others if being mixed with fuels or for unknown reactions. I have seen people who
mishandled them, and it was not good.
woelen - 16-11-2020 at 00:22
Chloric acid ios a very strong oxidizer and it can react dangerously with many reductors. Unfortunately it is not a very "clean" oxidizer, in the
sense that it has a single reaction. It can produce ClO2, Cl2 and Cl(-) ions in its reactions, depending on pH, concentration, and the compound to be
oxidized. Often, it produces a mix of different compounds.
Aqueous perchloric acid usually is not strongly oxidizing at normal temperatures. Concentrated perchloric acid (70% by weight) in most cases still
only works like a strong acid, but when it is heated with certain reductors, it can become very reactive. But with many other reductors it hardly does
anything. I did experiments with hot 70% HClO4 on KI and Na2SO3 and with these, no redox reaction occurs (in the latter case, one simply drives off
SO2). With many metal powders, simply H2 is produced besides the metal perchlorate, but with bismuth metal, an explosion occurs. Bismuth oxide, on the
other hand does not lead to an explosion, one simply gets bismuth perchlorate.
When perchloric acid is diluted to 40% or so, then it is generally safe, much more so than 40% HNO3.