I have read many essays on how to destroy the chlorates and not the prechlorates unfortunately most had a great lack of the procedure and the one that
gives a procedure was too dangerous and costly. Certainly there is other ways.
If any one knows of a synthesis or a site pleas post
Thank youguy - 31-12-2004 at 00:03
Do you just want to seperate the two? If it were NaClO3 and NaClO4, according to this msds, you can have a solution of 209% NaClO4 in water, and NaClO3 only dissolves 100g in 100ml of water, so you can evevaporate the water to
isolate one of them.kclo4 - 31-12-2004 at 00:06
That would work but I do not know how much naclo3 is in my solutionBromicAcid - 31-12-2004 at 00:08
You could heat the chlorate to give perchlorate and chloride and then re-crystalize since the chloride is much more soluable than the perchlorate.guy - 31-12-2004 at 16:11
Quote:
since the chloride is much more soluable than the perchlorate.
That would only work for the potassium salt. The sodium perchlorate is highly soluble.
Solubilities
MadHatter - 1-1-2005 at 16:02
From CRC 62<sup>ND</sup> Edition(1981-1982) in grams per 100 ml H<sub>2</sub>O:
The NaClO<sub>4</sub> is the monohydrate: NaClO<sub>4</sub>.H<sub>2</sub>Okclo4 - 2-1-2005 at 00:57
I research and I thick I am going to destroy the chlorates with Sodium metabisulphite
The reaction that happens is
3Na2S2O5 + H2O + 2NaClO3 => 3Na2SO4 + 3H2SO4 + 2NaCl
Than I am going to and naoh to neutralize the acid
But I till don’t know how to get the pure naclo4 with out the other contaminations
Is any one knows a way or another reaction that creates a insoluble salt
And how can I make NaS2O5 or the acid. Could I just add h2o2 to H2SO4 I read that’s how you can make it but I don’t see howguy - 2-1-2005 at 01:34
You could fractionally crystallize all of them since NaClO4 will stay in the solution. (209%)garage chemist - 2-1-2005 at 05:02
You could add KCl solution, this would precipitate KClO4, which is highly insoluble in cold water.
KClO3 is also sparingly soluble in cold water, but not nearly as insoluble as KClO4. You should work in a highly diluted solution to keep the KClO3
from precipitating!
[Edited on 2-1-2005 by garage chemist]chemoleo - 2-1-2005 at 05:03
I think the idea is to first destroy the chlorate by chemical means, and then to extract the perchlorate by precipiation with a potassium salt.
The resulting KClO4 precipitates and is free of KClO3, if the chemical destruction of the chlorate was complete. I suggest to take stoichiometric
quantities of bisulphite vs the amount of NaCl you started off with, just to make sure you aren't left with KClO3 - particularly if you obtained
your solution of sodium (per) chlorate by electrolysis or heating of chlorate.
Hodges - did you try the heating method to disproportionate the chlorate to perchlorate and chloride? From what I read this requires to keep the temp
of the chlorate/perchlorate mix at the MP for several hours. If too high, the chlorate/perchlorate decomposes, if too low it solidfies, or the
reaction is slow.
It's more of a hassle than it looks. I'd also prefer chemical means, although I wouldnt be happy about the contamination with other salts.neutrino - 2-1-2005 at 06:34
Quote:
Originally posted by kclo4
And how can I make NaS2O5 or the acid. Could I just add h2o2 to H2SO4 I read that’s how you can make it but I don’t see how
This creates the peroxysulfate (SO<sub>5</sub>-) ion, not the one you’re after.budullewraagh - 2-1-2005 at 07:23
[nitpick]peroxysulfate is a -2 anion[/nitpick]hashashan - 16-5-2007 at 03:42
Say guys, i have a problem. NaS2O5 is quite expensive in my country and i dont want to use tons of it.
So i thought that ill use small quantities till there is no chlorate left. But another problem is that i dont have any indicators. what is the best
way to make an indictor or another method of safely destryoing the chlorate so i can make ammonium perchlorate?woelen - 16-5-2007 at 05:44
Acidify the liquid a little bit and make it warm (60 C or so) and add the Na2S2O5 in small steps. As soon as a clear smell of SO2 can be observed, you
have converted all chlorate. This test is quite sensitive, the smell of SO2 is strong and pungent, and even an excess of a single spatula of Na2S2O5
will cause a very strong and pungent smell.hashashan - 16-5-2007 at 06:08
ok, but what if the SO2 shuts down my smell sense? i once had this, almost killed myself with a reaction ( or was it H2S?)
and how good is this method? would you dare to make AP with this perch?12AX7 - 16-5-2007 at 13:55
You could also use a dash of Fe(III) sulfate. Chlorate should be reduced in favor of Fe(III) until all chlorate is gone.
Obviously, you don't put in so much Na2S2O5 that it gasses you. If it's so strong that you can't smell the shit anymore, chances are you're
done............
Tim
[Edited on 5-16-2007 by 12AX7]hashashan - 17-5-2007 at 03:34
Hi,
i saw on the geocities site a method for cheching for chlorates that uses MnSO4 and phosphoric acid. Anyone tried it, or care to explain in a fully
detailed manner how to preform this test, because for some reason it doesnt work for me.
Manganous Sulphate-Phosphoric Acid
Manganous Sulphate in syrupy (concentrated) Phosphoric acid solution reacts with Chlorates to form the violet coloured Mangani-Phosphate ion:
6Mn++ + 12PO4- - - + 6H+ + ClO3 - === 6[Mn(PO4)2]- - - + Cl - + 3H2O
Persulphates, nitrates, bromates, iodates and aslo periodates react similarly.
A drop of the test solution is put into a micro crucible and a drop of the reagent is added. Warm rapidly over a micro burner and allow to cool. A
violet coloration appears. Very pale colorations may be intensified by adding a drop of 1% alcoholic diphenylcarbazide solution when a deep violet
colour, dur to an oxidation product of the diphenylcarbaxide, is obtained.
Sensitivity: 0.05 ug (micro grams). ClO -.
Concentration: 1 in 1,000,000.
The reagent is made up by adding equal volumes of concentrated phosphoric acid and saturated manganese (II) sulphate solution.woelen - 17-5-2007 at 11:41
If you have the chems, then you could use that method. It is based on formation of a manganese (III) complex, which is deep purple.
If you don't have the chems, simply use the method with the SO2 smell. If you use just a small spatula at a time, then believe me, you won't gas
yourself. SO2 is pungent, and you definitely notice the smell. It is H2S, which kills your sense of smell, not SO2.
So, if you think you smell SO2 remains and do not really trust, then simply add one other small spatula of Na2S2O5 if that makes you feel more
comfortable. Simple, effective and only requires a spatula or 2 excess of Na2S2O5.
[Edited on 17-5-07 by woelen]hashashan - 17-5-2007 at 12:07
ok well, probed some of my batch, gassed myself with SO2 . i just don't smell
this thing, just started to hurt in my nose and then in the lungs.
I do have all the chems for the test, but for some reason it doesn't work.
how much of the sulfate should i add to the acid? if i add in in stoichiometry it just gets really really gooey.
anyone?woelen - 17-5-2007 at 13:15
If the SO2 burns in your lungs, then you really don't have to worry about remains of chlorate .
I did the other test with some chlorate, and it works nicely. Simply add a tiny pinch of solid MnSO4 to 2 ml of phosphoric acid (85%). Then I added
some chlorate and heated all of this, until the water was boiling away. The color quickly changes from colorless to a beautiful purple, quite close to
the color of a not too concentrated solution of KMnO4.
The use of a 1 : 1 mix of concentrated H3PO4 and conc. MnSO4 solution does not seem very good to me. The only thing you get in that case is a lot of
black crap. The complex is very dark and even tiny amounts already have a deep purple color.G.i.B. - 17-5-2007 at 15:30
I just tried the MnSO4-H3PO4 (85%) (add solid to syrurpy liquid) method, and it works just fine. Even with a tiny amount of chlorate I got a nice
purple colour. I forgot to heat the first time, ahum... Make sure it boils a little !!hashashan - 17-5-2007 at 15:40
The problem is that my MnSO4 is already a little bit purple(but really a bit).
now you say you take 2ml of phosporic and add there really a small amount of MnSO4(i mean like at the tip of the knife? or some more?) and then you
add couple of drops of your solution and then you boil it down? because i did tget the colorless liquid... can you please be more specific because i
do have some problems with this methos
thanksG.i.B. - 17-5-2007 at 15:54
I think MnSO4 should be a little purple (mine is anyway) Take 2 ml H3PO4 (85%) add a little MnSO4 (I tried a tiny bit up to a good knife tip full, it
all worked) Add 0.5 ml chlorate solution, and then heat to a boil. If this doesnt work for you, I think you should check your chems.
Go to that link, 3rd post down for a demonstration of chlorate detection. This is what
happens when 5 ppm threshold is reached and why I only used 3 drops instead of a
full 5 ml. Even at 1 ppm a decoloration of the mix shows. Granted, the MnSO4 test is
more sensitive but 1 gram of IC can provide 1000 tests. Just keep it out of sunlight
because I found out it does react to UV. As for chlorate destruction, use ferrous sulphate
as mentioned earlier. It's cheap and available at most garden shops.woelen - 18-5-2007 at 11:35
MnSO4 exists in two forms. The monohydrate is very pale pink. The tetrahydrate also is pale pink, but the color is a little bit stronger.
Manganese (II) ions have a very pale pink color. In solution it looks colorless, only the most concentrated solutions have a faint pink color.
Manganese (III) is red/brown in the form of the aqua ion. The phosphato-complex, however, is intensely purple. This purple color differs quite a lot
from the pink color of manganese (II).
@G.i.B.: Leuk dat je hier ook bent! Ik hoop je hier nog veel vaker te zien. Weer eens wat anders dan chemieforum !
[Edited on 18-5-07 by woelen]hashashan - 20-5-2007 at 04:00
Fianlly managed to preform the test. After reducing all Chlorate I got some nice perchlorate.
Now I am seriously curious how accurate that test? I want to make Ammonium perchlorate but i dont want to risk my fingers/hands/legs and any other
body part, the chlorate makes me really worry.woelen - 20-5-2007 at 06:26
Which test are you talking about? The sniff SO2 test or the test with the pyrophosphate of manganese (III)? The latter test is more accurate. If you
first had a nice purple color and after reduction it remains colorless, then you are sure that hardly any chlorate is left. If the test indicates no
color, then just add a small additional pinch of bisulfite, and then you 100% sure have no chlorate left.
But the SO2 test also is quite sure. If the stench of SO2 is strong, then no chlorate is left.hashashan - 20-5-2007 at 06:39
The SO2 is problematic .. i just can hardly feel that smell. I feel some faint smell only after i gass myself with it ... and the seizure in the lungs
after it isnt the most pleasent(probably not very healthy also) thing.
I was talking about the manganese sulfate method.
Destroying Chlorate with Iron Sulphate
dann2 - 26-2-2008 at 12:36
Hello,
I had always presumed that you can destroy Chlorate with Iron Sulphate (on its own). This is not the case.......I think. You need acid (Sulphuric). Is
this true.
Since you must have acid when doing a titration for Chlorate with FeSO4 then it must be true that you need acid if destroying it.
6FeSO4 + NaClO3 + 3H2SO4 ------> 3Fe2(SO4)3 + NaCl + 3H2O
Yes, you will likely need acid to reduce it via ferrous sulfate. I've always thought of it something like dichromate and permanganate, where the real
oxidising species is more prevalent in acidic solution. The same would be true here I would imagine. Many oxidisers are more potent in acidic
conditions.12AX7 - 26-2-2008 at 13:29
The redox reaction requires H+ to proceed, or Fe(OH)3 drops out of solution and, with a lesser concentration of H+, the reaction proceeds very slowly.
Timhashashan - 26-2-2008 at 22:03
The sulfate, sulfite and all the methods leading to SO4 are problematic.
If you had about 20% chlorate in your solution .. you will get loads of sulfate while crystallizing out any perchlorate, there must be a method to
reduce the chlorate and not to get any sulfates or risk gassing your neighborhood with the HCl methoddann2 - 1-3-2008 at 16:28
Quote:
Originally posted by 12AX7
The redox reaction requires H+ to proceed, or Fe(OH)3 drops out of solution and, with a lesser concentration of H+, the reaction proceeds very slowly.
Tim
Hello Tim,
When you say the reaction needs the H+ to proceed, does this mean the Stiochemetric amount is needed or just enough to keep the solution at a low pH.?
The reaction I am talking about is from a Titration (below):
6FeSO4 + NaClO3 + 3H2SO4 ------> 3Fe2(SO4)3 + NaCl + 3H2O
(You then titrate the excess Fe Sulphate with Potassium Permanganate.)
Perhaps with the Titration you must end up with Fe Sulphite otherwise the Fe(OH)3 would interfere and thats why you need a large amount (enough) of
Acid in the Titration?
Dann212AX7 - 1-3-2008 at 18:10
H+ is the driving force for two reasons: one, oxidation of a metal requires H+ to keep it in solution. This is especially true of Fe(OH)3! (Note if
you were oxidizing something like S, it would give H+ and be prone to runaway like many oxidations can be.) Two, ClO3- is resonance stabilized and
rather stable; an excess of H+ is necessary to form HClO3, which is unstable and reacts. HClO3 has a pKa around 1, so pH less than say 3 is needed.
. i just don't smell
this thing, just started to hurt in my nose and then in the lungs.