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MadHatter
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Indigo Carmine Test
Indigo Carmine Test For Clhorates
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.
From opening of NCIS New Orleans - It goes a BOOM ! BOOM ! BOOM ! MUHAHAHAHAHAHAHA !
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woelen
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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]
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hashashan
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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.
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woelen
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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.
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hashashan
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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.
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dann2
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Destroying Chlorate with Iron Sulphate
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
Titration here:
http://www.geocities.com/CapeCanaveral/Campus/5361/chlorate/...
Any and all comments welcome
Dann2
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Mumbles
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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.
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12AX7
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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
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hashashan
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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 method
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dann2
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| 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?
Dann2
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12AX7
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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.
Tim
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