Question about the Purity of KClO3

[Cesium]

I recently magically made some potassium chlorate via a secret process that I cannot reveal.

The stuff burns with a nice lilac flame, but it tests positive for chloride ions when silver/copper nitrate is added (a white precipitate forms). The flame color, however, indicates that they are very few sodium ions in the product. I guess there could be KCl but doesn't make too much sense due to its higher solubility. There were some other impurities listed in the salt substitute (calcium or magnesium silicate and something else that I can't remember) so perhaps this is what is precipitating. Or else there are just so few Na+ ions in there that there are overpowered by the K+? This is doubtful because flame tests are very sensitive to sodium ions...odd. What do you think?

 

[mrjeffy321]

Interesting

The Ag+ solution test seems to reveal the presence of the Chloride ion in solution and the two obvious choices for the source of this would be from Sodium or Potassium Chloride impurities. The solubility of KCl and NaCl in cold water would imply that NaCl would more readily come out of solution and remain as an impurity on the KClO3 crystals.

Obviously there are K+ ions in solution (either from potential KCl or at least from KClO3), but what we don’t know is if there are any Na+ ions. A flame test would ordinarily be very sensitive to the presence of any Sodium ions, but you do not see a yellow flame indicative of Sodium.
So maybe you need another type of test to detect the presence of Sodium ions. But I don’t know of any tests offhand which can differentiate between Na+ and K+ compounds that easily, besides of course a flame test.

My guess is that the Chloride is coming from KCl due to the flame test evidence.
Perhaps a very small amount of KClO3 decomposed when you heated the KClO3 to dry it, thus leaving KCl? But I don’t think you got it up to a temperature where this decomposition occurs very readily.

 

[Cesium]

Perhaps a very small amount of KClO3 decomposed when you heated the KClO3 to dry it, thus leaving KCl? But I don’t think you got it up to a temperature where this decomposition occurs very readily.

Good idea; didn't think of that. I heated it at 400F (204C) in a toaster oven. Wikipedia says it doesn't decompose to around 400C, but maybe this occurred very slightly and this is what the silver nitrate is picking up.

If you wouldn't mind mrjeffy321, could you look at a picture of it burning? I have never really seen super-pure chlorate burn before so I don't know precisely what color I should expect. Can you tell if there are any sodium ions by looking at this flame color? Is that reddish-purple color at the bottom of the flame normal? (sorry for the poor picture quality...it's a video frame).

http://img472.imageshack.us/my.php?image=burninggi7.jpg

 

[mrjeffy321]

I think 400 °F was a little excessive if all you wanted to do was dry the crystals.

Although it is difficult to tell much from your picture, around the edge of the flame (where it is not so bright) it looks like the familiar color I have seen with my clean KClO3.
Actually, I have seen two distinctive flame test colors which I attribute to Potassium compounds. I have done flame tests using very pure (>99%) Potassium Chloride and I see a lilac/violet color in the flame (I have pictures), as one is suppose to see due to the K+ ion. But when I use (which I deem to be clean) Potassium Chlorate and burn it with sugar I see a much brighter flame (obviously since it is actually burning) that has a color which is shifted more toward the reddish purple end of the spectrum. I was not totally sure what was causing this color difference…I figured it was probably either due to the intense brightness of the flame which impaired my ability to see the color, or that I didn’t have all the Na impurities out.

 

[Gokul43201]

If the source of KCl was Losalt (or equivalent), remember that that has about 30% NaCl.

 

[Cesium]

No it was not LoSalt but that is irrelevant because there is NaCl from the original table salt used. (I can't get too specific in the procedure I used because my original post and the epic thread it was in was sadly deleted for this very reason. I therefore have to assume that you are "one who is skilled in the art"). The KClO3 was produced via electrolysis and in this process not all NaCl goes to NaClO3 because at 10% w/w chloride concentration perchlorate begins to form so you have to stop electrolysis short.

The whole point of fractional crystallization is to get rid of the NaCl which is more soluble than the KClO3. After the first crude crystallization, I ended up with an impure product that burned with a yellow sodium-contaminated flame. A second crystallization gave the purer product with the lilac flame.

 

[mrjeffy321]

This picture is of a plain ethanol flame,
As you can see it is practically colorless.

This picture here is of a flame test done with very pure Potassium Chloride,
This should be the lilac / violet color due to the K+ ion

This picture shows a rainbow, so to speak, of colors using a variety of flame colorants (Li+, Sr+2, Ca+2, Na+, Cu+2, K+ salts),
This demonstration was done using Potassium Chlorate as an oxidizer and sucrose sugar as a fuel. To this mixture a flame colorant was added in the desired location.
Due to the intense brightness of the flames it is easier to see the colors reflected off the wall behind the reaction.
The color located a little right of the middle is the color I often see when burning ‘clean’ Potassium Chlorate and sucrose. You can see that this is not really the same color as the above flame test picture of KCl, but it is what your picture reminds me of.

 

[chemisttree]

Was the pan you used to dry the chlorate also used in your manganese dioxide experiments? Was the pan made of aluminum? These two metals catalytically decompose chlorate to oxygen and KCl at temperatures significantly lower than their melting points.

Any of the transition metals that contain partially-filled d orbitals will act as a catalyst to decompose chlorate as well. I'm not sure what silver (I) or copper (II) will do to it in solution. Did the solution effervesce slightly?

 

[Cesium]

The drying container was a tuna fish "tin" can- that is a very thin coating of tin on steel. I could try the procedure again in glass and see if it makes a difference.