woelen - 6-10-2006 at 14:55
I have some HClO4, NaClO4 and Mg(ClO4)2. I do aqueous experiments with these chems, but when I'm done with it, I keep the waste. Every time, when I
have collected a few tens of ml of waste, I work it up with KCl to make KClO4. Perchlorate is too good to throw away, the more because it can be so
easily separated (quite good solubility in boiling water, almost insoluble near the freezing point).
I have worked up waste from HClO4 and NaClO4 and that gives me nice perfectly dry crystals, The powder is absolutely dry and with red P it gives an
exploding mix, even when not confined.
Yesterday evening, I worked up waste from Mg(ClO4)2. It is a white crystalline powder. I rinsed the powder another time with some distilled water, and
then dissolved the powder in an as small as possible amount of boiling water and let the KClO4 crystallize from this. Yesterday evening I had humid
crystals. I put them on a warm place (appr. 45 degrees) and now, 24 hours later, they are _almost_ dry, but not totally dry. These crystals also seem
not to dry further 
. After 24 hours I have a sticky and clumpy mass. I have
crystallized them from water twice, with intermediate rinses between the recrystallizations. The finely powdered mix with red P does not explode, it
only gives a bright flame and a whoosh noise. That also is an indication of humidity I think.
In contrast, the other batch from neutralized HClO4 and NaClO4 I have is only recrystallized once and that really is nice dry, free flowing.
Is this because of remains of Mg(2+) in it? Is that tiny trace still causing some hygroscopic behavior of the material I now have? Or does KClO4 form
a double salt with Mg(ClO4)2, which is much more hygroscopic than KClO4 alone? If that is the case, then it would mean that waste from Mg(ClO4)2
hardly can be used as a nice source for KClO4.
Although I am not a pyro man, I intend to use these small quantities of KClO4 for funny little pyro displays (e.g. for the kids, they like it when you
make nice little flashes and pops).
[Edited on 6-10-06 by woelen]
The_Davster - 6-10-2006 at 15:02
What is the pH of the rinse water of the perchlorate? I think it could be some residual perchloric acid or base, depending on what the magnesium
perchlorate was waste from. And of course you know how strong a drying agent magnesium perchlorate is, so any residual of that would absorb water
from the air.(not sure if magnesium perchlorate is deliquescent though?)
12AX7 - 6-10-2006 at 15:48
Did you use KCl? If so, residual MgCl2 could indeed be hygroscopic.
A slightly basic pH might precipitate the Mg, from which the KClO4 can be recrystallized.
Tim
not_important - 6-10-2006 at 21:36
Given that magnesium and potassium form a number of mixed salts, usually from pretty concentrated solution, it is possible that you actually have a
mixed Mg-K perchlorate.
The idea of using some basic salt of potasium, hydroxide or carbonate, when stating with Mg(ClO4)2, would seem to be worth trying. Possibly using an
excess of KCl might make a difference; as well as working with a bit more dilute solution, getting a first crop of crystals, then concentrating the
liquor for a second crop that goes back into the workup stock.
Another thing that might be tried is a wash with cold, saturated KCl solution.
[Edited on 7-10-2006 by not_important]
woelen - 7-10-2006 at 01:34
I already tried something myself. I added some of the humid KClO4 to distilled water and dissolved it. It does not give a precipitate with NaOH. So,
it does not contain much magnesium, but of course, Mg(OH)2 is slightly soluble, so there might still be some magnesium in it. I also tested the
material with AgNO3. There is a very slight blue/white opalescence, so it only contains trace amounts of chloride from the KCl.
What about precipitating with K2CO3 instead of KCl? Could that be the solution of this problem? I could try that, but the K2CO3 I have is quite impure
(ceramics/pottery supplier, it is off-white). My KCl is much purer, it is so-called photo-grade quality from a photography raw chemicals supplier.
not_important - 8-10-2006 at 01:05
I would think you could do the KCl precipitation method, then add a bit of K2CO3 to the KClO4 with enough water to dislove the KClo$ when hot, heat it
up, filter hot, and cool to recover the KClO4. If you start out using K2CO3 either you need it dilute enough that all the KClO4 stays in solution,
then boil it down after filtering, or you filter off the MgCO3 when recrystallizing the KClO4.
Pottery K2CO3 seems to contain K2CO3, KHCO3, KCl, and sometimes traces of Mg/Ca salts. It's fairly free from d-block metals, as those are likely to
colour or otherwise affect glazes made using the K2CO3. Just because it's pottery grade doesn't mean that it's impure, given that parts per thousand
can shift glazes colours.