AJKOER
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Proposed New/Old Faster Method for KClO3 with Better Yield(?)
I was wondering whether one could use the instability of some metal hypochlorites (like Magnesium and Zinc) to more readily form the chlorates and
then, via KCl, create Potassium chlorate.
A similar technique is actually currently employed to make KClO3 starting with Ca(OH)2 and treating with Chlorine, which first creates CaCl2 and
Ca(OCl)2. Then, the Chlorine water, via HClO, actually pushes the disproportionation in a multi-stage reaction under mildly acidic environment (pH
5.5):
Ca(OCl)2 + 2 HClO = Ca(ClO2)2 + HCl
Ca(ClO2)2 + 2 HClO = Ca(ClO3)2 + HCl
(see: "Handbook of Detergents: Production Volume 142", pages 444 to 445, by Uri Zoller and Paul Sosis) and then upon adding KCl to form KClO3:
Ca(ClO3)2 + 2 KCl => CaCl2 + KClO3
Now, of interest for the current thread is the alluded to properties/stability of Zinc hypochlorite. To quote:
"C. Hypochlorite Of Zinc-oxide or Zinc-hypochlorite.—The solution of zinc-oxide in aqueous hypochlorous acid, decomposes spontaneously, if it
contains excess of acid, into chlorine gas, chloride of zinc, and chlorate of zinc-oxide; and even when the zinc-oxide is in excess, the solution
cannot be evaporated without decomposition. When heated, it gives off hypochlorous acid, probably mixed with a small quantity of free oxygen, and
deposits a white, pearly oxychloride, which decomposes spontaneously into chloride of zinc, chlorate of zinc-oxide, oxygen gas, and a small quantity
of chlorine gas. (Balard.) Zinc-vitriol mixed with excess of hypochlorite of lime, gives a precipitate consisting of zinc-oxide and sulphate of lime,
and a liquid which contains no zinc-oxide, but hypochlorite of lime with excess of acid. (Balard.)—1 At. chlorine in the state of aqueous solution
dissolves 1 At. oxide of zinc. The transparent and colourless solution bleaches tincture of indigo strongly, even after a quarter of an hour's
boiling. (Grouvelle, Ann. Chim. Phys. 17, 37.) The solution is resolved by distillation into hypochlorous acid, which passes over in small quantity,
chloride of zinc which remains dissolved, and oxychloride of zinc which separates in the solid state. (Balard.)".
Source: "Hand-book of chemistry", Volume 5, by Leopold Gmelin, Henry Watts, page 32.
LINK:
http://books.google.com/books?pg=PA32&lpg=PA32&dq=hy...
Now, it is my speculation that the new technique via ZnO in either Chlorine water (or upon treating with excess HClO), and then adding KCl, would be a
better method in some regards. I then came across the following reference which completely supports my contention and only predates my supposition by
a mere 200 years. To quote:
"Zinc Chloride and Potassium Chlorate.—An interesting method for the preparation of zinc chloride in connection with potassium chlorate has been
described by K. J. Baver [1]. In carrying out this process chlorine gas is passed into zinc oxide suspended in cold water, whereby zinc chloride and
zinc hypochlorite are first formed, the whole of the zinc going finally into solution. According to the following equation:
2 ZnO + 4 Cl = ZnCl2 + ZnCl202.
1. Chem. Ztg., 1805, XIX, 1453 to 1453; Journ. Soc. Chem. Ind., Jan. 31, 1806, p. 32.
The zinc hypochlorite is converted into zinc chloride by potassium chloride, potassium chlorate being formed in the reaction as follows:
3 ZnCl202 + 2 KCl = 3 ZnCl2 + 2 KCl03
The potassium chloride may be added to the zinc oxide in the beginning or after the solution of the zinc has been effected, the temperature being
maintained at 90 to 95° C. In either case a solution of about 30° B. is obtained which on cooling yields fairly pure crystals of potassium chlorate.
On concentrating the mother liquor to about 60° B. the greater part of the remaining potassium chlorate crystallizes out. The mother liquor, then
containing only very little potassium chlorate, is treated with chlorhydric acid and evaporated to dryness to obtain zinc chloride. The apparatus
required differs but little from that which is used in the manufacture of potassium chlorate by the lime process, the two processes being indeed
identical in their reactions with the only difference that zinc oxide is substituted in one for calcium oxide in the other. The evaporation of the
zinc chloride solution may be carried out in iron vessels, preferably in a vacuum, if traces of iron are not detrimental. The zinc process is claimed
to give a better yield of potassium chlorate than when either lime or magnesia is employed, and of course the whole of the chlorine used is btained as
a marketable product. The zinc chloride which is produced is very pure but may contain a little potassium chloride, which, however, is not detrimental
for many purposes."
Source: "The metallurgy of zinc and cadmium", by Walter Renton Ingalls, pages 686 to 687.
http://books.google.com/books?id=RnhUAAAAYAAJ&pg=PA687&a...
I am not sure of the authors last equation as it is my speculation that the formed Zn(ClO3) just immediately reacts with the KCl to form KClO3. I will
attempt this procedure and invite others to comment/try as well.
Note, HOCl can be formed quickly (with a Sodium Acetate presence) by adding vinegar to bleach (NaClO), or CO2 to aqueous Ca(OCl)2 and filtering out
the CaCO3.
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Pulverulescent
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I haven't read all of your post but in the metathesis of NaClO<sub>3</sub> and KClO<sub>3</sub> the K salt precipitates simply
because it's less soluble than the Na salt!
To produce a particular chlorate this way you must start with chlorate . . .
P
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AJKOER
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Pulverulescent, thanks for your comment.
To be clearer, I now state the creation and disproportionation reaction for Zinc hypochlorite. The textbook creation of a hypochlorite is by the
action of HOCl on an oxide or hydroxide:
ZnO + H2O + 2 HClO --> Zn(ClO)2 + H2O
Then, in the case of Zinc hypochlorite, being highly unstable, there rapidly occurs a multi-stage disproportionation reaction forming the chlorate in
the presence of HOCl in a slightly acidic environment (pH 5.5 for NaOCl):
Zn(OCl)2 + 2 HClO = Zn(ClO2)2 + HCl
Zn(ClO2)2 + 2 HClO = Zn(ClO3)2 + HCl
The common industrial process for Chlorate production does not directly employ HClO, but instead passes Chlorine into solution (forming HClO):
Cl2 + H2O <---> HClO + HCl
A slower reaction in the presence of a base (pH 11) and heat (around 80 C) is given, in the case of NaClO (a more stable hypochlorite) by:
NaClO + NaClO --> NaClO2 + NaCl
NaClO2 + NaClO --> NaClO3 + NaCl
where the formation of the chlorite is said to be a limiting step.
Now, I agree that the classical lab preparation techniques for many chlorates usually start with a chlorate, but the goal here is the more
efficient/faster formation of KClO3 from the basic starting elements.
Here is another reference:
http://www.scribd.com/doc/30121142/Dichlorine-Monoxide-Hypoc...
[Edited on 6-1-2012 by AJKOER]
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weiming1998
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What about this method?
174 grams of potassium sulfate is added to 142 grams of calcium hypochlorite is added to a beaker. about 500mls of water is added to the mixture.
filter out the calcium sulfate (the solution attacks filtering paper and even cloth! You will then get 180 grams of potassium hypochlorite in
solution(assuming you lost none during the filtering process) The solution is boiled until water have been drive out, then water (about two-three
times the thickness of the layer of white powder on the bottom of the beaker is added. The mixture is chilled in a freezer, then taken out, and
filtered again. The powder on the filtering paper is dissolved in another beaker of water, and then is boiled to near-dryness, then put in the sun to
dry completely. I tried this and the yield is pretty good. I've got a thread on this as well. http://www.sciencemadness.org/talk/viewthread.php?tid=18548
There is no reference on it. It simply involve a double displacement reaction http://en.wikipedia.org/wiki/Salt_metathesis_reaction
Then extracting the chlorate because of the solubility differences.
[Edited on 16-1-2012 by weiming1998]
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entropy51
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Quote: Originally posted by S.C. Wack  | 20 g. of pool hypochlorite (label says: 52% Ca(OCl)2, 49% active chlorine) was stirred with 200 g. warm tap water until everything that was going to
dissolve dissolved. This mixture was vacuum filtered through fritted glass, leaving solids that weighed 5 g. after air drying with heat. The filtrate
was placed in a 250 ml. glazed and covered porcelain crucible and heated in a water bath to 70C for 4 hours, a random time and random temperature. It
was then removed from the bath and 20 g. KCl (a likely ill-advised very large excess) was added with stirring, and the mixture was again vacuum
filtered and placed in the crucible, which was then placed in the freezer. After 4 hours there was another vacuum filtration (precooled fritted
filter), and the precipitate was recrystallized by dissolving in 25 ml boiling tap water and cooling in the fridge.
This precipitate was, guess what, vacuum filtered, air dried with heat, and weighed (2.7 g. of familiar uniform, dry, and free-flowing glittering
plates, isolated yield 45% of theoretical). After powdering, 2.39 g. of this was heated to full evolution of oxygen in an weighed test tube with a
certain amount of ignited MnO2 (homemade CMD). The loss in weight (0.93 g.) corresponds to 100% purity.
Covering the solution during heating seems important, so that it doesn't react with CO2. There did not seem to be any odor of chlorine during the
heating, but then I have a cold.
Theoretical considerations:
20 g. x 52% = 10.4 g Ca(OCl)2 (mw 142.98)
= .0727 mole Ca(OCl)2
3 Ca(OCl)2 -> 2 CaCl2 + Ca(ClO3)2
= .0242 mole Ca(ClO3)2
2 KCl (mw 74.55) + Ca(ClO3)2 -> 2 KClO3 (mw 122.55) + CaCl2
= .0485 mole KClO3
= 5.943675 g. |
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kmno4
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The best methode is visiting any good pyro shop.
1 kg ~ 4 american dollars, 25 kg bag ~ 70 american dollars.
Yield is usually better than 99%.
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weiming1998
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| Quote: Originally posted by kmno4 |
The best methode is visiting any good pyro shop.
1 kg ~ 4 american dollars, 25 kg bag ~ 70 american dollars.
Yield is usually better than 99%. |
Sorry, but in some countries like Australia, the best oxidizer you are ever going to get from the shops is calcium hypochlorite. The rest you'll have
to make yourself.
You can get potassium permanganate in Perth's pharmacies, but they are very very expensive. You get two 50 gram bottles for 20 dollars, which is the
same price for 2 kilos of calcium hypochlorite. Other oxidizers... no, don't even think about it, unless you have the time and equipment to
fractionally crystallize fertilizers, or wash off the waste in hundreds of matchheads.
[Edited on 16-1-2012 by weiming1998]
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Pulverulescent
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| Quote: | | 25 kg bag ~ 70 american dollars. |
What? You can just walk in and walk out ─ with 25kg of potassium chlorate??? 
Where? ( )
P
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weiming1998
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There is no pyro shop around me, and possibly in the entire city of Perth. I just went to Sigma Chemicals today and bought 2kg oxalic acid and 1 kg
sodium thiosulfate. The workers there keeps on asking me questions on what it is going to be used for. Imagine the paranoia from them when I order
potassium chlorate, an obvious illegal fireworks/low explosives precursor! So buying is not an option, even from chemical suppliers.
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kmno4
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I am just against making some compounds on largher (>10 g ) scale, because of their availability (via internet shops).
This includes KNO3, NaNO3, KMnO4, KClO3, KClO4, Na benzoate, acetate...... produced in China and exported to other countries for pyrotechnic purposes
- in very nice prices and quality. I realize that in some countries they are not easily available because of "paranoia".
What can I say - pity.
ps. Sigma is the last place where I would buy anything.
ps2. additional info only for "trusted" members - arbitrary trusted in my opinion. Mentioned prices after recalculation from internet shop in
not-very-eastern-europe.
[Edited on 16-1-2012 by kmno4]
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Pulverulescent
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| Quote: | | Sigma is the last place where I would buy anything. |
And Sigma is the last company that would actually sell you anything! 
But you haven't answered my question upthread; where exactly is this "silly, stupid country" that sells KClO<sub>3</sub>, OTC, in
quantity?
P
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weiming1998
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But I looked around the forums once and somebody said that Sigma Chemicals is good for individuals to get chemicals from. They were wrong, and I would
had better luck buying chemicals if I went to another pool shop around my house. You can even get concentrated HCl there without workers asking you
anything. Anyway, this is getting off topic. Let's change the subject back to the preparation of KClO3.
[Edited on 16-1-2012 by weiming1998]
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AJKOER
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OK, yes there are good methods that involve the disproportionation of KOCl, NaOCl and Ca(OCl)2, and many thanks for the descriptions.
One of the original point of the thread, however, is that these hypochlorites are more stable than either Zinc or Magnesium hypochlorite. And as
such, perhaps an unstable hypochlorites is a better candidates for forming a chlorate salt of choice for the home chemist.
In addition, the Zn and Mg chlorate formed are sufficiently stable as to permit the double replacement into another chlorate. As an example of an
unstable candidate, Silver hypochlorite otherwise decomposes into AgCl and Silver chlorate, being only stable in the presence of Ag2O. However, in
addition to being an expensive salt, Silver chlorate itself breaks down with time into AgCl and O2.
Finally, please avoid storing an unstable hypochlorite as there is a history of self-sustaining fires (in one case having occurred in the marine
transport of NaOCl with an apparent Magnesium impurity).
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Poppy
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Seems the whole process need approval, as wel as experimentation on manufacturing such large ammounts without melting plastic containers ...
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woelen
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That story about NaOCl with magnesium impurity must be nonsense. Probably there is some mistake in this.
NaOCl only exists in aqueous alkaline solution at room temperature and then the best you can get is around 15% active chlorine. If you go above this
concentration then there will be decomposition and the more you go above this concentration the faster the decomposition. Solid NaOCl decomposes
explosively at room temperature and only at temperatures many tens of degrees below zero this solid can be kept around for any length of time.
There are very few hypochlorites which are sufficiently stable to have them around at room temperature. The only two about which I am sure are
Ca(OCl)2 and LiOCl. The first usually is present as a dihydrate, the latter usually is mixed with LiOH and LiCl to make it more stable on storage.
Probably you are mistaken with NaClO2. The latter also is unstable, but a mix of NaClO2 with appr. 15% by weight of NaCl in it is stable and can be
kept around indefinately at room temperature. Pure NaClO2 initially slowly disproportionates to NaCl and NaClO3 but this process stops when the
concentration of NaClO2 has gone down to appr. 80%.
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AJKOER
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Poppy:
The Zinc process as described and presented in the referenced book (http://books.google.com/books?id=RnhUAAAAYAAJ&pg=PA687&a... ) was a commercial process. Note, the caption at the top of page 687, "Manufacture
of Zinc Dust, Zinc White, etc."
Also, to quote from above:
"The apparatus required differs but little from that which is used in the manufacture of potassium chlorate by the lime process, the two processes
being indeed identical in their reactions with the only difference that zinc oxide is substituted in one for calcium oxide in the other. The
evaporation of the zinc chloride solution may be carried out in iron vessels, preferably in a vacuum, if traces of iron are not detrimental. The zinc
process is claimed to give a better yield of potassium chlorate than when either lime or magnesia is employed, and of course the whole of the chlorine
used is btained as a marketable product. The zinc chloride which is produced is very pure but may contain a little potassium chloride, which, however,
is not detrimental for many purposes."
Source: "The metallurgy of zinc and cadmium", by Walter Renton Ingalls, pages 686 to 687.
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AJKOER
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OK, here is a reference on the alleged cargo ship disaster relating to a Japanese supplier and commercial hypochlorites with a Magnesium oxide
impurity in the lime used to produce the hypochlorite (with the lime reference this is Ca(OCl)2 and not, sorry, NaOCl). Discussed in: "Bretherick's
handbook of reactive chemical hazards" edited by P. G. Urben, page 1358.
Also a free online Google book.
Link:
http://books.google.com/books?id=_dW_2XPbo_oC&pg=PA1358&...
The suspected cause remains minute amounts of Magnesium hypochlorite formed along with Ca(OCl)2. These dried salts in the presence of moisture is much
more dangerous, I suspect, than any aqueous Mg(ClO)2 solution (even though both have the potential to accumulate an explosive Chlorine oxide, but in
the dried mixed salt case, there is both an elevated concentration and propagation risk). In addition, per the same reference, page 1473, there
appears to be an interaction effect as Magnesium hypochlorite is linked to the explosive and spontaneous decomposition of Calcium hypochlorite (as
produced from magnesia-containing lime derived from dolomite).
I believe the depicted account as it was probably a major insurance loss, hence the reference to an investigation as an insurance company, after it
pays the claim, can sue the responsible company's insurance carrier, to recoup its payout (so called subrogation).
[Edited on 7-3-2012 by AJKOER]
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Vikascoder
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I prefer making potassium chlorate via electrolysis of saturated potassium chloride by graphite electrodes yield is about 30 grams a day which is
good for me. No need of again and again filtering , freezing . This is the best way
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