I was wondering a simple question how would I figure out weight and or approximate volume measurements of kcl to add to a packet of 50% calcium
hypochlorite? When stated on package 50% hypochlorite rest inert chemicals to make a complete reaction?Formatik - 27-6-2008 at 19:15
100% / 50% = 2 that's the mass conversion factor.
Get the equation: 3 Ca(OCl)2 + 2 KCl -> 2 KClO3 + 3 CaCl2 (hopefully that's the reaction!)
Get the molar mass of the reactants. Ca(OCl)2: 142.98 g/mol. KCl: 74.55 g/mol.
Know the moles of the reactants, in this case, 3 and 2.
The mass ratio of reactants is 2*74.55 / 3*142.98 then *2 = 0.174 g KCl to 1 g 50% Ca-hypochlorite.
Now how well that reaction will work, especially with the 50% "inert ingredients", I'm not sure.lucky123 - 29-6-2008 at 16:04
Maybe I confused here?? Wouldn't the inert chemicals weight depending on what they are throw off the 1 gram weight of hypochlorite... I mean what if
the inert chemicals were lead dust? That 1 gram would be mostly lead in weight and not even half a gream hypochlorate? Or am I getting something
wrong? Is the packetlabeled 50% by weight or volume hypochlorite? And is there a way to calculate the volume measurements perhaps if you don't have
a scale? Perhaps I'm complicating something here... Any help with this???lucky123 - 29-6-2008 at 16:17
I think i see my confusion the packet is by weight not volume percentage...duh... So if I got this solids labels will always be labled weight and
when talking about like liquids it will be by volume correct?Formatik - 29-6-2008 at 16:20
What they mean with 50% is most likley by mass, so that means 1 g of the entire material has 0.5 g calcium hypochlorite.
The inert ingredients could complicate it, and who knows what it is. But there's no way there is lead dust in there.
When talking about liquids it should be in mass too, unless it is mentioned otherwise.
Ca(OCl)2
MadHatter - 30-6-2008 at 14:03
The inert ingredient is usually CaCl2 so nothing to worry about there. Unless I can see
a reference, I just don't see how hypochlorite will convert to chlorate without some thermal
decomposition of the hypochlorite as an intermediate. Dissolve and boil it in water for
about 20 minutes.
3 Ca(OCl)2 --> Ca(ClO3)2 + 2CaCl2
Calcium Chlorate is far more soluble than the potassium cation so crystallizing it
out is fairly easy.
I think I mentioned this before somewhere, but Ca(OCl)2 is not disproportionable to Ca-chlorate (as with KOCl)... you'd indeed first have to do the
metathesis reaction from Ca-hypochlorite to KOCl, which is then heated, to precipitate the KClO3.
Also, I think what's important in the Ca hypochlorite is the percentage of active chlorine/oxygen! Not all hypochlorite is present as Ca(OCl)2 but
also as Ca(OH)OCl, and other things. Better use an excess of hypochlorite to start off with, or titrate the Ca-hypochlorite to determine the
percentage of active chlorine/oxygen.DerAlte - 30-6-2008 at 18:48
Use the sodium hypochlorite in bleach or HTH pool. These contain about the same weight of NaCl as NaOCl plus a very small (insignificant) amount to
NaOH to keep pH somewhat above 11 for stability.
Boiling this for enough time gives hypochlorite by disproportionationat pretty close to 100% yield:
3ClO- --> 2Cl- + ClO3- . You land up with a lot of salt and on precipitating out with KCl get even more (do the math).
Most calcium products contain Ca(OH)2 which is poorly soluble, as well as CaCl2 IMHO (and experience). Or maybe it's the Ca(OH)OCl chemleo mentions
above. If you heat calcium 'hypochlorite' dry it does produce oxygen, and maybe when wet, too.
Regards
Der AlteScienceSquirrel - 1-7-2008 at 04:08
Making potassium chlorate from bleach requires heating a lot of bleach.
A rough back of an envelope calculation will show you that boiling a litre of 10% w/v bleach and then concentrating it to ca 100ml and adding
potassium chloride will yield ca 50g of potassium chlorate.
Boiling the bleach should be done in glass or Teflon lined pans using plastic stirrers and preferably outside as you may produce quite a lot of fumes
in the early stages.
There is a useful solubility table here which should help in the final recrystallisations; http://en.wikipedia.org/wiki/Solubility_tableS.C. Wack - 6-7-2008 at 20:08
Quote:
Originally posted by chemoleo
I think I mentioned this before somewhere, but Ca(OCl)2 is not disproportionable to Ca-chlorate... you'd indeed first have to do the metathesis
reaction from Ca-hypochlorite to KOCl
Not being at all pissed that a moderator said that heating solutions of calcium hypochlorite does not produce chlorate, two weeks after I said that it does, the following happened:
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
I think that you will find that on heating dry calcium hypochlorite decomposes to yield oxygen and a solution of calcium hypochlorite decomposes to
yield the chlorate.
Calcium chlorate has been used as a weedkiller.blogfast25 - 8-7-2008 at 04:30
Very interesting indeed. Never stopped to realise hypochlorite provides a route to chlorate. I use K chlorate (but any chlorate will do) as a heat
booster (with Al powder) for thermally challenged thermite reactions (Ti, Si and others) but am not sure my regular supplier will want to play in the
near future.
S.C. Wack: do by vacuum filtrations you mean Buchner filtrations? Do you use vacuum as a fast and convenient alternative to gravity
assisted filtration or is there, in the context of this thread, a more specific reason to resort to vacuum?
As a means to generate significant quantities of chlorate, this may not be enormously convenient but it appears a lot easier to develop than the
electrolytic oxidation of chlorides...ScienceSquirrel - 8-7-2008 at 05:14
The easiest approach to potassium chlorate is from sodium chlorate weedkiller if you can obtain it.
This is a blend of sodium chlorate with a fire depressant in most parts of the world, France is an exception they can buy it neat.
Mixing a hot solution of this with a hot solution of potassium chloride and cooling it will precipitate potassium chlorate in very high yield and
purity.
Vacuum filtration is an advantage as filtration is very fast. You do not want the solution to warm up as your potassium chlorate will start to
redissolve cutting your yield.
Using glass frits as opposed to filter paper is an advantage as it avoids the problem of filter papers soaked in chlorate solution. Wet they are
harmless but you really do not want to let them dry out.blogfast25 - 8-7-2008 at 09:48
Quote:
Originally posted by ScienceSquirrel
The easiest approach to potassium chlorate is from sodium chlorate weedkiller if you can obtain it.
Well, well, live and learn, I guess. Didn't know I'd probably be able to get a cheap(ish) source of chlorate from the nearest garden centre. Googling
shows chlorate based weedkillers to be easily available in the UK.
And it'd be interesting to see if chlorate based weedkiller, as such w/o refining to KClO<sub>3</sub>, could work as a heat booster (with
Al) in thermites, just for a laugh...ScienceSquirrel - 8-7-2008 at 14:58
Nothing to do with chlorates is a laugh.
Chlorate salts contaminated with transition metal salts or mixtures of chlorates with sulphur etc can be dangerously explosive.
Recrystallisation of chlorates requires good laboratory technique and a scrupulous approach.S.C. Wack - 8-7-2008 at 15:07
The filter was a Pyrex 36060 10-15 micron. These and similar Kimax filters get a lot of use. It is silly to use filter paper if you can avoid it.
After drying the whole filter, solids come off pretty easy.
From Wouter's Website
MadHatter - 8-7-2008 at 15:32
Procedure when using calcium hypochlorite
warning: On one occasion an small explosion occured when I was doing this preparation. I am not sure exactly what caused the explosion. It seems to
have been a steam explosion. I was also not sure wheter I was using calcium hypochlorite or trichlorohydrocyanuric acid, another common pool
chlorinating agent. It seems to be very uncommon that explosions happen and they can probably be prevented by vigorous stirring, but I thought
everyone attempting this method should know so proper precautions can be taken. The procedure below has been optimised to reduce the chances of an
explosion happening.
1. Place 250 ml of water in a heat resistant glass or stainless steel container, large enough to hold twice that volume.
2. Bring the water to a boil.
3. To the boiling water, add 125 gram of calcium hypochlorite in 10 gram portions. The calcium hypochlorite usually comes in tablets, which need to be
crushed first. Stir vigorously during this step, occasionally scraping over the bottom to prevent the formation of a cake of calcium chloride. The
solution will foam a lot. If too much foam is developed, do not add any more calcium hypochlorite and boil untill the foam subsides. Then continue
adding calcium hypochlorite.
4. When all calcium hypochlorite has been added, continue boiling untill no more foaming is observed. Stir continuously.
5. Allow the solution to cool down, and filter to remove the precipitated calcium chloride.
6. In a separate container, dissolve 68 grams of potassium chloride in the smallest volume of water possible (approximately 195 ml). This can be done
by dissolving the potassium chloride in about 200 ml of water, and allowing it to cool. If crystals form, add some more water, boil again to dissolve
the potassium chloride, and allow to cool again. If crystals form, repeat. If not, the solution is ready to use.
7. Mix this solution with the boiled calcium hypochlorite solution. A white precipitate of potassium chlorate should form.
8. Bring the solution to a boil and add water untill all potassium chlorate has dissolved.
9. Allow the solution to cool slowly. Crystals of potassium chlorate will form. Cool to 0 deg C.
10. Filter to obtain the raw potassium chlorate. Rinse the crystals in the filter with ice-cold water.
Just in case someone wanted to view Wouter's method.blogfast25 - 9-7-2008 at 10:04
Interesting. Wouter's website being...? (url, please)dann2 - 9-7-2008 at 15:31
[Edited on 9-7-2008 by dann2]S.C. Wack - 31-8-2008 at 01:08
I found an interesting (and therefore in German) article at Gallica on the subject of hypochlorite decomposition, containing a large number of
experiments. If you want a better version, you'll have to get someone with Wiley archive access to fetch it for you. Or you'll need a library that
hasn't closed its journal access for more than a year, and which will continue to do so for at least another year.
But anyways, it says that if there isn't free alkali present, you don't even need heat, it's so unstable. Heating with free alkali is not only more
work, the amount of chloride produced instead of chlorate is higher. They used homemade hypochlorite, not some commercial product containing traces of
metal oxides, etc.
Using an amount of chlorine to combine with all of the alkali, the resulting solution easily decomposes to chlorate in the expected manner, compared
to hypochlorite containing free alkali. They point out that this was pointed out many years before by Gay-Lussac. Addition of acids also works.
Sorry to dredge up an old thread, but I'd like to report my results with the Ca(OCl)2 route:
Into a 1 liter beaker containing 500mL of distilled water, 100.0g of 54.6% calcium hypochlorite pool shock was slowly added with stirring. The
hypochlorite foamed somewhat as it was added. The beaker was covered with the watch glass. Heating was carried out in a microwave oven on maximum
power. The mixture was watched at all times as it is highly prone to “boiling” (probably from oxygen release) over, especially at the beginning.
The solution was brought to a hard boil and was kept boiling for 20 minutes. During this time the solution reduced to about half of its initial volume
and the pale green-yellow of dissolved hypochlorite ion vanished.
The solution was vacuum filtered after cooling. Around 235mL of clear filtrate was collected. The solids were relatively dry and were discarded. The
filtrate was transferred to a 600mL beaker and 19.0g of potassium chloride was added. This was covered with the watch glass and boiled until the
solution volume was about 130mL. During the boiling, the solution became cloudy, likely due to the reaction of dissolved Ca(OH)2 or traces of
remaining hypochlorite with CO2. The solution was gravity filtered while hot through a loose cotton plug to remove the cloudiness and placed in a
salted ice bath. Potassium chlorate separated as flat, colorless, glittering plates.
The crystals vacuum filtered and washed with two 25mL portions of ice cold distilled water. Air was drawn over them for several minutes to dry them as
much as possible. They were dried completely by leaving at room temperature for about 24 hours (the solid is not at all hygroscopic and readily dries
without the need for a dessicator) to afford 15.2g of odorless, free-flowing, glittering plates. This was 48.7% of the theoretical yield based on
calcium hypochlorite.
I would tend to believe that basic calcium hypochlorite (which is still included as part of the weight percent of hypochlorite) decomposes to oxygen
and basic calcium chloride when heated, which means that there is an upper limit for yield. I would suspect that it is around 65-70%. One of the nice
things about Ca(OCl)2 is that there is no glass etching, which I noticed when boiling down bleach.
[Edited on 1-18-09 by UnintentionalChaos]Ebao-lu - 19-1-2009 at 23:17
yes, the pH of solution is very essential. probably the reaction HOCl + 2OCl- = HOCl3 + 2Cl- is much more preferable then 3ClO- = ClO3- + 2Cl- due to
both reacting particles have negative charge..
i was trying to boil bleach for hours, but nothing happened (i am morteover sure, there was no decomposion of hypochlorite) because i did not acidify
it to proper pH. there is a plenty of free alkaline in a bleach.
[Edited on 19-1-2009 by Ebao-lu]Formatik - 26-3-2009 at 16:47
I can only agree with the comments about pH. Hypochlorites will also decompose into the chloride and oxygen lost in the form of O2, the more alkaline
it is (e.g. Gay-Lussac, Lieb. Ann. 43 [1842]178). Commercial ClO- is all alkaline. That's all underlying of the point of forming chlorate from a base
like KOH and Cl2, once the neutralization point is crossed.JohnWW - 26-3-2009 at 17:44
Commercial "chloride of lime", sold mainly for swimming pools, an off-white powder in appearance, is made by chlorination of technical-grade slaked
lime, Ca(OH)2, made from crushed and calcined and then slaked limestone, with Cl2 gas. Byproducts are HCl vapor which is collected either by
condensation or passing the vapor into water, and CaCl2 which remains mixed with the hypochlorite. In view of Formatik's first post on this thread,
and subsequent posts, it appears likely that, when used, some of the admixed CaCl2 may react with the hypochlorite to produce chlorate in the first
instance.woelen - 27-3-2009 at 04:59
I do not agree with JohnWW. Modern calciumhypochlorite does not contain half of its amount as calciumchloride. The "chloride of lime" as described by
JohnWW has formula Ca(ClO)Cl and only has 43.5% of active chlorine if pure (which means that for each 100 grams of pure solid 43.5 grams of elemental
chlorine can be formed if excess HCl is added). Modern calciumhypochlorite is fairly pure Ca(ClO)2.2H2O. Theoretically this has an active chlorine
content of 79.2%, but the commercial preparations usually have around 70% (the remaining solid being some CaCl2, mainly Ca(OH)2 and some CaCO3).hunharcos - 26-5-2009 at 03:01
Hi,
I faced some stupid problems, while making chlorate from Ca-bleach (Wouter's method).
I started out with 500 ml water and 250g hypochlorite (378g with all impurities).
When there were just a few gramms left to add, I came to realize that my thermometer had broken . I know that I don't really need it in this reaction, but I like to feel having control on the reaction.
So, my questions are:
-Is the Hg reacts with anything important - paying no regard to the impurities?
-If yes, is the product soluble in water? (I know HgCl2s solubility at all temp.)
-How does the product affects the chlorate or the following procedures?
-The solution turned slightly pink after heating - even without Hg (first picture/SM1). Any guess why?
Now the Hg or Hg-compound is on the bottom (black coloured) and when boiling, some of those (veeery little) particles are flows up and sinks back
again (second picture/SM2).
If you guys say that there will be something trouble, i will discard it- regardless to the nature
All answers, suggestions, hints-tips and tactics etc... are appreciated!
Thanks
Formatik - 26-5-2009 at 19:19
The pink is the hypochlorite in aq. form (the same way NaClO takes on a yellow color). Hg from a thermometer will be oxidized. Some literature on
this: Hg when shaken with chloride of lime solution will form a grey, chloride-free powder and HgCl2 in solution (Souberain, Ann. Chim. Phys. [2] 48
[1831] 125; Lieb. Ann. 1 [1832] 267). Ca(ClO)2 solution in contact with Hg changes it to red oxychloride (Ballard, Ann. Chim. Phys. [2] 57 [1834]
301). Shaking Hg with hypochlorite solution forms CaCl2 and yellow HgO which eventually turns reddish (Wolters, J. pr. Ch. 115 [1873] 469; Dingl. J.
215 [1873]). That's the kind of impurities you're looking at.
That container is too narrow to be carrying out that reaction, try using a big glass pot since a lot of foaming occurs initially. Ca(ClO3)2 is very
impure with CaCl2, though you can effect some separation in the aq. phase. Separation can be done with acetone (solid phase separation, CaCl2 is
insoluble). Though alkaline ClO- you can boil till kingdom come and still have significant amounts of hypochlorite. Thermal decomposition of
commercial ClO- is a really a piss poor way to obtain chlorates. The only plus is the convenience of not generating toxic Cl2 gas. phys_geek - 2-9-2020 at 05:50
gr8 thread everybody.
Is this method good to make sodium chlorate also or the high solubility of sodium chlorate makes separation(and synthesis as result) impossible?metalresearcher - 2-9-2020 at 09:04
. I know that I don't really need it in this reaction, but I like to feel having control on the reaction.
