Sciencemadness Discussion Board

CaCl₂ solubility/precipitation

Keras - 1-8-2024 at 05:26

Folks,

this maybe a weird question, but here it is.

I’m planning to refill my stash of potassium chlorate using calcium hypochlorite as the base reactant.

The reaction will generate CaCl₂ as a byproduct, and I don’t want it to contaminate my chlorate. So I basically need to know what is the solubility of calcium chloride.

Problem: calcium chloride exhibits various degree of hydration, and each one has its own solubility value. So which one applies when calcium chloride is generated in situ? The most hydrated form? The anhydrous one? I didn't do the maths, but maybe they all correspond to the same mass of CaCl₂, the difference being only due to the water molecules attached.

Any idea? Thanks!

B(a)P - 1-8-2024 at 14:03

Doesn't it just depend on how much water is available (and temperature)? If you have excess water it will crystallise as the higher hydrate. So if you have excess water it will be the hexahydrate. However, if you crystallise your product at greater than 30C you will get the tetrahydrate, greater than 45C the dihydrate and so on.

j_sum1 - 1-8-2024 at 17:11

CaCl2 is very soluble. And deliquescent. It is sold around here as "damp rid" to absorb moisture from the atmosphere.
As such, it is difficult to precipitate. It does want to stay in solution.
So... if you precipitate your chlorate, it should not have much CaCl2 contaminant in it. Wash with chilled water as usual.

Keras - 2-8-2024 at 00:24

Hi!
Quote: Originally posted by B(a)P  
Doesn't it just depend on how much water is available (and temperature)? If you have excess water it will crystallise as the higher hydrate. So if you have excess water it will be the hexahydrate.


That makes sense, for sure, but I'd like to be sure of it :p

Quote: Originally posted by j_sum1  
CaCl2 is very soluble. And deliquescent. […] It does want to stay in solution.
So... if you precipitate your chlorate, it should not have much CaCl2 contaminant in it. Wash with chilled water as usual.


Right, that was what I expected, but still I was curious, because the question can apply to any salt able to form multiple hydrates, like magnesium sulphate or sodium sulphate, for example.

yobbo II - 2-8-2024 at 07:52


Calcium Chloride has hydrates of 0,1, 2, 4 and 6.
Calcium Chlorate has hydrates of 0, 2, 4 and 6.
The Perchlorate has hydrates of 0, 4 and 6. No sign of a 2 but it probably has.
Figures for Chlorate are from SDS Vol. 14. Figures for the Chloride are from Seidell. For the Perchlorate, SDS Vol. 61
Three patents that have phase diagrams for Ca Chlorate/Chloride are US 1,887,809 and 1,893,740 and 1,949,204. There are phase diagrams in SDS Vol. 14 too but they are not copied to web .pdf satisfactorily. The green lines are drawn to the Chloride solubility points from Seidell. I believe it takes a long time for Calcium Chloride/Chlorate solutions that are ppt'ing salts to stabilise (reach equilibrium) as they are viscous due to high salt percentages.

100 grams saturated aqueous solution contain 64.0 grams Ca(ClO3)2 at 18C. Density of solution is 1.729. That's 177g Chlorate per 100ml water = 64%. (Seidell).

casolubi.gif - 14kB

[Edited on 2-8-2024 by yobbo II]

Boffis - 2-8-2024 at 08:39

Hi Keras,

There is no need to worry. Potassium Chlorate solubility in a strong solution of calcium chloride is pretty low so practically all of the chlorate crystallise out on addition of potassium bicarbonate and then potassium chloride long before the saturation point of calcium chloride solution is reached. This is my preferred route to potassium chlorate and i have pretty well perfected it. Although it is a little messier than the sodium hypochlorite it gives a product free from sodium. The final purification is by by adding a little potassium carbonate or bicarbonate and heating to decompose the bicarbonate slowly (this gives a more readily filterable calcium carbonate ppt) and cooling. I have never tried to isolate the calcium chlorate first.

I'll post my recipe if you like.

bnull - 2-8-2024 at 10:27

You can precipitate calcium using potassium carbonate and forget the chloride issues. It also removes calcium hydroxide and calcium chloride, which are the major impurities in Ca(ClO)2. Edit: Never mind. You'd be exchanging one problem for another. @Thanks, Boffis.

@Boffis: What brand of HTH do you use? The ones I'm used to buy have a few percent of NaCl.

[Edited on 3-8-2024 by bnull]

Boffis - 2-8-2024 at 11:51

Quote: Originally posted by bnull  
You can precipitate calcium using potassium carbonate and forget the chloride issues. It also removes calcium hydroxide and calcium chloride, which are the major impurities in Ca(ClO)2.

@Boffis: What brand of HTH do you use? The ones I'm used to buy have a few percent of NaCl.


No! Exchanging potassium for calcium simply creates an excess of potassium chloride. The chloride concentration is the same and KCl has a steeper thermal solubility curve and lower overall solubility than CaCl2 so tend to ppt if you over evaporate the solution.

Potassium bicarbonate is added initially to lower the pH to about 7-8 as this increases the rate of decomposition of the hypochlorite into chloride and chlorate. But if you add too much it causes the liberation of chlorine. Potassium carbonate simply replaces calcium chloride with potassium chloride, this doesn't actually change anything apart from increasing the number of recrystallisations required to obtain a pure product.

I use a bleaching powder brand called "Relax" chlorine granules.

Keras - 3-8-2024 at 01:15

Thanks for all the input.

I’m going to lower the pH using a small quantity of sodium bisulphate, maybe. Problem is, I have no pH paper here.

I use some brand called PCH as a source of calcium hypochlorite. This is marketed (as written on the package) as pure calcium hypochlorite. I’m a bit skeptical, because it’s given as 70% free chlorine, whereas pure hypochlorite would have nearly 99% according to Wikipedia. So I guess there are 20% of impurities, but on the other hand the hypochlorite is hydrated (5-10% water contents) so that may explain the difference.

Boffis - 3-8-2024 at 03:00

Actually, pH paper is not very useful and it tends to bleach too quickly but excessive acidification results in chlorine generation. Potassium bicarbonate is a better buffer than sodium bisulphate and doesn't introduce sodium.

As has been discussed before on SM, most standard commercial calcium hypochlorite is between 65 and 75% "available chlorine" and since it is preparaed from industrial slaked lime from calcined limestone it is never pure and during the production process calcium chloride is a byproduct. So Bleaching powder is always a mixture of stuff.

I have typed up my notes from some recent K chlorate preparations, I hope this is helpful:

Preparation of potassium chlorate from bleaching powder.

The experimental details presented below represent the near optimal conditions for the conversion of commercial calcium hypochlorite (bleaching powder) into potassium chlorate. As with other hypochlorites, the calcium salt slowly disproportionates into calcium chloride and calcium chlorate in solution and this process is greatly accelerated by heating. Attempts to isolate calcium chlorate were not successful; while this may be possible, the high solubility of both calcium chloride and chlorate makes this a difficult process to operate. The much simpler route is to precipitate the chlorate as the potassium salt by the addition of a small excess potassium chloride. The pH of a calcium hypochlorite solution is >12 but formation of chlorate is favoured by a pH of 7-8. I have found therefore that the process can be improved by neutralising the excess alkali by slowly adding potassium bicarbonate to the initial calcium hypochlorite solution. The precipitated calcium carbonate does not adversely affect the process and is easily filter off before crystallisation. The advantage of the calcium hypochlorite route to potassium chlorate over the sodium hypochlorite route is that the product is free of sodium.

Experimental

250g of commercial bleaching powder were added to 1 l of cold water in a suitable glass or stainless-steel vessel on a stirrer hotplate. The slurry was stirred for about half an hour to allow the hypochlorite component to dissolve. About 33 to 40g* of solid potassium bicarbonate was then slowly stirred into the suspension in small portions at a time until the mixture starts to foam slightly and smell of chlorine. The pH should now be about 7, if it is too low add a little potassium hydroxide solution to raise the pH to 7-8. Cover to limit evaporation and warm with stirring, the white suspension slowly to about 90° C over half an hour and then maintain this temperature for 7 to 8 hours. The is usually some evaporation and fresh water was added periodically to maintain the initial volume.

After the allotted time the mixture was cooled slightly and 30g of fine potassium carbonate were added slowly. The solution was stirred for 20 minutes and then filtered hot through a large Buchner funnel (12.5 or 15cm). The result was a beautiful clear pink solution. Reheat to 80-90° C again and add 60g of pure potassium chloride** powder and stir until dissolved. Allow to cool to room temperature, say 15° C, overnight and filter of the crystals, wash with a little ice-cold water and air dry in a warm place. The yield of crude KClO3 is usually about 50-52g of slightly rusty coloured crystals***.

The combined filtrate and washings were then evaporated down to about 300-350ml, filtered if necessary and cooled to room temperature and then chilled to 4° in the fridge. The resulting crystals ere filtered off, washed as before and dried. The yield is a usually a further 50g or so of less pure potassium chlorate contaminate with a little calcium chloride that makes the crystals difficult to dry.

The two crops were combined and dissolved in 350ml of hot water. When solution was complete 2g of potassium carbonate, 3g of decolourising charcoal and a 2.5ml scoop of cellite or kieselguhr added. The solution was stirred for 15 min at 90° and then filtered hot through a pre-heated Buchner funnel and cooled. The solution was mechanically stirred until at room temperature to control the grain size and then chilled to 4° C. The product filtered and washed with a little ice-cold water and dried. The yield is usually about 72-75g. A further 12-15g may be recovered by evaporating the filtrate and washings to about 100ml.

*) The amount of potassium bicarbonate required for the neutralisation varies from one purchase of calcium hypochlorite to another, so the first time you use calcium hypochlorite from a new source you need to calibrate the procedure. My current stock of bleaching powder claims to be 65% available chlorine and requires about 35g of KHCO3.
**) Some commercial potassium chloride contains anticaking agents that make the solution cloudy. If your potassium chloride is of this type I recommend making up a saturated solution, treating with charcoal and kieselguhr and filtering. An appropriate amount of this saturate solution is then used.If you are going to use a saturated KCl brine you can be a little less generous with the top-up water during the initial heating stage.
***) The colour is thought to be due to the decomposition of the pink ferrate salt that is thought to cause the pink colour of the initial filtrate. See:
http://www.sciencemadness.org/talk/viewthread.php?tid=159662...

[Edited on 3-8-2024 by Boffis]

bnull - 3-8-2024 at 06:09

@Boffis, @Keras: It seems that all brands contain sodium chloride (16% or less). It comes form the manufacturing process and improves stability. Oh, well. :(

Keras - 3-8-2024 at 12:45

Quote: Originally posted by Boffis  
Cover to limit evaporation and warm with stirring, the white suspension slowly to about 90° C over half an hour and then maintain this temperature for 7 to 8 hours.


What? 7 hours? Isn't that a bit of an overkill?

Boffis - 4-8-2024 at 07:26

Quote: Originally posted by Keras  
Quote: Originally posted by Boffis  
Cover to limit evaporation and warm with stirring, the white suspension slowly to about 90° C over half an hour and then maintain this temperature for 7 to 8 hours.


What? 7 hours? Isn't that a bit of an overkill?


Is It? How do you know?

My criteria for the end of reaction was no further evolution of chlorine on addition of potassium bicarbonate. Its not a particularly sensitive test I admit but in this case over cooked is better than undercooked! If you know better please enlighten us.

Keras - 20-8-2024 at 02:48

Quote: Originally posted by Boffis  
Quote: Originally posted by Keras  

What? 7 hours? Isn't that a bit of an overkill?

Is It? How do you know?
If you know better please enlighten us.

Oh no, I don’t. I was just surprised it took so long. I have downloaded a few papers about hypochlorite decomposition kinetics, so I’ll read them and see what I can make of them.

Working with calcium hypochlorite is not easy without the full paraphernalia. Dissolving commercial calcium hypochlorite releases very thin particles of calcium hydroxide/calcium carbonate which will pass through any sufficiently coarse filter (in my case, a glass wool plug). Lowering the pH will destroy the dust, but this cannot be done with hydrochloric acid, since it’ll also destroy the hypochlorite. I used sodium bisulphate, which precipitated calcium sulphate by double displacement (and I got standard bleach). Calcium sulphate is easier to filter out but one must wait for it to settle, which takes about a day, and then the filtration is very slow, assuming you only can do gravity filtration. I’m not sure vacuum filtration is the way to go, since removing calcium sulphate stuck in a Büchner funnel sounds like a nightmare, given the inertness of the chemical. I think the best way is to use a coarse porcelain funnel with a glass wool filter paper disc.

In any case, I lost quite a lot of product because plaster retained an appreciable amount of water, and probably hypochlorite too, and I had no vacuum filtration equipment at hand.

At that point, you’re left with standard bleach to work with (though somewhat purple, which might be caused by manganese impurities), and that’s pretty straightforward.