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janger
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[*] posted on 29-9-2005 at 21:53
Pool chlorine and purity


Is it possible to roughly calculate the amount of Ca(OCl)2 in pool chlorine from it's 'available chlorine' figure? IIRC, this figure is a measure of sanitizing power as compared to free chlorine. But I think it's the hypochlorous acid, HOCl, which is produced via reaction with water that does the sanitizing.

Say if the label says '650g/kg available chlorine as calcium hypochlorite, hydrated.'
That means 1kg has the same sanitizing power as 650g free chlorine.
There is 1mol of HOCl produced for every mole of Cl2 via the reaction:

Cl2 + H20 ---> HOCl + HCl

n[HOCl] = n[Cl2] = 650 / 70.91 = 9.17mol

but 1mol of Ca(OCl)2 produces 2mol HOCl:
n[Ca(OCl)2] equiv. = n[Cl2] / 2 = 4.58mol

if the hypo is the dihydrate (MW = 179.01g/mol), then

m[ Ca(OCl)2.2H2O ] = 4.58 * 179.01 = 819.88g
which makes it about 82% pure.

Does this work or am I on the wrong track?


A related matter. When I add some pool chlorine to water and shake until all the granules have dissolved, then let it settle, there is a huge amount of sediment at the bottom. Is it all impurities? Calcium hypochlorite is meant to be readily soluble but I'm wondering if it might take some time or maybe a little heat in high concentrations. Just seems way too much sediment to be all impurites/filler.
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woelen
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[*] posted on 30-9-2005 at 01:06


Here follows an interesting quote from http://www.intox.org about the possible impurities:

Quote:

Calcium hypochlorite can exist as the anhydrous material and as a dihydrate. The strength of calcium hypochlorite products is often expressed in terms of percent available chlorine (% av Cl2). The dihydrate is produced on a large scale and marketed as a 65-70% av Cl2 product. The commercial product is a complex mixture containing variable amounts of related salts, such as sodium chloride, calcium chloride, calcium chlorate, calcium hydroxide, calcium carbonate and other impurities. Calcium hypochlorite is also sold with 75-80% or 50% av Cl2, 35-37% av Cl2 (bleaching powder) and 33-35% (tropical bleach). Historically, calcium hypochlorite sometimes contained impurities such as magnesium hypochlorite, which made the product dangerously reactive.


"Available chlorine" of compound X is defined as the ratio of iodine, produced by adding a certain weight of compound X to excess acidified solution of an iodide, and iodine produced by the same weight of chlorine.

So, in reaction form for dihydrated calcium hypochlorite:

Ca(ClO)2.2H2O + 4I(-) + 4H(+) ---> CaCl2 + 2I2 + 4H2O

For chlorine we have:

Cl2 + 2I(-) ---> I2 + 2Cl(-)


Now let's see how many grams of iodine we get from 1 kilo of the calcium hypochlorite: This is 2835.66 grams.

From 1 kilo of chlorine we get 3579.50 grams of iodine.

So, pure dihydrated calcium hypochlorite has an "available chlorine" content of 2835.66/3579.50*100%, which is around 79.2%

Your stuff has a chlorine content of 65%, so the purity of this is 65/79.2, which equals 82%

So, indeed your reasoning seems to be correct, but the industry standard is through the iodide calculation. The reason for this is that the amount of "available chlorine" only is released when the material is immersed in HCl. Half of the "available chlorine" is needed from an external source in the case of Ca(ClO)2.

That purity seems to be reasonable. I have the same stuff with 65% available chlorine, but from another company, I also have the stuff with 70% available chlorine, which corresponds to 88% purity.

The "funny" thing is that the 65% bottle has a very strong smell of chlorine and emits gas. I store it in a very tightly capped container, but every time I open the bottle there is quite some pressure and a very sharp odour of chlorine. The 70% stuff is completely odourless. Any explanation for this? The MSDS of calcium hypochlorite states that it must be stored in a tightly closed container, but I feel a little uncomfortable with that 65% stuff, because of the buildup of pressure. Is there is chance of rupture, or will there be an equilibrium at a certain point, such that pressure does not rise without limit?

The presence of CaCO3 in the stuff also is somewhat a pity. This makes the chlorine gas, produced with the stuff, impure. CO2 cannot easily be removed from the chlorine gas (e.g. bubbling it through NaOH-solution also removes the Cl2). For many purposes, however, a few percents of CO2 should not do any harm.




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neutrino
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[*] posted on 30-9-2005 at 02:27


I would guess that there's some kind of stabilizer in the more concentrated hypochlorite. Maybe the 70% stuff has Ca(OH)<sub>2</sub> as its imputiry, while the 65% material has CaCO<sub>3</sub>.
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[*] posted on 30-9-2005 at 14:06


Great. Thanks for that. Maybe my calculations work well enough for what I need.

The other way of determining 'available chlorine' which I don't think is used any more is to react the sample with H2O2 and measure the volume of oxygen liberated.

The only solubility figure I have for Ca(OCl)2 gives it as 21.4% at 76ºF (25ºC). Is this correct? I've been playing with solutions less concentrated than that, but the impurities may interfere with the solubility. How can I check if all the hypochlorite, or most of it, is ending up in solution without real chem equipment? Dilute a sample and check with those pool test strips that measure 1-7ppm? Hmmm.

I really need to use this stuff up so thought of creating some chloroform, but wanted to convert it to strong NaOCl first. I think the junk causes some problems with this. Last night I decided to try using NaHCO3 - not a good idea. After a few minutes the bubbling started, making a sort of crackling sound every so often and with the evolution of significant amounts of chlorine gas - enough to give me the 'chlorine wheeze'. I have read that NaOCl solutions at least 40% are possible, so why was chlorine given off? Could the evolved CO2 be driving it out of solution? When I tried the same with NaOH it seemed to work fine, so it might be better to use Na2CO3.
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[*] posted on 30-9-2005 at 15:14


Most dry pool chlorine products contain cyanuric acid as a stabilizer, (also called conditioner), which slows the reaction of sunlight and the oxidizer to keep it active in the water longer. Cyanuric acid is only slightly soluble, and may be a sizable portion of the sediment you are finding. A small amount does dissolve, tho, so it could interfere with your results.

Z
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[*] posted on 30-9-2005 at 16:21


Copy/paste from Inorganic Syntheses:

Procedure. A weighed sample (close to 0.1 g.) is added to an iodine flask containing 40 ml. of a solution prepared by dissolving 2 g. of potassium iodide in dilute (not over 2 N) hydrochloric acid. The resulting solution is titrated immediately with standardized 0.1 N sodium thiosulfate until the yellow color of the iodine just vanishes. Any yellow color which may then develop on standing as a result of oxidation of iodide ion by chlorate ion is ignored. Calculations are made as follows:

% "available chlorine" =
[35.457 × (ml. of Na2S2O3) × (N of Na2S2O3)]/[10 × g. of sample]

% Ca(OCl)2
= 1.008 × % "available chlorine"
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[*] posted on 30-9-2005 at 18:00


Iodometric titrations are the standard for determining the concentration. HOCl (and OCl- to some extent) is the active species (I think ~90%) in aqueous systems.

I've titrated aqueous solutions of Br2 and Cl2 and found that 10.0mL of your sample and .2 grams of KI work with 2N sulfuric acid. As you all know, after titration is finished, the clear solution will gradually redarken as it is exposed to oxygen, reverting to its characteristic iodine colour.
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[*] posted on 30-9-2005 at 20:07


Here's something interesting. The solution that NaHCO3 was added to had a greenish yellow color this morning but still stunk like chlorine was being given off. But a few hours later it has turned pink! Any idea why this would occur?
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[*] posted on 30-9-2005 at 21:31


- Is there any iron present?

Once after boiling down a batch of sodium chlorate electrolysis cell liquor in a steel can, after cooling to 5°F (-10C?) I was left with a pink solution. In fact, I think I took a picture... ah, here we go.
http://webpages.charter.net/dawill/Images/CooledSaltSolution...
(Ya, frost forming on the outside.)
My best guess is ferrate ion due to the intense hypochlorite content. After a few weeks at room temperature, the color disappeared.

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[*] posted on 30-9-2005 at 23:12


Quote:
Originally posted by 12AX7
- Is there any iron present?
Tim


Well that's interesting. It was about the same color as your stuff. This was all done in plastic and glass. But I notice when adding the pool chlorine to water there are some browny specks of something in the sediment. Maybe there is some iron present from the production?

Oh, and I found an old post by Rosco Bodine about converting Ca(OCl)2 to the Na form. There's some good info in it.
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[*] posted on 1-10-2005 at 00:16


I have seen swimming pool calcium hypochlorite turn pink when the aqueous solution is heated for a prolonged period. It reminded me of dilute permanganate or phenolphthalein, but as 12AX7 has pointed out, I suppose it could be dilute ferrate as well. I always found it puzzling because I couldn't imagine any colorful organics in those conditions, and I thought that calcium hypochlorite was supposed to be transition metal free to prevent decomposition. But maybe they're using technical materials made from limestone without too much purification, and traces of iron or manganese remain.



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[*] posted on 2-10-2005 at 19:38


Polverone, I think it was you who made some chloroform from Ca(OCl)2, and dissolving the precipitated hydroxide with HCl. The problem I'm having is the ppt takes way more than stoichiometric amounts of HCl to dissolve, like ten times as much. It must be due to all the other stuff in my pool chlorine. There's no way of purifiying it, is there? Apart from filtering there's nothing much else possible, correct?
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[*] posted on 2-10-2005 at 20:06


I never had difficulty dissolving the precipitate. I do think the best approach is to grind your hypochlorite finely with cold water, then filter through glass wool or some other nonreactive filter. This should leave you with nothing that is water-insoluble to start with, yielding easily acid-soluble precipitate after reaction.



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[*] posted on 2-10-2005 at 23:04


I've been experimenting a little with my Ca(ClO)2 (with the 65% and with the 70% available chlorine stuff). Both samples dissolve very nicely in cold dilute HCl without any problem, leaving a totally clear liquid. I've not been able to obtain any pink liquids as others reported here. With my samples, I can get colors ranging from colorless to pale green, that's all.

ASo, I indeed think that the color is due to impurities, either with iron or with manganese. Hypochlorite certainly is capable of oxidizing iron to ther +6 oxidation state and manganese to the +7 oxidation state. I've done this in an experiment with bleach.




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[*] posted on 3-10-2005 at 16:37


woelen, isn't that a bit of a mistake dissolving it in HCl? It releases chlorine gas.

It still seems strange to me that adding sodium bicarbonate causes release of chlorine gas and an eventual color change.
Anyway, letting the mixture settle gives a nice greeny yellow solution. But the sediment is so fine that it's impossible to decant without clouding up again. Of course the liquid could be pipetted off, but that would be for someone with patience. A sintered glass crucible could probably catch the sediment. Wish I had one.
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[*] posted on 3-10-2005 at 23:43


Quote:

woelen, isn't that a bit of a mistake dissolving it in HCl? It releases chlorine gas.

Of course I know that (in fact that is my preferred may of making Cl2 gas for experiments), but what I wanted to say is that the remaining liquid is totally clear. From others I understood that they were left with some solid residue and with colored (pink) stuff. Well, that is something which I have never observed.




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[*] posted on 4-10-2005 at 00:04


I've never seen the pink color in acidified hypochlorite, but in samples that have been mixed with water and subjected to extended heating.

I have a hard time thinking of any calcium-like compounds that are insoluble in dilute hydrochloric acid, so the dissolving problems are a mystery. Are you able to check the pH of your chloroform reaction products as you add acid, or do remaining traces of oxidant discolor pH paper?

I would guess that NaHCO3 is too acidic to be optimal; it will easily give up CO2, which I believe will react with Ca(OCl)2 with the liberation of hypochlorous acid, which as well known is unstable. Do you get a greenish/yellow color if you prepare an aqueous solution of Ca(OCl)2 and then bubble CO2 through it?




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