FireLion3
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OTC Calcium Hypochlorite is labeled as ~70% - what are the impurities?
I am unable to find any MSDS for any products of this kind. It is largely a pool product and pool suppliers for some reason never label their product
percentages in terms of weight of molar mass so it makes it difficult for the amateur chemist to use these products for chemistry purposes.
Someone suggested that its labeled as 70% chlorine, but is in fact 99-100% calcium hypochlorite. Another person suggested that it is 70% calcium
hypochlorite with some added base to stabilize it. Does anyone have a suggestion? I even called the manufacturers and they were surprisingly not able
to give me any answers.
[Edited on 8-9-2014 by FireLion3]
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Amos
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Does it say "70% calcium hypochlorite" or does it say "70% available chlorine" or what? Please be more specific or post a picture. As for the calcium
hypochlorite for pools(70% is what I've seen in 2 products, and they mean 70% of the hypochlorite), the most common impurities I've seen are calcium
oxide/hydroxide, which you can remove most of through filtration.
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FireLion3
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I am not exactly sure. I haven't bought it yet because I don't want to waste my money. As my memory serves me, I remember that among the products they
seem to claim a variety of those. Some 70% hypochlorite (referring to the ion by weight? Which would be 99% Calcium Hypochlorite, correct?), and some
say 70% calcium hypochlorite, which may or may not be referring to hypochlorite ion, and some say available chlorine.
I can't imagine all the products to differ. They are all pool products used for the same purpose, and if they all are claiming 70%, then I imagine
they are referring to the same thing. These pool companies do not seem to have the best chemistry knowledge it would seem. I remember one product I
saw had "111% available bromine".
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woelen
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It always says % available chlorine. If you can buy 70% available chlorine, then you have top grade stuff.
The term available chlorine is defined as follows:
If you have 100 grams of the compound/product and you add excess amount of hydrochloric acid and you obtain N grams of elemental Cl2 when all of it
has reacted, then the compound/product is said to have N % available chlorine.
In this way, some compounds can have more than 100% available chlorine, or compounds can have available chlorine, while they do not have any chlorine
in them. An example of more than 100% available chlorine is LiOCl. An example of a compound having no chlorine in it, but with larger than 0
available chlorine is KMnO4. Pure TCCA has 92% available chlorine, Na-DCCA has appr. 60% available chlorine.
Calcium hypochlorite usually comes as the dihydrate, Ca(ClO)2.2H2O. If this is 100% pure, then you have appr. 85% available chlorine. Commercial
material contains CaO, Ca(OH)2, CaCl2 and sometimes some CaCO3 as impurities. The latter is annoying, because it causes the Cl2, made with it, to be
contaminated with CO2. But usually, it is at most a few percents.
[Edited on 8-9-14 by woelen]
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FireLion3
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Great post woelen.
I will check tomorrow on this product and report back.
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careysub
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Quote: Originally posted by woelen  | It always says % available chlorine. If you can buy 70% available chlorine, then you have top grade stuff.
The term available chlorine is defined as follows:
If you have 100 grams of the compound/product and you add excess amount of hydrochloric acid and you obtain N grams of elemental Cl2 when all of it
has reacted, then the compound/product is said to have N % available chlorine.
In this way, some compounds can have more than 100% available chlorine, or compounds can have available chlorine, while they do not have any chlorine
in them. An example of more than 100% available chlorine is LiOCl. An example of a compound having no chlorine in it, but with larger than 0
available chlorine is KMnO4. Pure TCCA has 92% available chlorine, Na-DCCA has appr. 60% available chlorine.
Calcium hypochlorite usually comes as the dihydrate, Ca(ClO)2.2H2O. If this is 100% pure, then you have appr. 85% available chlorine. Commercial
material contains CaO, Ca(OH)2, CaCl2 and sometimes some CaCO3 as impurities. The latter is annoying, because it causes the Cl2, made with it, to be
contaminated with CO2. But usually, it is at most a few percents.
[Edited on 8-9-14 by woelen] |
The manufacturer's MSDS for the widely available brand HTH Dry Chlorine Granular states that their product is 60-80% calcium hypochlorite, 10-20%
NaCl, 5.5-10% H2O, with possible minor amounts (0-5%) of CaCl2, Ca(OH)2, CaCO3, and calcium chlorate.
Another widely available brand, Zappit by PPG, states that is >70% calcium hypochlorite, lists no NaCl, but <2-3% Ca(OH)2, CaCO3, and calcium
chlorate, and 0-0.8% lanthanum carbonate (!):
http://www.kellysolutions.com/erenewals/documentsubmit/Kelly...
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FireLion3
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Okay here is what I have...
It is called Refresh +
Calcium Hypochlorite - 73%
Other Ingredients - 27%
Minimum Available Chlorine - 70%
I checked the MSDS...
CALCIUM HYPOCHLORITE 7778-54-3 60 - 80
SODIUM CHLORIDE 7647-14-5 10 - 20
CALCIUM CHLORATE 10137-74-3 0 - 5
CALCIUM CHLORIDE 10043-52-4 0 - 5
CALCIUM HYDROXIDE 1305-62-0 0 - 4
CALCIUM CARBONATE 471-34-1 0 - 5
Water 7732-18-5 5.5 - 10
Should I just weigh this purely off the 73% listed in the title? The reactions I am doing wont be harmed by most of that stuff in there. Not sure
about Calcium Chlorate.
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woelen
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What is important is the minimum available chlorine. Weigh it according to this. So, assume that for each 100 grams of weighed solid you get 70 grams
of elemental Cl2. Be sure to use excess HCl at a concentration of 20% or so.
Chlorate is annoying, but hard to avoid. I forgot to mention that in my previous post. Hypochlorite easily disproportionates as follows irreversibly:
3ClO(-) ---> ClO3(-) + 2 Cl(-)
In a chlorate cell this reaction is exploited intentionally by favouring its occurence, but in swimming pool calciumhypochlorite it is not desirable.
If you make Cl2, then the chlorate does react very sluggishly and not all of it reacts to Cl2. Part of it reacts to ClO2. However, you do NOT have to
fear explosions due to the formation of ClO2. The amount of chlorate at most is a few percent and hence you won't get much ClO2. It can, however,
disturb some sensitive reactions. An example is that with manganese(II) solutions in conc. HCl, the ClO2 causes formation of a very dark brown color,
while hypochlorite and chlorine do not do this. For all practical purposes, if you use this product for making chlorine, then you get perfectly useful
stuff. Make the chlorine, bubble it through water and then use it in your reaction. Bubbling through water removes adhering HCl from the hydrochloric
acid and traces of ClO2 (if any are formed at all) also are held back by the water. Of course, if traces of HCl and ClO2 are not disturbing your
reaction, then you do not need to pass it through water. If the chlorine must be dry, then you can pass it through conc. H2SO4 (AFTER passing it
through water, not BEFORE) or through a long column of anhydrous CaCl2.
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FireLion3
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I don't plan to be generating elemental chlorine. I want to use the Hypochlorite as it is. Had used TCCA before but it is just so messy in the sense
that it is very difficult to remove from the reaction without distillation. I didn't want to use hypochlorite because I read explosive things about
it, but I figure with careful addition and good cooling then it should be fine. I will be using a phase transfer catalyst so the reaction will be
quicker than it usually is.
How does the reactivity of chlorate differ from hypochlorite in relation to alcohols and aromatics that are getting chlorinated/oxidized? I looked at
the wiki pages but couldn't find much information other than the explosive capacity of chlorates.
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woelen
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Please do not mix hypochlorites with alcohols. That is asking for trouble. Mixes of alcohols and calcium hypochlorite are abused by k3wls to make
so-called chlorine bombs. Such mixes slowly heat up and then suddenly self-ignite and if confined in a bottle, they explode!
Hypochlorites are much more reactive than chlorates. The few percents of chlorate in the mix will not do anything at all.
Could you please explain what compound you want to chlorinate and what procedure you have in mind to accomplish this?
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FireLion3
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Aside from chlorinating various aromatics.... chlorination of alcohols with Hypochlorites is a very common reaction. Exothermic, yes, but I would not
be doing it without the much needed cooling.
I have chlorinated several different primary alcohols, methoxybenzyl, benzyl, and dimethoxybenzyl alcohol with TCCA as an alternative to the synthesis
linked below
https://www.erowid.org/archive/rhodium/chemistry/ether2ester...
The reaction with TCCA is indeed exothermic. The first time I ran it on a half molar scale (stupid in retrospect) in an open beaker, I added the
entire quarter molar amount of TCCA, walked away for 10 minutes, came back, and found the contents of the flask has overflowed from the mixture not
long after I left. The table was stained with the cherry-type smell of aromatic aldehydes, metal rusted from the generated HCL, but other than that no
real damage, no broken glass, etc.
When I repeated this TCCA reaction with 1/10th additions and a decent cooling (jacketed beaker cooling), I found the heat was generated almost
immediately upon addition, and the heat generation stopped about 5 minutes after each addition. Giving me a form of experimental data to calculate how
much heat was being released with each portion I was adding, allowing me to perfect the reaction. With the jacketed beaker cooling it, I was able to
add the portions every 3 minutes, allowing for a fairly fast reaction. The biggest problem I have with this is that I have to vacuum distill the
mixture since Isocyanuric Acid and TCCA appear to be moderately soluble in the organic layer.
I imagine the reaction with the hypochlorites may be more exothermic considering that along with the oxidation energy release, is a simultaneous
solvation of NaOH and neutralization of formed HCl. Speeding this latter reaction rate up with a PTC catalyst will make it more exothermic by nature
of its faster speed, allowing for a more controllable reaction by careful addition, instead of having to add small portions every 30 minutes. I of
course will be doing this on a small scale first with slow additions and documenting temperature increases.
Are you telling me that in spite of hypochlorite alcohol oxidations being a very common reaction, that this reaction will explode [ignite] on
me, even with cooling? Or is this only if I recklessly add everything at once?
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woelen
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If you add hypochlorite in small portions and check whether a portion has reacted or not, before adding the next portion, then you do not have to fear
fire or explosion. I warned you, just to be sure that you know of the risks, but from what I gather from your last post, I have the impression that
you do it in a safe way.
I have seen movies of stupid k3wls on Youtube of bottles, filled with denatured alcohol and calcium hypochlorite, which exploded, giving fire and
glass shrapnel all around. Hence the warning.
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FireLion3
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Thanks for the heads up. Question 1: Does the risk of fire or explosion only occur when large quantities are rapidly reacted in a
small medium, or is there a tendency to randomly spark upon energy release? Your response makes it sound like the former is the case, in that the risk
of fire is when large quantities chain react.
Question 2: Do you have a possibly idea on why hypochlorite takes so long to react with alcohols like this? It almost sounds like
there is an very exponential curve with its reaction rate, that causes it to be so explosive. If this is the case, to maintain
optimal reaction rates, the reaction would need to be held perfectly at that temperature, and cooled back to that temperature continuously, which I
can manage with my equipment.
The question is then, what is the temperature which the reaction rate really takes off?
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woelen
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There only is a risk when large quantities are mixed together.
The problem with slow initial reaction and progressively speeding up reactions is not specific for this reaction. It is a common risk with many
exothermic reactions. Each 10 degrees of increase of temperature may lead to a 2 to 3-fold increase in reaction rate. The increased reaction rate then
results in more heat production in a shorter time, which cannot escape from the mix quickly enough and you understand what can happen in such
situations.
Just read this thread for a really spectacular example of this phenomenon. At the end of the thread I present a web page I have written about this
reaction: http://woelen.homescience.net/science/chem/exps/hydroxylamin...
This is far from your experiments, but it is a nice example of what runaways and in general kinetics of reactions can do.
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FireLion3
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What I'm hoping to do is hook my temperature controller to the water pump attached to the jacketed flask, heat the flask to 40-70 degrees, and have
the pump turn on if it exceeds the set temperature. When the pump turns back off after the target temperature is reached, I will add the next bit.
Seems like an efficient way to go about it. I just need to figure out a good temperature that won't prove dangerous and that won't prove to be super
slow.
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FireLion3
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Needless to say I just discovered the "run away" temperature...
When the reaction hits 50 degrees with my jacketed beaker having ice water pumped through it, the reaction can be brought down to 30 degrees in under
a minute.... when the reaction hits 55 degrees, my iced beaker stands no chance. 10 seconds after it hit 55 degrees It began to gain almost 1 degree
per second, with my water pump still on. One minute later it was boiling.... I panicked I was when I realized I didn't have anything to dunk the
reaction in. Whoops. Managed to throw a hand full of ice cubes in there.
As far as I can tell, the reaction rate was 100x faster at 55 than it was at 45...
How annoying. Now I must figure out how to solve this problem. It's so difficult to tell how much is reacting at the different temperatures because
the rates are so exponential.
Tomorrow I am going to try using a phase transfer catalyst to see if I can get more a more immediate reaction at lower temperatures, so as to better
control it.
[Edited on 13-9-2014 by FireLion3]
I did find a paper that indicates using phase transfer catalysts allows these alcohol oxidations to take place at room temperature. They did use
sodium hypochlorite, but I do not think it matters whether it is sodium or calcium hypochlorite, since the anion exchange onto the phase transfer
catalyst will be where most of the magic happens.
[Edited on 13-9-2014 by FireLion3]
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