Jinc8
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How pure can CaCl2 extracted from eggshells be?
Greetings (first post on the forum!), I am planning on extracting CaCl2 from eggshells and HCl, but I have the following question: Is there
going to be so much organic/inorganic contamination in the final solution that makes the whole process not worth trying?
Sidenote: is there any chance the eggshell's CaCO3 won't even react with the HCl? I'm seeing surprisingly few mentions of this extraction
method online...
I actually dont yet have any use for Calcium chloride and i'm making it just for fun, so any ideas on fun experiments with it would also be highly
appreciated 
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B(a)P
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Welcome to the forum. I am not sure how long you have been storing your egg shells, but it would be worth giving them a good wash with water,
preferably as soon as you crack the egg to get rid of as much albumen as possible (I only suggest this to avoid odour in the next step). Then torch
them or stick them in an over for a while to get rid of the rest. Don't worry too much as the HCl should take care of the residual proteins. You will
then need to crush them up, having baked them will help with this. As per the below link, egg shells are mostly CaCO3. Once you have them
crushed weigh your sample of shells and work out how much HCl you require.
Calcium from egg shell link
As for usage, it is one of the chemicals that I get through the most of. I use it exclusively for drying, not particularly interesting. I have tried
drying it a few times to reuse it, but nothing else. I keep thinking there must be something interesting to do with it, but I have never spent much
time looking into it. Hopefully you get some interesting ideas from others.
I look forward to hearing how you go.
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CharlieA
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I don't know what your experience/knowledge/equipment/reagent levels are. Assuming that they are low or moderate, here are my suggestions (Rx: take
with a grain of salt).
An important aspect of chemistry is the qualitative and quantitative analysis of the product(s) of a reaction or other process. B(a)P's suggestion for
isolating CaCO3 from eggshells is a good one.
The CaCO3 you isolate can then be analyzed qualitatively and quantitatively for Ca, carbonate, and water content. After converting the CaCO3 to CaCl2
and isolating the product, it can then be analyzed for Ca, chloride ion, and water.
Aside from the intellectual benefit of identifying compounds both qualitatively and quantitatively, there are other benefits:
-researching the analytical methods is good learning experience;
-if your laboratory chemicals and equipment are limited, gathering the needed materials for the analyses provides a plan for systematically expanding
your laboratory.
Whether you have any use for the CaCl2 you prepare is really mostly irrelevant. The improvement in your chemical skills and your laboratory will be of
great benefit in future preparations.
Good luck in you endeavors.
p.s. Remember, I wrote the above assuming little experience/equipment. If I underestimated the situation, just write off the above as the rambling of
an old "wet" chemist.
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Jinc8
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Many thanks to both of you for your answers, first of all, I am quite new to the whole chemistry thing (I currently only have a bunch of pretty
inorganic copper salts to show but its been extremely fun), to be honest I do actually find just the learning experience to be worth my time.
I also somehow happened to forget the amount of times CaCl2 has been mentioned as a drying agent (thanks for the reminder), plus Ι could use it to
make Ca(OH) which is apparently also really useful.
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Texium
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Anhydrous calcium chloride is a good drying agent. You need to actually bake it in an oven to fully dry it though. You can also buy anhydrous calcium
chloride as “DampRid” or equivalent products at most hardware stores. You can also buy very pure food grade stuff from homebrew stores and online
for cheap. Making it from scratch, while not a bad experiment, is not incredibly practical.
My recommendation would be to dissolve the eggshells in HCl, filter the solution if anything doesn’t dissolve, and then re-precipitate the calcium
carbonate by adding an excess of sodium bicarbonate to the solution. This should yield quite pure calcium carbonate, which can be filtered off and
washed with water to remove any impurities that remain dissolved. Sure, calcium carbonate is cheap and readily available too, but in this way you
would be regenerating the original material of the egg shells, albeit in pure form, which I think is a cool concept.
You could also precipitate calcium hydroxide from the solution by adding NaOH instead of bicarbonate, but hydroxides tend to be difficult to filter
since they like to form gelatinous messes when precipitated, and they can dissolve/clog up both filter papers and glass frits. Calcium hydroxide, too,
is cheaper than dirt and available at many hardware and garden stores as “hydrated lime” or “slaked lime.”
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Texium
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Thread Moved
24-3-2022 at 18:32 |
SnailsAttack
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I've extracted calcium acetate from seashells. Probably a similar process to what you're doing. Even roasted seashells have a ton of protein fibers in
them though, has a pretty nasty smell. Requires filtration and decantation. Purity seems decent after a few recrystallizations. I am concerned that
HCl would be strong enough to dissolve more of the fibers and make purification harder though.
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Texium
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If you just re-precipitate the calcium carbonate as I suggested, you shouldn’t have any issues with purification using chloride or acetate.
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Jinc8
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Thank you all, I've started saving up my eggshells today, I'll probably stop once I have a theoretical yield of about 50g (calculated through random
search results but that still counts). Texium's idea sounds quite compelling too, since I don't really specifically need CaCl2, and I can convert it
back from CaCO3.
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B(a)P
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If you are enjoying learning you could work out the calculations for yourself. First you need to establish the reaction scheme, this will tell you the
molar ratio between your reactants and products. You then use the molar ratios to work out the mass of each reactant required and what your
theoretical yield is. This formula gives you the relationship between the number of moles (#M), molar mass (mm) and mass of product (m).
#M = m/mm
Note that I have got the symbology incorrect, I can't remember the correct terms just at the moment, but hopefully this makes sense to you.
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