sabbath06
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calcium oxide preparation
I have been attempting to prepare calcium oxide from calcium carbonate by heating the calcium carbonate directly with a propane torch. As I heat it,
it glows a dull orange and then larger and larger areas begin to glow very brightly. I continue to heat it for a few more minutes and then allow it to
cool completely. When I add the product to a small amount of water, no reaction is observed, and the water does not increase in temperature.
I have a few theories as to why this isn't working properly. One is I might have to heat it for much longer and break the carbonate up into finer
pieces. My other theory is that the calcium oxide is being formed but is picking carbon dioxide right back up from the combustion of the propane.
Should I be heating it for far longer and in a closed container? I'm relatively sure that the propane torch is indeed hot enough. Is there anything
else that I'm doing wrong?
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Mr. Wizard
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Did you weigh the calcium carbonate before and after heating it? If your sample was the size of a match head you may not see much heat released, as
the sample is small. Concrete gives off heat when it cures, but the effect is hardly noticeable in 1 gram quantities.
At about 900C to 1000C the CO2 is driven off, so it may be possible you are adding CO2 and water to your sample with the propane flame combustion
products. Burning lime was usually accomplished with sealed piles, allowing the charcoal to react with the lime and a limited amount of air.
Allowing the lime to overheat produces dead burned lime, which reduces it's effective surface area, if I understand correctly.
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entropy51
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I've never tried to make CaO by heating directly with a torch, but I have made it many times by heating CaCO3 in a crucible until the weight loss is
close to the theoretical value for loss of CO2. That method works fine, but is rather slow. You don't want to heat it in a completely closed
container because the CO2 can't escape and will be reabsorbed by the CaO.
I would trust the weight loss more than exothermic hydration as an indication that you have CaO at the end.
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AJKOER
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I believe an easier way is to add NH4OH carefully to a solution of CaCl2 in water (or, add vinegar to your CaCO3 to work with Calcium acetate instead,
or try ammonia on CaCO3 directly).
2 NH4OH (aq) + CaCl2 (aq) --> 2 NH4Cl + Ca(OH)2 (s)
Filter out the Ca(OH)2, rinse and heat to dehydrate. Any residual NH4Cl (or other Ammonia salt) will decompose on heating.
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Mr. Wizard
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AJKER He wanted to make calcium oxide, not a solution of calcium hydroxide. I don't think your equation will work in the direction have it there.
When you add an ammonia salt with a calcium hydroxide solution, ammonia is liberated and can be smelled, or even driven off by heating the solution.
Even sweat or organic matter containing protein will give an ammonia smell when put in wet lime. Even freshly mixed concrete smells of ammonia to me.
Curiously, CaCl2 has an affinity for liquid (anhydrous) ammonia, and can actually be used as an absorbent in solar powered ammonia refrigeration
systems. I'm not sure what effect if any that would have on the mix.
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entropy51
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You might dehydrate it, if you mean dry it, but it
sure as sheet won't form CaO. Man, you are full of it. If you don't know just listen and learn something instead of passing gas.
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watson.fawkes
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Quote: Originally posted by Mr. Wizard  | | Allowing the lime to overheat produces dead burned lime, which reduces it's effective surface area, if I understand correctly. |
Dead-burnt lime doesn't slake quickly because impurities in the feedstock form glasses on the surface. Those impurities include
silica and sodium salts, which together are the main ingredients for soda-lime glass. This does indeed reduce the effective surface area. Grinding the
product will restore its ability to hydrate quickly.
I second @entropy51 and recommend using mass loss to gauge how far the reaction has run to completion.
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ScienceSquirrel
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Calcium hydroxide can be bought cheaply and easily. It readily dehydrates on heating to around 512C which is a lot lower than the temperature required
for calcium carbonate.
http://en.wikipedia.org/wiki/Calcium_hydroxide
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blogfast25
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Squirrel beat me to it.
But I have to say that an attempt to dehydrate homemade Ca(OH)2 (NaOH + CaCl2, filtered, dried then calcined) didn't yield what I expected: quick lime
is supposed to react quite violently with water, mine didn't. I wonder if along the way it had converted to CaCO3 (but I didn't test that) by CO2
absorption?
'Burning' limestone to quick lime takes forever. Britain's full of 'Lime Kiln Lanes' because of the Roman period (when lime production started).
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AJKOER
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Entropy51: As Squirrel and Blogfast25 have noted, one can indeed dehydrate Ca(OH)2. My reference gives the temperature at 580 C, which is lower than
898 C for CaCO3.
Now, on the reaction of NH4OH and CaCl2, per Modern inorganic chemistry by Joseph William Mellor, page 534, supporting your assertion of the reaction
direction:
"by heating an intimate mixture of commercial ammonium chloride or ammonium sulphate with twice its weight of quicklime, CaO, or slaked lime, Ca(OH)2.
The reaction is represented: 2NH4Cl + Ca(OH)2 = CaCl2 + 2H20 + 2NH3."
So apparently, the reverse of the reaction, I presented, proceeds upon being heated, per your point.
LINK:
http://books.google.com/books?pg=PA542&lpg=PA537&dq=ammonia+and+CaCl2&id=5t0uAQAAIAAJ&ots=eJwE9DU1_m#v=snippet&q=Ca(OH)2&f=fal
se
However, per the same source, page 664:
"CaCl2 + (NH4)2CO3 <===> CaCO3 + 2 NH4Cl"
which may support a possible argument for reversibility, as Ammonium carbonate is less basic than NH4OH. That is, it may be that:
2NH4Cl + Ca(OH)2 + Heat <==> CaCl2 + 2 NH4OH
But, if you still have doubts, check out this link, which only states the reaction in the direction I have written it:
http://faculty.plattsburgh.edu/tom.moffett/che111/chemreacti...
Also, my reading did uncover an interesting fact that rapid heating of CaCl2.xH2O (Calcium chloride hydrates) to a mere 200 C, will partially
decompose the Calcium chloride hydrate forming CaO (apparently, the rapid heating releases some of the HCl and forms CaO).
[Edited on 13-1-2012 by AJKOER]
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blogfast25
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This has very little to do with reversibility. But I'm not surprised you would think that when you're still spelling ammonia solution as 'NH4OH' and
comparing apples and oranges: ammonium carbonate and ammonia solution.
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byko3y
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I've just realized most of the CaO I purchased some time ago is just a crap. 74 g of it can be neutralized with 1 mole of HCl , while a lot of
something similar to CaCO3 now lies on the bottom of a clear CaCl2 solution. This means only half weight of the original substance is Ca(OH)2 (maybe
also some small amount of CaO, but I doubt it), which is 42:57 molar ratio of Ca(OH)2:CaCO3 (50:50 for CaO:CaCO3) provided there was no contaminations
and CaCO3 did not react at all.
I'm just wondering: is there anybody who actually bought CaO and got it usable? Looks like usage of a regular polyethylene bottle is pretty similar to
storing the CaO in open air. As far as I understand, some special non-permeable bottle should be used http://www.ebay.co.uk/itm/1kg-Calcium-oxide-High-purity-Quic...
In fact I would be glad to avoid all those problems and prepare the CaO right before using it, so I would never need to worry about its storage.
You guys have already mentioned the possibility for preparation of Ca(OH2) from the cheap CaCl2, ahowever you don't need ammonia for that because
reaction goes the opposite direction (ammonia flies away), NaOH works well for this purpose. In fact I think it's better to prepare a brand new
portion of Ca(OH)2 every time you need a pure one.
Ca(OH)2 can be dehydrated at relatively low temperature (350-400°C):
- either under vacuum http://onlinelibrary.wiley.com/doi/10.1111/j.1151-2916.1980....
- or using a dry inert gas (nitrogen here, oxygen will probably work too) http://pubs.acs.org/doi/abs/10.1021/ie00104a004
Other threads on CaO preparation from Ca(OH)2 and CaCO3 http://www.sciencemadness.org/talk/viewthread.php?tid=30122 and http://www.sciencemadness.org/talk/viewthread.php?tid=11969
What I really think to do is to obtain some salt of calcium which can be easily decomposed to obtain CaO right before its usage. Right now I have no
ideas about what substance can be used for that. CaCO3 is a no way because of high temperature; Ca(OAc)2 needs 450°C which is really close to the
maximum for pyrex glass, and it generates CO2 as a byproduct leading to CaCO3; Ca(OH)2 will be contaminated with a lot of CaCO3 unless you made a
fresh one; CaCl2 is a no way (high t°); calcium citrate decomposes into CaCO3 + C. Most researchers prepare CaO either via CaCl2 -> Ca(OH)2 ->
CaO way, or CaCO3 -> CaO.
And of course you can't allow CO2 to pass into the device where CaO is being prepared, because CO2 will ruin everything. For this reason you can't
prepare CaO using an open flame of propane torch:
C3H8 + 5 O2 => 3 CO2 + 4 H2O
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Big Boss
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This isn't really a reaction which is completed in a few minutes of heating, industrial methods heat limestone for days.
I once tried making CaO by putting some shells in a tin and putting it on our anthracite fire for a few hours, I then put some on the palm of my hand
and put a little water on it, it's a good way to test, just be sure you have a bucket of water nearby to dunk your hand into.
Kept you waiting, huh?
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byko3y
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Rate of CaCO3 decomposition strongly depends on temperature and particle size. 2 cm particles can be calcined in 1 hour at 1000°C, while 12 cm stones
require 40 hours at 1000°C. Raising the temperature to 1100°C cuts the time in half.
In fact, fine CaCO3 powder can be calcined at 550°C and vacuum in a matter of hour, but you need a really strong vacuum: at least 20 kPa/150 mmHg for
800°C, 3 kPa/25 mmHg for 700°C, 0,5 kPa/4 mmHg for 600°C.
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