bbartlog
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(Re)drying CaCl2?
I generally acquire my drying agents OTC - MgSO4, CaSO4, CaCl2 and so on. Most of these, when dried in a crucible or in a microwave, eventually turn
crumbly and are not too hard to deal with; for example MgSO4 seems to shrink a bit on final dehydration and can easily be removed from a glass
container, and is pretty brittle once dry in any case.
Calcium chloride, on the other hand, seems to turn into an intractable cement that attaches itself firmly to either glass or cast iron. Any tips for
drying this? Am I just not heating it quite enough (does it eventually turn crumbly at high enough temp)? The industrial answer to this is 'build a
drying tower' but I'm hoping there is some simpler solution. Does it adhere less to other metals (copper or aluminum)? The hardness is obviously
something that a hammer can address, if I can just get it out of the container it's in.
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vulture
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Does your CaCl2 melt before turning into this cement? If so, you could try shock cooling with compressed air for example.
One shouldn't accept or resort to the mutilation of science to appease the mentally impaired.
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aonomus
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What I have found sometimes is that if you have fully hydrated material, and if you heat too quickly, the salt dissolves in its own water of
hydration, forming a rock hard cake. Perhaps try heating very slowly, or with a large surface area to promote faster evaporation of water before
droplets and a liquid layer can form
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bbartlog
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Well yes, the calcium chloride does normally form a hydrate melt while I'm trying to drive off the water. So by the sound of it I could dry it
successfully if I just did it slowly enough, rather than trying to heat it past the point where it turns into a bubbling liquid...
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User
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I have the same problem with drying.
So yeah well to be honest is sticked to partially redrying it in a regular oven at 250 degrees C.
About 1/3 of the weight can be removed this way, well not ideal but i got insane of those rock hard lumbs.
I do like hammers but not really when doing chemistry.
A sexy solution to this would be nice.
[Edited on 22-1-2010 by User]
What a fine day for chemistry this is.
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entropy51
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I doubt that it's worth trying to dry hydrated CaCl2. I buy it for less than $1 per kg as ice melt. It appears to be nearly anhydrous based on the
heat of solution. Just find a good cheap source and you're good to go.
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bbartlog
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I have a big bag of ice melt as well. Now possibly my bag is just old and has absorbed moisture from the air, but it's not anhydrous. I would be
surprised if commercial CaCl2 icemelt was anything less than a dihydrate. Of course, that can still provide plenty of drying, so maybe I should just
use it as is...
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aonomus
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Now another question, is it possible to dry out NaOH prills used as a drying agent in a desiccator by heating?
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woelen
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You need very strong heating to get rid of water absorbed by NaOH. Even if the solid melts (around 300 C) then it does not completely give up its
water. Another problem is that NaOH, especially if humid, quickly absorbs CO2 from the air, forming Na2CO3. The CO2 certainly is not driven off by
heating, unless you go to temperatures of 1000 C or so.
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aonomus
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Thats what I suspected, just wanted to be sure. I will probably wait till the NaOH prills turn into puddles in a dessicator I have drying out a sample
of FeCl2....
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entropy51
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I think silica gel is one of the best general purpose drying agents for dessicator use. It can be dried in just a few minutes in a hot oven. I've
been using the same batch of silica gel, which contains a moisture indicator, for over 30 years. It still works like a charm. Drierite can also be
regenerated. I rarely need anything so dry that one of these two dessicants doesn't do the trick in a vacuum dessicator.
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rrkss
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I agree with entropy. I use silica gel for my dessicators and it is easy and simple to regenerate. Type 3A molecular sieves do an excellent job in
drying air as well but exlcuding organics. I do use CaCl2 anhydrous in my drying tubes. Can be bought by the lb cheap at home depot as a product
called damp rid. That stuff is definately anhydrous and turns slimy as it hydrates when left open to the atmosphere.
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User
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I just found out that CaCl2 dries very nice in the microwave.
It heats up when "wet" and stay's cold when dry.
I operated the microwave at 200 watts for a total of 5 minutes or so.
This was enough to fully dry a sample of 30 gr.
(calculated with use of analytic balance)
The 5 minute run was done in 3 times to let it cool down for a moment and prevent super heating of the water for hydrated crystals get very hot when
exposed to MW radiation.
Also it is very important to give it a good stir to make sure that no water gets trapped.
It is unpredictable and thus potentially dangerous.
The previous described procedure was done outside and remote controlled.
So far no nothing scary happend while dehydrating serveral substances.
What a fine day for chemistry this is.
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bbartlog
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Heating 30g of CaCl2 is unpredictable and potentially dangerous? Why? Also, nice that you used an analytical balance, but how then did you determine
the initial water content?
Lastly, what container did you use? I'm thinking that my problem may just be that I'm using glass (to which the CaCl2 sticks); do you have a
teflon/PTFE container, or what?
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User
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Well your indeed right I had no real evidence for the original hydrated state.
This could be tested though , adding a known amount of H2O next to dehydrated CaCl2.
When the water is absorbed into the crystal one measure out the amounts etc etc, you get the picture.
The only evidence I have is that the CaCl2 no longer heats up, even when the MW is ran at full blast.
I used a borosilicate beaker as well as a ceramic dish.
IMO the dish worked best.
The kind with the smooth coating for which i can't figure the right english word, the ones used in the kitchen nothing special really.
CaCl2 looses almost all stickiness when dry, at least that's my experience.
It could all depend on the amount of water in the crystals.
More experiments should be done before I can really say more.
*eidt*
I said potentially because I just dont know 
It could be if water gets trapped and builds up to much pressure, use your imagination
[Edited on 27-1-2010 by User]
What a fine day for chemistry this is.
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per.y.ohlin
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I doubt water buildup would produce any significant pressure, but a large bubble of water could cause the molten CaCl*H2O to splash out. What would
worry me more is superheating the hydrate. A pure hydrate could be superheated well beyond its dehydration point, and opening the door could cause a
bump, as with boiling sulfuric acid.
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fractional
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That's how I dry my CaCl2, too. The trick is to do this in a mortar. The dry, but cemented CaCl2 can then easily be broken up in pieces of the desired
size by just a few strokes with the pestle.
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Jor
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I use an old stainless steel cooking pan for drying my CaCl2. The technical stuff contains much water, and I want an effective drying agent for my
dessicator. So I heat with full blast of a burner, and it is quickly dry. Then it forms a VERY hard rock like cement, but I use hammers to break this.
Maybe I'll just use P4O10 in the future for dessicator, although it tends to clump together preventing more P4O10 from reacting with water. It is
pretty cheap though.
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woelen
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What I do when I want to dry something is taking 1 part by weight of P4O10 and 10 parts by weight of CaCl2 and then strongly shake this mix. The P4O10
then covers the CaCl2 and is spread out much better than when pure P4O10 is used. In this way the expensive P4O10 is well-combined with the much
cheaper CaCl2 and it is as effective as pure P4O10.
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fractional
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That's very clever! I have never seen this method mentioned anywhere. Your own invention?
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woelen
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Yes, this is what I have 'invented' myself. I noticed the nasty clumping of P4O10 when it is used as drying agent and it becomes less effective very
quickly and only a small part is really used. The dry powder becomes covered by a kind of sticky 'glaze' preventing further drying. With the CaCl2,
this glazing effect is absent and still the drying capabilities of the P4O10 remain intact.
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AJKOER
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OK, I just noticed something different about an old batch of CaCl2 that does not behave chemically the same. For example, adding new CaCl2 to aqueous
MgSO4 forms a milkly colored solution (the CaSO4, due to pH, is not precipitating out). However, my old batch apparently exposed to air, rapidly forms
a massive amount of a solid white precipitate!
Upon further research, I believe, per Atomistry.com (please see article at http://calcium.atomistry.com/calcium_oxychloride.html ), I now have some Calcium oxychloride, CaCl2.CaO.xH2O present. To quote:
"On fusing calcium chloride in moist air it becomes basic, due to the formation of an oxysalt, CaCl2.CaO."
Also, per Atomistry.com on Calcium chloride, to quote, "can be prepared by neutralisation of calcium carbonate or oxide with hydrochloric acid and
evaporation to dryness. To obtain the anhydrous salt the residue must be fused, but as the action of water vapour tends to decompose it and make it
alkaline, hydrochloric acid gas followed by nitrogen must be passed over it."
Reference link: http://calcium.atomistry.com/calcium_chloride.html
Some there appears there could be a propensity with Calcium also to form basic salts. As another example, I cited so called Bleaching Powder, which in
my somewhat novel opinion, perhaps could be describe as either containing a basic hypochlorite (in a fashion similar to Magnesium dibasic
hypochlorite) together with CaCl2, or the mixed salt of Ca(ClO)2 and Calcium oxychloride, CaCl2.CaO.H2O (which is the opinion of Wikipedia on Ca(ClO)2
and this source http://books.google.com/books?id=yZ786vEild0C&pg=PA74&am... ). There apparently has been a long history of differing opinions on the precise
chemical structure of Bleaching powder.
Interestingly, on dissolving my old CaCl2, the resulting solution is, for the most part, clear.
So be warned, as clearly the presence of CaO, may change not only the pH, but chemical reactivity as well.
[Edited on 11-7-2014 by AJKOER]
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