RogueRose
International Hazard
Posts: 1590
Registered: 16-6-2014
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MgCO3 production from MgCl2 - & use for drying ethanol?
So I had a solution of MgCl2 that I wanted to get rid of so I filtered it and added equimolar amounts (actually 2x of NaHCO3) to make MgCO3. This
stuff is extremely fluffy much like CaSO4 but is more of a mix of CaCO3 and CaSO4 in texture and feel. It filters nicely like CaCO3 as well.
When I was drying it I worried that it would get "too dry"like CaSO4 where if it is dried to anhydrous, it is difficult to rehydrate it from what I
have read. So I was worried that this might act the same way. Does anyone know if it has the same re-hydration problem?
While looking at the properties I found it is insoluble in ethanol and it forms the pentahydrate . The anhydrous is 84g/mole while the pentahydrate is
174g/mole, meaning it absorbs 107%of it's weight in water if the anhydrous can be re-hydrated. If not, then it is possible to dehydrate to di-hydrate
at 179-197C to keep it from totally drying out, and it still can absorb 3 water molecules pre mole.
Now it seems that K2CO3 was the most recommended salt to dry ethanol some time ago and IDK if MgCO3 has ever been tried, but I'm willing to give it a
go when I also test the TriSodium Phosphate. What is odd is that K2CO3 can only absorb one water molecule per mole, which is very low and I think a
lot of ethanol would be lost due to adhesion. So far the TSP, MnSO4 & this (MgCO3) looks to be the best as per quantity of water absorbed per
mole & density.
Anyone have any suggestions on this carbonate and it's hydrates?
As a side question, I can say that MgCO3 seems to be a much superior form of chalk when it comes to weightlifting and over-all feel. It also seems to
be less drying. I'm wondering if it is possible to use the pentahydrate form and have it be non-drying chalk as opposed to the other hydrates which
may be drying. Anyone have experience with chalk and know if it is a hydrate or anhydrous?
[Edited on 4-10-2018 by RogueRose]
[Edited on 4-10-2018 by RogueRose]
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AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
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Read Wikipedia at https://en.wikipedia.org/wiki/Water_of_crystallization . To quote:
" In molecular formulas water of crystallization can be denoted in different ways:
"hydrated compound⋅nH2O" or "hydrated compound×nH2O"
This notation is used when the compound only contains lattice water or when the crystal structure is undetermined. For example Calcium chloride:
CaCl2·2H2O
"hydrated compound(H2O)n"
A hydrate with coordinated water. For example Zinc chloride: ZnCl2(H2O)4
Both notations can be combined as for example in copper(II) sulfate: [Cu(H2O)4]SO4·H2O "
So, if extracting, say, just H2O from say a mix of CH3OH/H2O, I suspect you want a salt like Na2SO4 with a large number of coordinated water molecules
(like 7 or 8, see https://pdfs.semanticscholar.org/24ce/d94e1145f90c6b599db519... ).
Note, mixing MgSO4 and Na2CO3 both in solutions, per my experience, can produce a compound that consumes nearly all the water! However, my
understanding is that none of those waters are coordinated with respect to the formed Magnesium Alba (or MgCO3·Mg(OH)2·nH2O, see https://www.sciencemadness.org/whisper/viewthread.php?tid=74... , but not so for the Na2SO4), so this would not be an especially good path to
separate out water from CH3OH/H2O.
But try it out yourself.
[Edited on 10-4-2018 by AJKOER]
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