Can it be isolated at room temperature? Or even at any temperature?
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4 Answers
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Aluminium carbonate reacts easily with water/water vapour to form aluminium hydroxide as you can see on the Wikipedia page.
$$\ce{Al2(CO3)3 + 3 H2O → 2 Al(OH)3 + 3 CO2}$$
So, if you want to store aluminium carbonate, you have to make sure it is free from moisture.
Why it is unstable:
Smaller and more charged metal cations (i.e. metal cations with high charge density) tend to distort the electron cloud of the carbonate ions (polarisation).
Polarisation eventually leads to abstraction of oxygen from the carbonate ion, making one of the $\ce{C-O}$ bond in the carbonate weak, which makes it very prone to attack by water and heat (similar to the decomposition of copper (II) carbonate when compared to alkali metal carbonates).
The reaction itself doesn't release a lot of energy; hence why it is rather unstable and not reactive in chemistry terms.
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1$\begingroup$ So if it can be synthesised from the reaction on the Wiki page, how is it isolated? I've searched quite a bit, and can find almost nothing about it on the web, nor in any of my inorganic textbooks. Lots of reagents are sensitive to water, but that doesn't make them unstable, just reactive. Fluorine, for instance. You wouldn't think of that as unstable, although it's exceptionally reactive. $\endgroup$– ChrisACommented Oct 1, 2013 at 14:37
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$\begingroup$ @ChrisA I added some more information to my answer. I tend to miss the title of the question once I read the body :( For the production of Aluminium Carbonate, I can't say for sure, but that could be prepared without the presence of water. Both $\ce{Al2(SO4)3}$ and $\ce{Na2CO3}$ are anhydrous solids; the former is acidic and the latter basic, the reaction should occur without much difficulty even without water involved. $\endgroup$– JerryCommented Oct 1, 2013 at 15:24
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Aluminum carbonate, if existed, would decompose by the following reaction: $$\ce{Al2(CO3)3 → Al2O3 + 3 CO2}$$ to $\ce{CO2}$ and $\ce{Al2O3}$, with a release of energy, even in the absence of water. Therefore, it is unstable, and probably could not be isolated under any conditions.
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1$\begingroup$ This answer seems to be the only one so far that directly answers the question of instability even in isolation as opposed to reactivity (which could be viewed as instability in the presence of other compounds like water). $\endgroup$– Curt F.Commented Mar 2, 2015 at 5:00
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Aluminium carbonate is a salt of a very weak acid and a weak base, so it tends to hydrolyze. Similar may be said for a much easier available aluminium sulfide: it reacts with water in the air, so it is a common tool for rebellious teenagers if they can get the materials.
Moreover, aluminium cation is extremely small. It means, that it forms very strong bonds with small anions (oxygen). Since carbonate is a rather big cation, it is not a best partner for aluminium, so it should release carbon dioxide very easily.
For a better understanding consider second group carbonates. Barium carbonate is very stable, calcium carbonate dissociates around $1000^oC$ and berrilum carbonate does not form. Aluminium is an even smaller cation, than berillium.
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It does not exist in nature, but in theory you could make it in the lab by adding pressure.
From a traditional pencil and paper calculation of Gibbs energy I found the pressure needed for
Al2/3O (s) + CO2 (g) = Al2/3CO3 (s)
is about 10 GPa.
Or you could make an Al2(CO3)3 (aq) solution with concentration of $10^{-5}$ M.
Well, still not impossible. And there should be some special solvents and/or complex structures that easily allow this reaction (postulation).
