Can it be isolated at room temperature? Or even at any temperature?


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.

  • 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. $\endgroup$
    – ChrisA
    Oct 1 '13 at 14:37
  • $\begingroup$ @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. $\endgroup$
    – Jerry
    Oct 1 '13 at 15:24

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.

  • $\begingroup$ 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.
    Mar 2 '15 at 5:00

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.


aluminium form smaller cation,since carbonate ion is larger in size thus greater polarization lead to weak covalent bond hence at room temperature aluminium carbonate tend to break down.

  • $\begingroup$ Welcome to chemistry.SE. What has polarization to do with a weaker covalent bond?! Your answer is obviously incorrect, unfortunately; at least the reasoning is. $\endgroup$
    – M.A.R.
    Mar 30 '15 at 16:36
  • $\begingroup$ @MARamezani see this presentation cosweb1.fau.edu/~warburton/Fall2014/GLY4200C_F14/… $\endgroup$
    – DavePhD
    Mar 30 '15 at 17:10
  • $\begingroup$ @DavePhD I stand corrected. But maybe I should come back and add a more elaborate answer.... $\endgroup$
    – M.A.R.
    Mar 30 '15 at 17:24

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