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Topic: Aluminium Hydroxide Solubility  (Read 7773 times)

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Offline Amanita_Virosa

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Aluminium Hydroxide Solubility
« on: October 17, 2008, 04:20:46 AM »
Hi everybody,

I'm investigating aluminium toxicity but I'm slightly confused by some of the speciation data that I have.

At circum-neutral/slightly acidic pH nearly all aluminium ions (in the absence of other ligands) are going to be in the form of aluminium hydroxide.  As I am aware Al(OH)3 is so near insoluble as to be undetectable i.e the species Al(OH)3 (aq) does not exist.  As a result if one was to add a simple aluminium salt, say aluminium chloride, to a solution at pH=6.8-7.4 then pretty much all of the aluminium should precipitate out as amorphous aluminium hydroxide, presumably in a colloid form to begin with.

The problem I have is that I am trying to model the speciation using computer software (the program I am using is Visual MINTEQ - a windows version of MINTEQA2 - which I believe is a highly respection speciation modeller) and I am getting some results that I don't understand.  In order that I didn't confuse the situation I first simply modelled the speciation of Al in a solution at pH 6.8 with no other ligands.  As expected the amount of free Al3+ was very small, consistant with all the data I had previously seen but a large compontent was Al(OH)3 (aq) - a species that I thought didn't exist.  It wasn't a small amount either; it was the dominant species and no mention was made of any solid phase.  The overall solubilty was 100% with no precipitate.

The software contains both amorphous aluminium hydroxide and gibbsite in its database and if you model the dissolution of these at the same pH 99+% remains in the solid phase.  It produces a similar speciation model with the Al that does go into solution and, once again, Al(OH)3 (aq) is present.  Even if this was to represent a quirk of the software and a large amounts should be precipitaing from an Al salt, it still suggests that Al(OH)3 (aq) can and does exist in reality as even dissolution of gibbsite produces it (albeit in small quantites) in this model.

My question, therefore, is twofold.  a) does this mean that Al(OH)3 (aq) actually exists in reality b) if so, under what circumstances would this occur rather than precipitation that I would have expected?

Sorry to ramble on a bit but I wanted to be as clear as possible about the source of my confusion!!!

Thanks you in advance for any help you may be able to give me - I am very puzzled.

Offline nj_bartel

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Re: Aluminium Hydroxide Solubility
« Reply #1 on: October 17, 2008, 10:23:20 AM »
How do you mean Al(OH)3 exists?  You mean, the substance is dissolving in water somehow, but you're not sure how (no aluminum ions or hydroxide ions).  Have you looked at the possibility of the aluminum hydroxide forming a hydrate?

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