Lou
Active Member
PostNovember 6th, 2008, 7:37 pm
I'll give you the short answer:
In this case, Hoke is wrong. Aluminum is inconvenient to use because it is tough contaminant to remove because it forms sticky gels (like what tin does with conc. nitric, metastannic acid goo!).
To understand why it is such a royal pain, you must know something of its aqueous chemistry.
When you go to do a cementation reaction, you are basically using a reactive metal to displace a less reactive metal. That much you know. When the aluminum or zinc goes into an acidic solution, it produces hydrogen gas. It is both the nascent (new-born) hydrogen that does the actual reduction of these metals and the chemical potential caused by the flow of electrons. Gross simplification here.
The problem with aluminum is that it has problematic aqueous chemistry and can do several things. When you put aluminum metal into this hydrochloric acid solution, you get "aluminum chloride hexahydrate" (not really, actually it is a mixture of many things: [Al(H2O)6]3+ + H2O --><-- [Al(OH)(H2O)5]2+ + H3O+, and also the benign complex AlCl4-). In short, this is hydrolysis. This is driven by the low Ksp of hydrous aluminum oxide, Al(OH)3. Now you may ask, why am I getting a seemingly basic precipitate (note the OH function in it) when I am doing a reduction at low pH in hydrochloric acid? Well, alumina is called an acidic oxide for this reason, it precipitates at low pH! Recall that water autoionizes into both hydronium (H3O+) and hydroxyl (OH-) ions. This is because water is amphoteric and acts as both an acid and a base. Well, as luck would have it, so do many other things, in particular aluminum hydroxide. The hydroxyls from water's autoionization will hook up with any AlCl4- and any hexaaqualuminum (III) floating about (which is what happens when aluminum cation is present with excess water from HCl).
So what does this mean in practical talk? It means that even at low pH of around 1 (which is equivalent to 1M HCl) you will have some aluminum hydroxide beginning to drop out because of hydrolysis. I think the actual math (which gets nasty if solved exactly, 8 or 9 unknowns) says it starts at pH 1.4. Anyhow, this gets entrained and mixed up with your cemented values and becomes hell to remove. It requires a boil in base (to complex the aluminum hydroxide as soluble aluminate, Al(OH)4-) or in a very conc. solution of HCl. Aluminum hydroxide forms alumina upon heating, a very stable entity--it is used as refractory for lining steel ladles or as an abrasive, or in sapphires in rings.
Did I also mention that this aluminum hydroxide is very flocculent and slimy and will clog filters? It is a pain to remove from your metals. Pure Al(OH)3 made by an analytical chemist with good technique and know how won't do this, but I can guarantee you, it will when you try it. Not trying to insult, but there is a lot of math, technique, and art to it. It ties in with selective precipitation techniques, also a good tool for refining! I'd be willing to discuss this all in more detail with math and examples if there is enough interest...