[Kemistry Kaiser]

I have recently constructed a crude apparatus that isolates the H₂ gas from the following formula:

2Al(s) + 6HCl(aq) 2AlCl₃(aq) + 3H₂(g)

What's far more interesting to me though is the reaction itself. Naturally, I did a little research before I considered carrying this reaction out in my backyard and noticed that solutions of Aluminum (III) Chloride are known to appear clear and colorless (or clear and faintly yellow due to ppm of Fe impurities). My product solution was grey and completely cloudy.

Afterwards I looked into possible products that would make the solution grey and completely cloudy and found that Al(OH)₃ could potentially give off such a color in access amounts. However, I seriously doubt that Al(OH)₃ was formed in an acidic solution (and yes, the solution is still quite acidic after a standardized amount of added aluminum foil has completely reacted).

I also considered Al₂O₃ as the mystery product but the only way I understand it can be formed in my system is by substantial heating (~400°C according to Wikipedia) of the Al(OH)₃, and although this reaction is very exothermic, I seriously doubt this is happening, especially considering I keep the temperature of the glass reaction vessel regulated by a cool water bath. It is also worthy of note that my source of aluminum is store-bought aluminum foil and my acid is commercial Klean Strip brand.

Does anybody know what this mystery compound is? Is there any possible way Al(OH)₃ could be produced? Is AlCl₃ even a product? An expert's opinion/explanation would be greatly appreciated.

[Borek]

What was your source of Al?

[Kemistry Kaiser]

Store-bought aluminum foil, although from what I've read about aluminum foil I doubt impurities in the foil are responsible for the substantial amounts of unknown product produced.

[Arkcon]

You have AlCl₃, in a solution of water.  Is aluminum chloride stable in water solutions?  Warning:  this is not a yes/no question -- check wikipedia and this forums own search of aluminum chloride.

[Kemistry Kaiser]

You have AlCl₃, in a solution of water.  Is aluminum chloride stable in water solutions?  Warning:  this is not a yes/no question -- check wikipedia and this forums own search of aluminum chloride.

It's stable in water, very soluble in HCl, and reacts vigorously with bases. Direct contact with it can cause burns to the skin, but I always wear forearm length rubber gloves. It's pretty toxic stuff but I dispose of it properly. How are you sure it's AlCl₃? Is it AlCl₃ mixed with something else?

[Intanjir]

Sounds like a straightforward incomplete reaction. Take a small portion of it dilute it a bit and then add lots more HCl to that and see if it clears up. The product of your reaction is quite acidic so the acidity is no evidence that the reaction is complete.

You will in general get lots of OH groups on the surface of metal oxides exposed to water. 2M-O + H2O => 2M-OH
Note that this reaction doesn't rely on the existence of OH- in solution.
Do not confuse -OH with OH-. A hydroxide ion is not the same as a hydroxide group, at least when it is covalently bound.

In an aqueous reaction it is best to suspect the production of a metal hydroxide before the oxide. Forming the oxide requires O2 and water has only a small amount of oxygen dissolved in it.

[Enthalpy]

Aluminium foil to wrap food is less alloyed than construction aluminium, but it still contains about 0.5% of other metals like Mn, Cu, Zn, Fe, Cr and a whole lot more. This should suffice for a coloration.

[Kemistry Kaiser]

Sounds like a straightforward incomplete reaction. Take a small portion of it dilute it a bit and then add lots more HCl to that and see if it clears up. The product of your reaction is quite acidic so the acidity is no evidence that the reaction is complete.

You will in general get lots of OH groups on the surface of metal oxides exposed to water. 2M-O + H2O => 2M-OH
Note that this reaction doesn't rely on the existence of OH- in solution.
Do not confuse -OH with OH-. A hydroxide ion is not the same as a hydroxide group, at least when it is covalently bound.

In an aqueous reaction it is best to suspect the production of a metal hydroxide before the oxide. Forming the oxide requires O2 and water has only a small amount of oxygen dissolved in it.

Wow, very informative, thank you! What about the chlorine in this reaction? If the grey color is indeed due to the formation of Al(OH)₃, which by what you told me I assume to be a valid product in this system, why does it form along with AlCl₃ at what seems to be the same time? Allow me to explain my question with greater detail-

The reaction vessel contained 200mL of almost distilled water (not tap water) and I added 200mL of pure HCl to that (I did this to regulate the amount of heat the reaction produced). Here's the interesting part, After I add a scrunched up ball of aluminum foil, the reaction takes over 6 minutes to start bubbling (releasing H₂ gas). At this same time I start to notice the cloudy grey material rain down off of the aluminum ball. Does this indicate that both AlCl₃ and Al(OH)₃ are being produced at the same time? If so, why is this? Shouldn't one product be mostly favorable to produce at one time?

Also as a minor question, Why does it take so long for the reaction to begin taking place after I add the aluminum? Is there a word for this phenomenon?

[Yggdrasil]

Aluminium foil is covered in a layer of aluminum oxide.  Perhaps the precipitate you see is the leftover oxide coating after the pure aluminum inside the foil gets reacted away.

[Kemistry Kaiser]

Aluminium foil is covered in a layer of aluminum oxide.  Perhaps the precipitate you see is the leftover oxide coating after the pure aluminum inside the foil gets reacted away.

So you're suggesting Al₂O₃ isn't immediately affected by the acid as the Al metal within is... very interesting. I suppose a sheet of aluminum foil does have a lot of surface area... Do you think the Al₂O₃ layer could also be responsible for the perceived 6+ minute delay in reaction?

[Intanjir]

Aluminium Oxide is a very effective passivator. http://en.wikipedia.org/wiki/Passivation_%28chemistry%29
The initial reaction probably managed to dissolve only a small fraction of the alumina, but once it exposed the metal things happened from the inside.

Aluminium Hydroxide is not grey. I think the grey is probably metallic aluminium which has not yet reacted due to particularly intransigent pockets of passivation.
I only claimed that in an aqueous reaction if you are producing one of the two you are probably* not producing much Aluminium Oxide and instead producing Aluminium Hydroxide because any exposed oxide is converted to hydroxide.
Aluminium Hydroxide will eventually react with HCl which is why I suggested adding more. Indeed they add it to antacids precisely to be the slower-acting neutralizer. Anyways you used 200ml of HCl, which seems like over kill for a bit of crumpled aluminium foil, probably no need to add more(You didn't happen to weigh the foil did you?). Just requires more time I guess.

*If you carefully set things up you can precipitate an oxide and not the hydroxide. It is substantially a matter of the surface area of the solid you create. Metal oxide surface exposed in water will substantially hydrate to form hydroxide. On the other hand two Metal-Hydroxides can condensate and form Metal-Oxygen-Metal and release a water. Unless you very carefully grow your precipitate to minimize the surface area you very often will have so much surface area that the interior Metal-Oxygen-Metal bonds are few in number compared to the surface Metal-Oxygen-Hydrogen.

[Intanjir]

Oh yeah and I guess someone should eventually mention that so long as there is no barrier, Aluminium will vigorously react with H2O to make Hydrogen and Al(OH)3.

[Arkcon]

Oh yeah and I guess someone should eventually mention that so long as there is no barrier, Aluminium will vigorously react with H2O to make Hydrogen and Al(OH)3.

This is ultimately what you'll get from the aqueous reaction of aluminum foil and HCl in water.  Aluminum chloride is not very stable in water solution.  I used to think it was completely unstable, but as forum experts here described, that's not correct.  Still, what you have is aluminum hydroxide.

http://www.chemicalforums.com/index.php?topic=66913.msg241105#msg241105

Store-bought aluminum foil, although from what I've read about aluminum foil I doubt impurities in the foil are responsible for the substantial amounts of unknown product produced.

Don't be so sure, there is a small amount of other elements in aluminum foil, but they are significant.  They make the metal easier to roll into foil, that was a discovery that made household aluminum foil possible.  There may be more visible impurities than you expect.

[Intanjir]

I didn't realize it lost chlorine as gas. So yeah given that I agree: when enough chlorine escapes the PH won't be low enough to keep Aluminium dissolved and as this unfolds it will precipitate as Al(OH)3 while leaving a solution of decreasingly acidic Chlorohydrate. However so long as he started with a sufficient excess of HCl this won't happen for awhile. Then again the release of hydrogen was quite exothermic, he may have lost much of his excess.

However, aqueous AlCl3 isn't unstable like SiCl4. As long as it is sealed it will remain at least as stable as sealed HCl does and will eventually dissolve Aluminium Oxide/Hydroxide. I suspect however that the eventually is probably a very very long time at room temperature.

[Borek]

It is anhydrous AlCl₃ that decomposes in contact with water, hexahydrate is - as far as I am aware - stable.

But it is more complicated in general, as there are chlorohydrates - basically polymeric mixed basic salt, Al/Cl/OH.