Pure Aluminium

[airkyd]

ok sure we have lots of aluminium oxide around us, drinks cans and such, but how can i get pure aluminium and how can i get so that it will not react with oxygen to for aluminium oxide??? electrolysis??? reducing agent??? and ideas any one???

[woelen]

We have lots of pure aluminium around us. The cans and foil you know are not the oxide but the (almost) pure metal.

[airkyd]

can i use the foil when i make thermite or do i need to do something with it before binding it with iron oxide?

[xeluc]

Aluminum foil is waaaay too coarse for thermite. To get powder of the grade you'd need, you'd need a ball mill to crush the foil you put into it. Ball mills aren't really expensive if you build them yourself.

[airkyd]

hmmm, i want to make one but i cant find any where that sells the axel and the ruber belt

[xeluc]

just look around your house! You can use many house hold items for stuff like this. Maybe a trip to a hardware store will help you.

[Bio-Hazard]

And slowly it turns into totse...

[xeluc]

:-\ Maybe not.. I am interested in thermite (Got the Iron oxide but never the Aluminum) and I'm not a pyromaniac...

[akcapr]

You may have to just buy the aluinum. AN easy way to get aluminum is take one of those razor scooter, and take a file to it. File down however much Al you want. IT works pretty good if you use a fine enough file.

[xeluc]

unless you don't have a razor scooter... haha.. I used too.. Even had socks on it. But Aren't those a bit pricey for an aluminum source?

[RyanJ]

unless you don't have a razor scooter... haha.. I used too.. Even had socks on it. But Aren't those a bit pricey for an aluminum source?

 

Indeed, it would probably be cheaper just to but pure aluminium powder

 

Can't you just shred aluminium cans? I did that with one of those industreal shredders and it worked fine

 

Cheers,

 

Ryan Jones

[spudpeel]

Im glad its not just me who's been trying to make thermite. You can get hold of aluminium powder, mainly off ebay. I got iron oxide but havnt been bothered to get the aluminium. Its a wee bit pricey.

[RyanJ]

I've never actually done a thermite reaction yet but I've used Amilinium to test its reactivity with various acids or bases. That Aluminium Oxide is a pain, makes Aluminium seem rreally unreactive

 

Cheers,

 

Ryan Jones

[woelen]

I've never actually done a thermite reaction yet but I've used Amilinium to test its reactivity with various acids or bases. That Aluminium Oxide is a pain' date=' makes Aluminium seem rreally unreactive

 

Cheers,

 

Ryan Jones[/quote']

If you want to see the real reactivity of aluminium, then add some aluminium (foil, powder, any form will do) to a solution, containing both copper (II) ions and chloride ions. So, a solution of CuCl2 is OK, but a concentrated solution of plain table salt and some copper sulfate also is OK. You'll be surprised about the vigorous reaction. The chlorocuprate (II) complex at once destroys the oxide layer (I do not know why, in fact, this is one of the still badly understood parts of chemistry) and the exposed aluminium reacts with water very vigorously.

[RyanJ]

If you want to see the real reactivity of aluminium, then add some aluminium (foil, powder, any form will do) to a solution, containing both copper (II) ions and chloride ions. So, a solution of CuCl2 is OK, but a concentrated solution of plain table salt and some copper sulfate also is OK. You'll be surprised about the vigorous reaction. The chlorocuprate (II) complex at once destroys the oxide layer (I do not know why, in fact, this is one of the still badly understood parts of chemistry) and the exposed aluminium reacts with water very vigorously.

 

Sounds like some fun!

 

I'd also be interested to know why the oxide layer is removed - does this happen in other cases, I thought the oxide layer would protect it against most things :S

 

Cheers,

 

Ryan Jones

[xeluc]

The oxide layer protects the pure aluminum beneath from anything except for things powerful enough to disolve the oxide, lol. The oxide layer will protect from things such as Oxygen and water, but other things like HCl and NaOH are too pwerful for it.

[woelen]

The oxide layer protects the pure aluminum beneath from anything except for things powerful enough to disolve the oxide, lol. The oxide layer will protect from things such as Oxygen and water, but other things like HCl and NaOH are too pwerful for it.

Yes, but with the copper/chloride combination there is something special. That mix is not very powerful, yet capable of dissolving aluminium. In fact, just plain table salt and some non-acidic copper (II) salt are much more effective than both HCl and NaOH, it is really remarkable.

 

If you have some CuCl2 left from your copper experiments, then try the following:

- Prepare a solution of table salt, quite concentrated. If it does not become nice and clear, then add a few drops of HCl, but not too much. If you add some aluminium household foil to this solution, then it does not react. Even with a small amount of HCl in it, it does not react.

- Also dissolve some CuCl2 in that salt-solution. It need not be acidic, but if some acid is left in the CuCl2 it does not hurt.

- Immerse some Al-foil in the copper(II)/salt solution. Be careful, this reaction is quite vigorous at once!!

 

I made a web-page about this experiment:

 

http://woelen.scheikunde.net/science/chem/exps/cu+al/index.html

 

A really nice experiment is the last paragraph in smaller font. This, however, it quite dangerous, so be careful!!!!! You can do this with CuCl2 instead of the sulfate. So, take concentrated salt solution with a small amount of HCl to make it nicely clear. Immerse the Al-foil in this, and see that it does not react. Then put the solid copper salt on it and step back.

 

Dissolving of Al in HCl or NaOH can be explained easily (oxide form Al(3+) and water with HCl and it forms aluminate AlO3(3-) and water with NaOH). But the fact that essentially non-acidic copper(II)/chloride solution is so effective remains very remarkable to me.

 

With this Cu(II)/Cl(-) combo and Al-foil, you also have a nice and very cheap source of making lots of hydrogen gas! Fun assured!

[xeluc]

I still have my 20 Ounce bottle cookin up Cu Ions. So I'll have some soon.. or not. I have been doign this for like 4 days. I know that I have more Cu ions in solution, as I have to keep putting more wire in, but when I take off hte cap, the acid STILL fumes, indicating that its still very strong. I just can't find a good way to get the Cu+ Ions to oxidize. Putting new oxygen into the bottle works, but not very fast. H2O2 would work, but the lowsy 3% I have will dilute the acid pretty fast. So I don't know. Anyhow, I will try this experiemtn when I get a chance. I also found out that NaOH in solution will make H2 indefinatly as long as you have water and aluminum, but maybe the CuCl2/NaCl (Do I need both when I'm using CuCl2? (The CuCl2 has chlorine already...)) solution will work faster, although I'm not sure if I need that. The reaction with the NaOH solution is so exothermic that I had to put the bottle down because it got too hot. Soon after, the bottle actually shrunk, which was a cool thing to see.

[akcapr]

Woelen, you never mentioned what happened to the Na+ in that reaction, or is it unknown or did i just miss it?

[xeluc]

Woelen, you never mentioned what happened to the Na+ in that reaction, or is it unknown or did i just miss it?

 

No, I'm not Woelen, but basicly the Na+ is a spectator Ion. Woelen mixes NaCl and CuSO4 together because he needs Cu2+ Ions and Cl- Ions. The Na and SO4 Ions don't do much of anything. This is talking about his experiment though. Click on his link to view. When he said to mix CuCl2 and NaCl, I am pretty sure that was jsut a mistake, as CuCl2 already possesses the neede Chloride Ions, so adding NaCl will have no effect on the outcome of the reaction.

[woelen]

No, I'm not Woelen, but basicly the Na+ is a spectator Ion. Woelen mixes NaCl and CuSO4 together because he needs Cu2+ Ions and Cl- Ions. The Na and SO4 Ions don't do much of anything. This is talking about his experiment though. Click on his link to view. When he said to mix CuCl2 and NaCl, I am pretty sure that was jsut a mistake, as CuCl2 already possesses the neede Chloride Ions, so adding NaCl will have no effect on the outcome of the reaction.

Xeluc, you are very right with your explanation about spectator ions. They just are there, looking to all the interesting things the Cl(-) and Cu(2+) ions are doing .

 

The reaction indeed works with CuCl2 only, but as this is a scarce resource you can better use some NaCl/CuCl2. The latter also works and NaCl is cheap and available without limits. From my experiments I have the impression that the reaction rate is proportional to [Cu(2+)]*[Cl(-)]. So, having a low concentration of Cu(2+) and a very high concentration of Cl(-) may be as effective as a medium concentration of both, that it why I suggested you to use a very concentrated solution of NaCl with some solid CuCl2 added. I encourage you to experiment a little with both chems and see what is optimal.

 

Beware, this reaction also is quite exothermic.

[xeluc]

CuCl2 jsut became a little less scarce instead of opening my 20 ounce bottle and letting HCl fumes cloud up my room, I filled a prescription canister half full with my stuf and threw in excess copper wire. in a day or so I should have a very concentrated CuCl2 solution.. yay.. CuCl2 is very special to me, it is the very first inorganic compound I made myself not in Chem class. Also, of course you were the one to help me Woelen, on chemical forums. I need to buy some copper sulphate. I heard it's pretty easy to find in stores but i havn't had any luck..