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Author: Subject: Products of CuSO4+Al+NaCl?
Draeger
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[*] posted on 1-2-2020 at 04:20
Products of CuSO4+Al+NaCl?


So, uh, I need a little help.

Basically, I added aluminum foil to a solution of copper(II)-sulfate. I slept a night after that, and the next day I noticed it had actually been eating at the foil a bit and formed some copper at various places in the flask.

Now comes the part where I did a stupid thing.

I had heard that NaCl speeds up the reaction to the point of barely being able to control it. So, I added tiny amounts of normal household salt, and it suddenly started forming a gas, which made me slightly panic. After 20 minutes of waiting, I had decided I needed a way to vent it off, whatever it was. I managed to do just that.

After that, I added enough salt for the solution to turn green and to abruptly consume most of the aluminum in the flask. I then suddenly assumed it was probably HCl being formed, so I added some sodium hydroxide to neutralize it.

Stuff participated out, and I let it settle. Now there's a green powder sitting on top of what I assume to be chunks of copper in a greyish, though rather clear solution. There are also some orange and white chunks floating around.

I honestly have no idea what half of the stuff I have there is, and I hope you can help me. I can send a picture if needed. I have really learned my lesson now, though.

Thanks in advance.
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draculic acid69
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[*] posted on 1-2-2020 at 05:11


Copper sulphate and sodium hydroxide= sodium sulfate.do u have whitish rectangular shape crystals precipitated out? that'll be the sodium sulfate
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[*] posted on 1-2-2020 at 07:08


White chunks are probably Al(OH)3, orange chunks are Al(OH)3 with some copper. Grey stuff is carbon from aluminium. Gas which was released was hydrogen.

Good article about this reaction is on woelen's website.

If you want to separate copper, add some acid in to do solution - this will dissolve Al(OH)3. Then filter the solution and you'll have clear copper.
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Draeger
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[*] posted on 2-2-2020 at 05:47


Thanks for your answers. But what could the green stuff be?
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sciece nerd
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[*] posted on 2-2-2020 at 05:59


Basic copper chloride?
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Draeger
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[*] posted on 2-2-2020 at 14:33


I've looked up basic copper chloride and read that it reacts with bases. So I checked the pH with some test paper I had lying around and it read at 14, so it probably can't be that anymore..? Not sure anymore, this is so confusing to me.
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[*] posted on 2-2-2020 at 16:14


Maybe copper hydroxide mixed with copper?
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Draeger
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[*] posted on 2-2-2020 at 16:36


Quote: Originally posted by Bedlasky  
Maybe copper hydroxide mixed with copper?

Maybe. But I want to be as sure as I can be before I do anything else. Is there any way I could test some of the theories?
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DraconicAcid
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[*] posted on 2-2-2020 at 16:44


More likely to be a copper aluminate, or mixed copper/aluminum hydroxide.



Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
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MrHomeScientist
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[*] posted on 3-2-2020 at 09:06


The color of copper chloride is sensitive to the amount of chloride ions in solution. In dilute solution it is pale blue, but in concentrated solution (or with lots of added salt) it will be bright green. You could take a few drops of the green solution and dilute and see if the color changes.


You started with:
Al + CuSO4

The aluminum and copper sulfate react in a single displacement reaction:
2Al(s) + 3CuSO4(aq) == 3Cu(s) + Al2(SO4)3(aq)

You then added salt, NaCl. The salt helps speed this up by disrupting the oxide layer on the aluminum. If the copper sulfate was in excess, so some copper was left in solution after all the aluminum was gone, you now have (essentially) copper chloride in solution. If the concentration of chloride is high, which it sounds like it was, the solution will turn green.

Notice that no gases are formed from any of these reactions. The gas you observed indeed might have been hydrogen, from aluminum reacting with water:
2Al(s) + 3H2O(l) == 3H2(g) + Al2O3(s)

Finally, you added NaOH. The result here depends on which reagent was in excess, the aluminum or copper sulfate. One of these would happen:
2NaOH(aq) + CuSO4(aq) == Na2SO4(aq) + Cu(OH)2(s)
6NaOH(aq) + Al2(SO4)3(aq) == 3Na2SO4(aq) + 2Al(OH)3(s)

If the first was true, you'd see a pale blue precipitate that eventually turns black (decomposing into CuO). If the second, you'd get a white/grey solid. Sounds like the second was what happened.

So to summarize, your final mixture likely has:
Cu(s) - reddish powder
Al2O3(s) and/or Al(OH)3(s) - white/grey powder
Na2SO4(aq) - colorless in solution
Some leftovers of your initial reactants.
Any insoluble impurities from the initial aluminum foil alloy (Si, Fe, etc.).

If your final solution is colorless, you've dropped out all the copper either by reaction with Al or precipitation with NaOH. The green powder may be basic copper carbonate, coming from the copper that you precipitated that re-oxidized in the salty solution.

For disposal, filter off the solids and neutralize the liquid with HCl or another appropriate acid. Throw away the solids with regular trash and pour the liquid down the drain with plenty of water.
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Draeger
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[*] posted on 3-2-2020 at 09:37


Quote: Originally posted by MrHomeScientist  
The color of copper chloride is sensitive to the amount of chloride ions in solution. In dilute solution it is pale blue, but in concentrated solution (or with lots of added salt) it will be bright green. You could take a few drops of the green solution and dilute and see if the color changes.


You started with:
Al + CuSO4

The aluminum and copper sulfate react in a single displacement reaction:
2Al(s) + 3CuSO4(aq) == 3Cu(s) + Al2(SO4)3(aq)

You then added salt, NaCl. The salt helps speed this up by disrupting the oxide layer on the aluminum. If the copper sulfate was in excess, so some copper was left in solution after all the aluminum was gone, you now have (essentially) copper chloride in solution. If the concentration of chloride is high, which it sounds like it was, the solution will turn green.

Notice that no gases are formed from any of these reactions. The gas you observed indeed might have been hydrogen, from aluminum reacting with water:
2Al(s) + 3H2O(l) == 3H2(g) + Al2O3(s)

Finally, you added NaOH. The result here depends on which reagent was in excess, the aluminum or copper sulfate. One of these would happen:
2NaOH(aq) + CuSO4(aq) == Na2SO4(aq) + Cu(OH)2(s)
6NaOH(aq) + Al2(SO4)3(aq) == 3Na2SO4(aq) + 2Al(OH)3(s)

If the first was true, you'd see a pale blue precipitate that eventually turns black (decomposing into CuO). If the second, you'd get a white/grey solid. Sounds like the second was what happened.

So to summarize, your final mixture likely has:
Cu(s) - reddish powder
Al2O3(s) and/or Al(OH)3(s) - white/grey powder
Na2SO4(aq) - colorless in solution
Some leftovers of your initial reactants.
Any insoluble impurities from the initial aluminum foil alloy (Si, Fe, etc.).

If your final solution is colorless, you've dropped out all the copper either by reaction with Al or precipitation with NaOH. The green powder may be basic copper carbonate, coming from the copper that you precipitated that re-oxidized in the salty solution.

For disposal, filter off the solids and neutralize the liquid with HCl or another appropriate acid. Throw away the solids with regular trash and pour the liquid down the drain with plenty of water.

Thank you! Next time I'll definitely do more research into what I'm actually doing.

[Edited on 3-2-2020 by Draeger]
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[*] posted on 3-2-2020 at 11:39


While experimentation is fun and useful, you should definitely have a pretty good idea of what's going to happen before doing it! Think of possible reactions and what dangers they might present, and how you can prepare for them. Planning for neutralization, cleanup, and disposal is also a good idea.

Couple extra notes:
Quote:
The gas you observed indeed might have been hydrogen, from aluminum reacting with water:
2Al(s) + 3H2O(l) == 3H2(g) + Al2O3(s)

Normally aluminum won't react with water (if it does, it's very slowly), but the salt disrupts the protective oxide layer and exposes the bare metal to attack. For how much we use it in everyday life, Al is surprisingly reactive.

Quote:
Any insoluble impurities from the initial aluminum foil alloy (Si, Fe, etc.).

Any alloyed iron would, of course, quickly rust in the salty environment and you'd see a brown fluffy powder. Finely divided silicon would look black.
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