Sciencemadness Discussion Board » Fundamentals » Chemistry in General » Evaporation Question Select A Forum Fundamentals   » Chemistry in General   » Organic Chemistry   » Reagents and Apparatus Acquisition   » Beginnings   » Responsible Practices   » Miscellaneous   » The Wiki Special topics   » Technochemistry   » Energetic Materials   » Biochemistry   » Radiochemistry   » Computational Models and Techniques   » Prepublication Non-chemistry   » Forum Matters   » Legal and Societal Issues

Author: Subject: Evaporation Question
hodges
International Hazard

Posts: 524
Registered: 17-12-2003
Location: Midwest
Member Is Offline

Evaporation Question

I have some CuSO4 solution that was prepared by reacting a copper penny with a weak solution of H2O2 and H2SO4 (about 1M H2O2 and 1M H2SO4, reaction took several days). Can I let the solution evaporate, leaving CuSO4? Or will any remaining H2SO4 fail to evaporate? Is there any danger of any remaining H2O2 becoming concentrated and decomposing violently?

I know this is not the best way to get CuSO4. I already have plenty of CuSO4 - but I want to be able to say "Here is some CuSO4 I made from a copper penny."
t_Pyro
Hazard to Others

Posts: 120
Registered: 7-2-2004
Location: India
Member Is Offline

Mood: Volatile

H2O2 won't concentrate itself like that, don't worry (now don't we all wish that was possible!). However, the solution won't evaporate unless the acid is neutralised.
If you don't care about sodium impurities, you could use sodium bicarbonate to neutralise the mixture, else you could use barium carbonate and filter out the barium sulphate.
A better and much faster method to prepare copper sulfate would be to perform electrolysis of H2SO4 using copper electrodes.
thunderfvck
National Hazard

Posts: 347
Registered: 30-1-2004
Location: noitacoL
Member Is Offline

Mood: No Mood

I believe the H2O2 will decompose to O2 and H2O, so there's no danger there.
Yeah, I'd neutralize the acid until neutral (no more fizz fizz with baking soda), and let it evaporate.

fritz
Harmless

Posts: 49
Registered: 29-11-2003
Member Is Offline

Mood: No Mood

Try to precipiate the coppersulphate with EtOH.
I am a fish
undersea enforcer

Posts: 600
Registered: 16-1-2003
Location: Bath, United Kingdom
Member Is Offline

Mood: Ichthyoidal

Neutralise the acid with an excess of CaCO3. The resulting CaSO4 is only very slightly soluble, and so can be removed (along with the leftover CaCO3) by filtration.

If you have it, use BaCO3 instead, as BaSO4 is even less soluble than CaSO4.

1f /0u (4|\\| |234d 7|-|15, /0u |234||/ |\\|33d 70 937 0u7 /\\/\\0|23.
hodges
International Hazard

Posts: 524
Registered: 17-12-2003
Location: Midwest
Member Is Offline

 Quote: Originally posted by I am a fish Neutralise the acid with an excess of CaCO3. The resulting CaSO4 is only very slightly soluble, and so can be removed (along with the leftover CaCO3) by filtration. If you have it, use BaCO3 instead, as BaSO4 is even less soluble than CaSO4.

Sounds like a good idea, thanks. The penny is mostly gone now - about 1/4 its original thickness after 3 days. I'm going to put a mark on the glass once the penny is completely gone just to see if any water will evaporate without neutralizing. If not or once the evaporation stops, I'll add CaCO3 and filter. Had also thought of using NaOH to precipitate Cu(OH)2 and then adding more H2SO4 to this, but using CaCO3 saves this use of more H2SO4.
I am a fish
undersea enforcer

Posts: 600
Registered: 16-1-2003
Location: Bath, United Kingdom
Member Is Offline

Mood: Ichthyoidal

Digressing slighty:

What type of coin did you use as your source of copper? US\$0.01 coins are made out of zinc coated with copper. (They used to be made out of solid copper until 1981, when the price of copper rose above the face value of the coins.)

1f /0u (4|\\| |234d 7|-|15, /0u |234||/ |\\|33d 70 937 0u7 /\\/\\0|23.
meselfs
Harmless

Posts: 8
Registered: 7-2-2004
Member Is Offline

Mood: cf. FCl04

erm... if you neutralize the acid you'll wind up with some extra slat in your copper sulfate after drying.

I think you should just put it in a ceramic or pyrex container and boil it. Sulfuric acid evaporates at around 350.

meselfs himself!
hodges
International Hazard

Posts: 524
Registered: 17-12-2003
Location: Midwest
Member Is Offline

It was a 1961 penny, so it should have been all copper. The solution has turned dark blue. A copper penny weighs 3 grams, which is 0.0469 moles of Cu. That many moles of CuSO4 should weigh 7.5 grams when I'm done (if anhydrous), or 11.7 grams if *5H2O. Hopefully I won't lose too much of it due to difficulty filtering out the CaSO4.
Friedrich Wöhler
Harmless

Posts: 42
Registered: 7-2-2004
Location: Germany
Member Is Offline

Mood: No Mood

 Quote: but I want to be able to say "Here is some CuSO4 I made from a copper penny."

Ahhhhh! So speaks a scientific heart! I love it!

OK, ahm fish's method is perfectly. I do it always with sodium carbonate when build heavy-metal-salts. But not only neutralize excessed acid, I convert all Me-salt to Me-carbonate, wash it so thoroughly as I need and let it react with acid again (with a little excess of carbonate). Well, so you need more acid, but so it can be purer, when you need it as an analytic stuff.
Any amateurs repeat this "carbonizing" even more times.
Or...if its indifferentely, what the anion of your salt is, use such as acetic acid or so, from that an excess evaporates easily.

From this topic one comes to next:
1.) Oxidizing acides are not easily to get.
2.) For preparing heavy-metal-salts you need doubled quantity of acid (for oxidizing and for building salt)
3.) Emmision of NOx when using HNO3 and SO2, when using H2SO4, stinking of Chlorine, when using HCl + H2O2

Therefore I'm sure, most amateur chemists did already think about an alternativ oxidizing method. One solution can be anodic oxidation of that metal, if you have a good enough amperage.

I discovered another interesting method:
hot high concentrated NH4NO3-solution or molten NH4NO3 attacks metals like copper, silver etc. very rapidly (similar like in hot conc. acides).
But that can be dangerously. Use only small quantities of it and heat it in a large vessel. So that exothermic reaction cannot cumulate its energy. For much more security: hold a lot of water ready to use to stop reaction in emergency.

Kein Schwanz ist so hart wie das Leben.
fritz
Harmless

Posts: 49
Registered: 29-11-2003
Member Is Offline

Mood: No Mood

good idea!
the danger of an explosion can be reduced if you use ammoniumchloride instead of -nitrate. This will work too and isn´t that dangerous (O.K. HCl is produced which should not be inhaled)
Mumbles
National Hazard

Posts: 436
Registered: 12-3-2003
Location: US
Member Is Offline

Mood: Procrastinating

 Quote: Originally posted by I am a fish Neutralise the acid with an excess of CaCO3. The resulting CaSO4 is only very slightly soluble, and so can be removed (along with the leftover CaCO3) by filtration. If you have it, use BaCO3 instead, as BaSO4 is even less soluble than CaSO4.

Wont this result in precipitation of some Copper Carbonate as well? I suppose losing a little is better than not getting any.
hodges
International Hazard

Posts: 524
Registered: 17-12-2003
Location: Midwest
Member Is Offline

Once the entire penny had dissolved, I added NaOH until the solution turned clear. I washed the precipitate several times, then added dilute H2SO4 again to this precipitate until it just dissolved. Here is what the crystals looks like now that the solution has dried.
Marvin
International Hazard

Posts: 994
Registered: 13-10-2002
Member Is Offline

Mood: No Mood

Hot ammonium nitrate dissolving copper, and also mixtures of copper and ammonium nitrates during crystalising have been known to detonate.

Electrolysis is a pain becuase you need a salt bridge to prevent the copper redepositing on the cathode.

If you want another 'nice' experiment to do with your hard won copper sulphate, as opposed to the stuff you just bought, add ammonia, (light blue ppt, then intense dark blue/violet solution) in a much larger bottle. Lay ethanol, or meths carefully on top of the solution (try equal volumes) so it forms a seperate layer, cap it so it doesnt lose ammonia and wait a few days. The interface will have crusted over and crystals of the violet salt will slowly grow down. Its the only way I know of actually growing crystals of the ammonia complex.
hodges
International Hazard

Posts: 524
Registered: 17-12-2003
Location: Midwest
Member Is Offline

Thanks for the suggestion, Marvin. What purpose does the alcohol serve?

I do have some CuSO4 solution that I added NH4OH to drying (this was from bought, vs. made, CuSO4). I added an excess of NH4OH and I have a dark blue color. As the solution evaporates it leaves behind a green deposit but this eventually turns dark blue again. Still have a while to go before all this solution has evaporated. I am noticing a coating forming on the surface of the liquid, which is slowing down the evaporation. I'm guessing this is ammonium bicarbonate from CO2 in the air (I can still smell some ammonia in the solution). I believe ammonium bicarbonate will decompose on its own, leaving me just [Cu(NH3)4]SO4.

 Sciencemadness Discussion Board » Fundamentals » Chemistry in General » Evaporation Question Select A Forum Fundamentals   » Chemistry in General   » Organic Chemistry   » Reagents and Apparatus Acquisition   » Beginnings   » Responsible Practices   » Miscellaneous   » The Wiki Special topics   » Technochemistry   » Energetic Materials   » Biochemistry   » Radiochemistry   » Computational Models and Techniques   » Prepublication Non-chemistry   » Forum Matters   » Legal and Societal Issues