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12AX7
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New method for Cu2O:
Electroplate some colloidial copper. I happen to have some on hand from dissolving American nickels (25% Ni, balance Cu) at an absurd current density
in HCl. Some is also produced by oxidation of copper in NaCl solution, if your electrodes are spaced too closely. The copper sponge is broken up,
washed, dried and sifted.
About 100-200 mesh and greater appears to react with HCl acid, producing a white precipitate. I got best results by lightly moistening the powder,
stirring to make a suspension, then adding HCl to form Cu2Cl2. (If you don't use much water, it'll end up a thick milk shake consistency,
which appears to be just fine.) Pour this off into Na2CO3 solution slowly (CO2 bubbles are produced instantly, almost at the rate of splashing!) and
wash the yellow (looks like lead oxide) precipitate.
Well, if you can. I'm waiting for it to settle right now.
I'd still like a method that produces CuO though; HCl doesn't appear to oxidize CuCl further, even though it should?
Edit: well duh, of course it does, but it appears metallic copper present is able to spontaneously reduce it and leave shiny metallic crystals.
- Can Cu2O be roasted to CuO? (Being careful not to get a flame anywhere near it of course, lest it reduce to Cu.)
Tim
[Edited on 12-5-2005 by 12AX7]
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12AX7
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| Quote: | Originally posted by 12AX7
- Can Cu2O be roasted to CuO? (Being careful not to get a flame anywhere near it of course, lest it reduce to Cu.)
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Yep-- works great. Turns black slowly at medium temperatures, probably possible to spread Cu2O on a cookie sheet and bake it at 500°F for a few
hours. I heated it to red heat while stirring (at this temperature, any CuClx also evaporates) then dropped the powder through the air. Oddly I
didn't have much trouble with oxidation in the flame, raw copper oxide appears to reduce relatively slowly in a neutral flame.
Mixed with magnalium it burns quite a bit faster than average thermite. 
Chemleo, you said yours burned at a normal rate? Maybe it wasn't stoichiometric? Low grade of CuO?
Tim
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12AX7
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More blah (albeit off the original topic of electrolytic production).
Dissolved nearly 100g Cu wire in HCl acid, with the help of some Ca(OCl)2. Stoichiometry works out to around 140 grams Pool Shock (which is listed as
68% hypo; any idea what the remainder is?) and 300 grams 31.5% acid (no I don't have any graduated glassware). Instead of soaking wire in acid
and adding bleach (mmm, green bubbles..) a better apparatus would probably be a tall column with bleach on the bottom and dripping HCl over the copper
wire piled in tangles over the bleach, so that the acid causes Cl2 fumes which condense with the Cu directly, and the acid dripping down washes the
formed salt to the bottom.
After about an hour I had most of the hypochlorite added; it was slowing down and getting thicker (about right, CuCl2 isn't real stable in a
stoichiometrically neutral solution) so I added water and 100g acid. After adding all the hypochlorite and removing what was left (5-10g of sharp,
spiny, thinned copper wire) I crudely assessed the solution: dark yellow, brownish color; black in any thickness. Moderate dilution gives a dark
green color (copper tetrachloride ion, correct?) with no white precipitate,
further dissolution gives a light sky blue/cyan color (hydrated copper ion). Seems to be pretty pure divalent, at least.
But, it's cut about half with calcium chloride - any suggestions on seperating them? Could someone dredge up solubility vs. temperature for the
two salts for me? Quick google says they're about equally soluble at room temp...
Tim
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12AX7
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Just forwarding some photos tonight...
I removed the calcium in the previous solution by adding sulfuric acid. After adding a few pounds baking soda (in a five gallon bucket to handle the
bubbles  ) I had a suspension of pale green basic copper carbonate (unless
there's some chloride in there; the solution was green chloride complex at the time). After some settling and washing, this was obtained. Maybe
100-200g?
Tim
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Pyridinium
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I always made Cu2O by heating a copper pipe in the torch flame to form the black CuO, then quenching it by spraying water from a hose. It turns from
black to red very quickly, then flakes off.
It might make a difference that the torch was acetylene, which tends to impart carbon into the substrate. This might have had an effect; I'm
not sure it works as well if it's been done with propane.
I know, copper and acetylene, not a good pair.  but plumbers use it all the
time.
Edit: yah I know the post was first about CuO, but you could have that too if you don't quench the copper pipe.
[Edited on 24-6-2005 by Pyridinium]
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The_Davster
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I had a surplus copper nitrate solution a few weeks ago from a silver purification experiment, so I decided to turn it into copper(II) oxide. Sorry
about no exact numbers, but the experiment is not finicky at all and easily reproducable in my experiance.
The dilute copper nitrate solution(around 500mL) was reacted with 10mol/L sodium hydroxide solution until Cu(OH)2 stopped precipitating. I did not
wait for the ppt to settle, I just put the beaker on the hotplate on high(with a boiling stone) and stirred it vigorously. After a few minutes the
copper hydroxide began darkening a few minutes later the ppt was completly black and would settle very quickly upon ceasing stirring. The solution
was boiled with stirring for a few more min to ensure all the hydroxide had been converted to the oxide, then filtered(filtrate was completly clear)
and washed with plenty of distilled water. Air dryed then dried in a beaker on the hotplate, I was rewarded with a few grams of nice black CuO.
[Edited on 24-6-2005 by rogue chemist]
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12AX7
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Hmm odd. I've been making Cu(OH)2 by electrolysis in NaSO4 solution, it's turning out just fine; what I don't get is, I boiled the
suspension today and it didn't change color? It's the same drab teal color. (Hmm odd, I wonder where the green came from. No Cl- that
I'm aware of..)
Tim
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Lambda
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Maybe a little Cu+ in your Cu++ solution Tim.
[Edited on 10-7-2005 by Lambda]
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12AX7
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Undoubtedly; there is some orange on the copper anodes. But that would turn it reddish, not green. CuOH isn't stable and autolyzes to Cu2O
(orange), produced by NaCl electrolysis. By and large, the stuff is blue to green Cu(OH)2, so boiling should've caused a change.
Tim
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chemoleo
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Look, 12AX, it's not like you have to disprove all of science by observing odditites.
Instead, try to explain it yourself.
For instance, did you take your precipitate, and filter it, and wash it free from left over NaOH and so on?
How does it behave then? What happens if you dissolve it in a weak acid such as HAc?
I have done this very experiment myself, and it works fine (even though the CuO produced is not as active as commerical CuO in the context of a
thermite), the solution goes from turquoise to greenish to olive green to black - eventually.
Never Stop to Begin, and Never Begin to Stop...
Tolerance is good. But not with the intolerant! (Wilhelm Busch)
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Lambda
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It can allso be contaminants like Nickel in your Copper electrodes. It is not uncommon that melted recycled Copper has Nickel, Iron....and a very long
list of contaminants.
Are you using the same Copper electrode material as your previouse one ?. I mean a cutoff section.
[Edited on 10-7-2005 by Lambda]
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The_Davster
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Really Chemoleo, your home made CuO is less active in a thermite than commercial? Strange, mine always produces a thermite that seems to burn faster
than blackpowder.
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chemoleo
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Yes. I think I mentioned that over in the exotic thermites thread. I don't know why; I even I roasted it at red heat.
Regardless, the conversion of Cu(OH)2 to CuO seemed to work very well.
But yes, commercial CuO/Al is much faster than commerical BP, almost instantaneous. Did you honestly get this with homemade CuO and Al? I am baffled.
PS yet another scientific oddity eh?
[Edited on 11-7-2005 by chemoleo]
Never Stop to Begin, and Never Begin to Stop...
Tolerance is good. But not with the intolerant! (Wilhelm Busch)
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The_Davster
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Yeah, I honestly got thermite that fast with my homemade CuO. I only dried it at room temperature and then on a alcohol burner for 10 min or so,
perhaps your heating it to red heat reduced some of the CuO to Cu? My CuO used in a thermite test was CuO produced from CuSO4 and NaOH, I should test
my CuO made from Cu(NO3)2 +NaOH in a thermite, just to see if it makes any difference.
EDIT: I should probally mention my Al was 400 mesh spherical. Anyway we should probally take this over to the exotic thermite thread if we go more
in-depth into thermites.
[Edited on 11-7-2005 by rogue chemist]
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12AX7
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Hm, it turns drab green then finally black? Maybe I didn't boil it long enough... Is this a half hour thing or a ten seconds thing?
It's not filtered yet, still saturated with NaSO4. I don't see why any NaOH would be present, I mixed it well after removing the electrodes
and power.
I had the lye out this evening so I dropped some of the gloop into a jar and added some NaOH and water. In a few minutes of settling, I got the deep
blue 2Na<SUP>+</SUP> CuO<SUB>2</SUB><SUP>-2</SUP> solution.
If nothing else, I can wash, dry and calcine the Cu(OH)2, but since it isn't much in the habit of settling, I'd rather boil it down a bit
and save the Na2SO4 solution.
My CuO burns nice and fast with MgAl, though not flash powder fwoomplike; I'd guess my (and everyone else's) results are due to large and/or
irregular grain size. It does produce lots of copper metal vapor though!
Tim
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The_Davster
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It is usually a 20min thing. A good indication your Cu(OH)2 has been boiled long enough to convert it all to CuO is that the black ppt will settle
very fast. Unlike a Cu(OH)2 ppt whick takes forever to settle.
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12AX7
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Alrighty. Then, I'll need something better than the soda-lime glass jar I've been nuking it in, then set it on the gas burner for a lil
while...
Edit: I've been boiling it lightly for the last two hours, and...somehow... it's still colored. Now a drab dark gray greenish thing.
I tossed in a little alkali to try to hydrolyze it better, it does make a bit of black on contact but the whole thing has yet to turn. If it is
oxychloride, I have no damned idea how it got in there. (Maybe a trace from the porous graphite cathode but not nearly enough to chloridate all the
sludge present.)
Tim
[Edited on 7-11-2005 by 12AX7]
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kine
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I started the thread for the CuO formation but i have done some experiments lately with copper and i have a question:
I left a copper rod in 35% HCl for a day to form copper chloride. The copper metal that was in the acid disolved and turned the liquid to black color.
The top of the rod is filed with green powder wich is??? I am not sure as i am new to chemistry.
Is it Copper carbonate? Is it copper chloride (metal outside of the acid)? Is it copper hydroxide?
Only thing i am sure is that it is a litle bit caustic to the skin and iritating.
Also if it is copper cloride, is it dangerous on direct contact with the skin? I washed hands repeatedly but isome remains on..
thanks.
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neutrino
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Was the solution still acidic when you found this precipitate?
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12AX7
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| Quote: | Originally posted by kine
I started the thread for the CuO formation but i have done some experiments lately with copper and i have a question:
I left a copper rod in 35% HCl for a day to form copper chloride. The copper metal that was in the acid disolved and turned the liquid to black color.
The top of the rod is filed with green powder wich is??? |
Most likely copper hydroxychloride (written Cu(OH,Cl)2 ).
| Quote: | | Is it Copper carbonate? |
Unlikely above an acidic solution. I'm sure you've noticed the smell of the acid.
| Quote: | | Is it copper chloride (metal outside of the acid)? |
Unlikely, being deliquescent it would dissolve from ambient humidity (unless you left it in a warm, dry place) or dissolve in the solution by capilary
action.
| Quote: | | Is it copper hydroxide? |
No, Cu(OH)2 is blue and cannot be made from a chloride solution.
| Quote: | Only thing i am sure is that it is a litle bit caustic to the skin and iritating.
Also if it is copper cloride, is it dangerous on direct contact with the skin? |
http://www.google.com/search?q=cupric+chloride+msds
Before combining chemicals, investigate all possible reactions and check the safety hazards of each chemical involved and produced.
Tim
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The_Davster
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For everyone who did an electrolysis of MgSO4/Na2SO4 with copper electrodes and got a precipitate of some light blue compound which *did not*
decompose upon heating into CuO, I think I figured out why.(by figured out I mean found an old post which could explain it)
| Quote: |
Originally posted by Organikum
And after my quick and dirty experiments the hard truth from the MERCK-index:
- copper sulfate dibasic: Cu3H4O8S, blue-green, rhombic, bipyramidal crystals. Practically insoluble in water.
- copper sulfate tribasic: Cu4H6O10S, very finely divided, light-blue, gelatinous particles. Practically insoluble in water.
- Copper hydrate: Cu(OH)2, blue to blue-green gel or light blue crystalline powder. Stability is dependent on the method of preparation, may decompose
to black CuO on standing a few days or upon heating. Practically insoluble in water. Sol. in concentrated alkali when freshly precipitated.
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So I imagine that before electrolysis produces Cu(OH)2, these mixed basic salts form, and being insoluble, precipitate out, preventing conversion to
the pure hydroxide.
[Edited on 16-12-2005 by rogue chemist]
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Darkblade48
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The odd thing is that I carried out electrolysis in a MgSO4 solution with copper electrodes and got lots of fluffy blue precipitate. After I heated
it, in hot boiling water, it eventually turned black. But it did take a while, so perhaps you need to boil/heat your solution for a bit longer?
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The_Davster
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I tried to boil it for a while, after an hour not much had happened. I tried it with a KNO3 electrolyte today while boiling the solution, CuO was
produced nice an quick. So for anyone who wants CuO, KNO3 electrolyte is definatly the way to go.
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12AX7
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Na2SO4 works well, too. The key being heat that decomposes Cu(OH)2 > CuO + H2O.
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The_Davster
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I have not tried Na2SO4, but I would have thought there would be problems with the sulfate creating the annoying basic copper sulfate salt, just like
MgSO4 does?
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