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Ramiel
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I don't know why I'm doing this...
Jus' a littlé searcha later, and we have at our greedy little fingertips;
http://mineral.galleries.com/minerals/oxides/cuprite/cuprite...
That's not very helpful, but it gives a little interesting information...
the page
http://www.wm-blythe.co.uk/WMBLYTHE/CSDS.nsf/7eb5bed35d8ca35...
reveals that Cupuric hydroxide has the "Colour: Blue/green"
we even have a .edu site regarding the REACTIONS of copper...
http://dwb.unl.edu/Chemistry/MicroScale/MScale04.html
so, here I've demonstrated how useful google is - just 2 minutes yeilded all this bountiful information. Don't apologise - just don't
do it again. 
Sincerely
-Ramiel
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Organikum
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Actually I used Google and found lot´s more of information which indicated that this thread contains some strong but wrong opinions on the matter.
"Use Google" is no argument as the board can be closed following this logic.
I hoped our brandnew moderator would contribute some information to solve the contradiction in this thread.
[Edited on 23-6-2003 by vulture]
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vulture
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Instead of usually stirring up the hornest nest myself, it seems that I landed in an already stirred up one.
Now honestly, I can't see what's wrong with organikum asking these questions.
Sites to refer to are nice, but the reactions are sometimes ambigous and in this case it's not very clear either.
And ofcourse, an explanation of someone likeminded usually makes things alot more understandable than a dry internet text.
And yes, H2O2 can act both as an oxidizer and a reducer. With KMnO4 it acts as a reducer for example.
BTW, google is turning up more and more crap and sponsored results. I prefer www.ixquick.com, mostly yields more relevant results.
[Edited on 23-6-2003 by vulture]
One shouldn't accept or resort to the mutilation of science to appease the mentally impaired.
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Organikum
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thanks
Ramiel: Peace!
You might believe me that I usally inform myself before I post:
Firsthand so possible 
phantastic blue this blueish blue, isn´t it?
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Organikum
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I want to add that in the bottle is seen the copper(II) ion ammonia complex as there is quite a lot of ammonia present. Should be as ammoniawater was
added to oxidised copperwire as seen.
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chemoleo
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interesting... when I mixed CuSO4 and NaOH, I got the very pretty sky blue Cu(OH)2. After filtering this many times ( to remove the Na2SO4) - which
is very tedious, i boiled the hell out of this... with the result that I got a dark green/olive coloured precipitate. After drying this, I flamed it
under the bunsen - and it turned black, and very fluid (like activated charcoal). Believe me, I heated it strongly so that it started to glow! Anyway,
the point was to use this for a thermite mixture (as I was told this is one of the most powerful ones), but to my dismay, the reaction was by no means
more vigorous than with Fe2O3. Comments?
By the way, whats the chemistry behind the H2O2 precipitiation method, i.e. whats the mechanism to cause CuO to form?
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chloric1
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Electrolytic cupric oxide
I have abook on how to make your own lab. One chemical proceedure states on how to make basic lead carbonate by electrolyzing a sodium carbonate
solution with a little sodium chlorate dissolved. Now lead corrodes electrolytically slowly so i would think you could use a copper anode instead and
use NaOH instead of the carbonate. Then add about 5 or 6 gramms of sodium chlorate and hook up your voltage and go to town!
An alternative is something I discovered by accident about 14 years ago, mix 3% H2O2 with clorox bleach until the fizzing stops and then insert a
copper wire and go to bed. when you wake up the next day you will have copper oxide and salt water.
Fellow molecular manipulator
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Ramiel
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CuO?
I just got a great amount of copper oxide from the following method:
A large amount of (blue) Copper Sulphate was dehydrated over a bunsen for an hour - forming 100g of the mostly anhydrous Cupric Sulphate (white).
To 100g of ah. Copper Sulphate was added 400mL of distilled water. Not all the CuSO<sub>4</sub> had dissolved before 50g of solid NaOH was
added in equal parts of ten (the reaction is quite vigorous).
A dark black colloid immediately formed which was found to be Cupric Oxide (I just filtered a tiny bit, dried it and burnt it in a bunsen flame).
Now, I'll admit that filtration will be a bitch without serious vacuum apparatus or similar, but it is the price one pays for using only two
chemicals for the reaction - two easily acquirable chemicals too!
Play safe & always brush your teeth
-Ramiel
ps. sorry for being a bit aggro, but it wasn't @ you, Organikum.
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unionised
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A couple of points,
1 Why carefully remove the water, then add water?
2 It seems to me that the reaction of copper with hot air to form CuO is relevant here.
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Mr. Wizard
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washing soda alternate
Couldn't sodium carbonate, washing soda, Ph-Up pool chemical, be used instead of expensive NaOH? You would get the copper carbonate, which could
be roasted to get the oxide, IIRC. Couldn't it also be used instead of NaOH in making Na from Aluminum? Maybe this should be on another thread
but Na2CO3 is a lot cheaper and more available than NaOH. In many cases Na2CO3 plus cheap slaked Lime in water will give a solution of NaOH with CaCO3
precipitate. My old chemistry set book used it in about every reaction
Just a thought.
[Edited on 14-12-2003 by Mr. Wizard]
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Ramiel
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To unionized
I noted in some of the posts above that chemoleo had to go to a lot of bother in seperating the hydroxide of copper and the sodium sulphate salt. : |
So, I thought "If I dump large chunks of anhydrous cupuric sulphate in water, and then let the NaOH go nuts on it, I think the Copper (II) Oxide
will form in vitro as it were"
and it worked
No messing about with repeated filtrations, boiling or blasting... just literally dump two powders together, and then filter - viola, as they say.
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unionised
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Erm, Did I misunderstand the bit about filtration being a bitch?
Anyway, why not use hydrated CuSO4 rather than waste effort drying it then adding the water back? Even fewer steps.
BTW, Na2CO3 would work, the hydroxide has the advantage that it will decompose to the oxide on boiling in water I'm not sure the carbonate will
do this.
If you do this slowly there is some hope of getting larger particles of CuO which will be easier to filter.
[Edited on 14-12-2003 by unionised]
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KABOOOM(pyrojustforfun)
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unionised: I <b>deeply doubt</b> (because of his Ramielness ) that
he didn't know the amount of crystallization water (which is 5) so he decided to dehydrate it to measure the exact required amount
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12AX7
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Electrolytic copper oxide
Okay so I want some copper and lead oxides... burning is useless at best (especially since the slightest contact with a gas flame's H and C
content (even lean mixture) will reduce CuO!), and impossible to get all the metal oxidized. I don't have any nitric acid (or nitrates) and
HCl/H2SO4 aren't powerful enough, so that leaves electrolysis. This morning I hooked up some sodium carbonate electrolyte to a power source
(unknown voltage and current, just some turns of 8AWG on a spare MOT with a diode
) and "anodized" some copper wire with it (using a graphite cathode).
Now, oxygen and hydrogen appear to be coming from the respective electrodes..... but why is the solution turning blue? I mean, yes copper, but
perfectly clear, not precipitating? There ought to be plenty of carbonate ions left in solution (boiled down bicarbonate), and last I checked, copper
carbonate is rather insoluble. WTF?
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chemoleo
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12AX7 - SEARCH! It's not the first time I have to be your personal thread merging service!!
Your mood 'lazy' seems an accurate assessment!
[Edited on 4-4-2005 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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12AX7
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DOH! I had searched, guess I was too specific.
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12AX7
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...So, I left the soda cell running (between copper electrodes) for a while and it seems to have plated copper sponge. WTF?
The solution is still blue... how can copper ions be stable in a basic solution?! It ain't cuperate...
Tim
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uber luminal
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I just went into the lab next door and made this work.
plastic dish, 120 ml water. 3 grams NaOH crystal form.
copper anode, copper cathode (why not? i already have sheets of copper)
open any chemistry book and find the reduction potential for Cu(2+) + 2e --> Cu (s) == +0.337 V. now find it for all the spieces in soln. looks
like Cu and water are the highest. Great. add enough volts over waters reduction potential and away we go.
I walked away for about 10 min and came back to find... black precipitation around the anode.
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Ramiel
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Thank you uberluminal. What kind of setup did you have? is the current density important, I've had a lot of success with my electrodes at right
angles to each other in the past, as suggested by <html><a
href="http://67.15.145.24/~sciencem/talk/viewthread.php?tid=621">Chris Owen</a></html>
Uh, in response to earlier posts, my chemistry teachers (three different teachers) have quoted three different crystalisation structures of these
sorts of ions... 4, 5 and 6 all seem to be 'the truth'. Just thought that was interesting.
- D
ps. now Orgi, there's no need to yell at people to use the search function, it only gets them angry.. 
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uber luminal
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what kind of setup? erm... a lab power supply i picked up from surplus, allowing you you set voltage and amperage. I would strongly suggest you find a
cheap power supl on ebay or at a university sale or soemthing. Using just the ... 12v dc power supply limits your ability to do specific reductions.
I clamped the copper pieces in with SS clips, and clamped the clips with more clips to the side of the plastic dish. I also have a binder full of
reduction potentials from CRC online handy.
I would like to add that cheap acid will also oxidize the copper with a bit more ease...
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12AX7
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Followup for CuO-ism...
Been electrolyzing Cu wire in salt water, seems to work well, what I don't get is why it doesn't make CuCl(2) solution but rather a loud
orange precipitate. Another thread says it's hydroxide, but ain't no copper hydroxide that's anything but green or blue...
Tim
[Edited on 26-4-2005 by 12AX7]
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chemoleo
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Try to electrolyse it in HCl/NaCl. You need surplus HCl.
Then, try to dissolve the 'loud orange precipitate' in HCl. Does it dissolve? or does it just turn white, and green with time? If so, the
orange precipitate is Cu2O, the white stuff CuCl, and the green stuff CuCl2 * nH2O.
Edit: Just reading of your troubles from above: You want CuO, but are unable to find CuSO4 anywhere?
Try dissolving the Cu in vinegar (acetic acid concentrate) and H2O2, and put it on a radiator. This should do the job.
[Edited on 26-4-2005 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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12AX7
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"Acetic acid concentrate"
I thought vinegar was only 5% or so? (Or is that pH, I forget.) Seems to me it'd take a lot of vinegar to dissolve a pound of Cu metal!
I took a spot of the orange ppt and added some reagents, HCl turns it into a clear green liquid, CuCl(2) (whichever one it is), so it's probably
an oxide. Cu2O doesn't make sense to me as there should be only chlorine at the anode, or does it form an unstable hypochlorite on contact which
then redoxes to Cu2O? Also, the streak for cuprite (natural Cu2O) is listed as reddish brown, uh rusty I suppose, but this really is a bit brighter.
Could be though.
I haven't tried a mix of acid and salt but straight acid only plates out a nice mossy deposit as soon as the solution gets enough CuCl2 in
solution to do so.
(It does however strip nickels nicely, leaving emerald NiCl2 in solution.
I'm currently plating that solution out, with an aluminum anode, of all things...)
Tim
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chemoleo
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Yes, I found 20% concentrate in many shops around here. You just need to go searching. Alternatively, make the Na salt with Na2CO3 (baking powder  ), and react this with 30% HCl, and distill the acetic acid azeotrope. It's all
possible if one tries hard enough!
Also, vinegar is cheap. And a pound of Cu is really not that much, because of its high atomic mass you need not thaaat many moles of acetic acid. Do
the calculations.
It probably is Cu2O. It can have all shades of dark orange to nearly red brown. Think of it that way: NaCl is electrolysed, producing NaOH which in
turn could react with nascent CuCl(formed at the anode) to form CuOH, which immediately reacts onwards to Cu2O. That's why I suggested an excess
of HCl, to make sure that the free OH is scavenged (which it is not in your system). It's messy anyway. I'd rather use acetic acid.
Cu2O reacts with HCl to form white CuCl, which, if the solution is sufficiently oxygentated, forms immediately green CuCl2. Aside from that, you are
getting a green solution...would you like to tell me what else it could be?
Never Stop to Begin, and Never Begin to Stop...
Tolerance is good. But not with the intolerant! (Wilhelm Busch)
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12AX7
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Hmm, odd...(apologies for the long drawn out post, I'm using this time to reflect on it as well...)
Yesterday, I was anodizing a copper ingot. Its story: last year I ingotted my copper wire (about 10 pounds), so it has to be at least 98% pure. The
bars are full of bubbles, from absorbed oxygen released on freezing, testament to this.
Well it was proceeding slowly, mostly due to a buildup of something on it. After a short time, it's mostly orange (like the rest of the
suspension was at the time, the Cu2O I presume); after a longer time, a white layer formed - CuCl I guess, and after a long time without disturbance,
the surface became black (CuO?), totaling a thick layer of these three materials.
What I don't get is, why was it building up, unlike the straight copper wire before? Low current density not keeping it clean or something?
Looked like a pretty good current though.
The behavior of the suspension after was kind of weird too, although now that I reflect on it, it makes a little sense. That which doesn't...
I decanted and washed the orange suspension a bit, then I added some muriatic acid which produced a range of colors on contact, finally yielding a
large amount of white precipitate, I'm guessing 2HCl + Cu2O = H2O + Cu2Cl2, and a greenish solution.
Thinking PbCl2 (how it could get in there, I have no idea... one is I don't clean my crucible well ), I decanted it once again and attempted to wash the white stuff - now I've got a solution that's turning
orange again! DOH! But that would suggest that CuCl is only stable at lower pH; wouldn't a lower pH oxidize it further to CuCl2???
Anyway, I took a portion of the white suspension and added some lye. If it were lead, it should turn brown on contact (plumbate, no?) and eventually
clarify to a white lead hydroxide (something I've performed before on my PbCl2 yield). Instead it gave a brown precipitate (iron??) and a blue
solution (okay, so I got some ammonium contamination in there, somewhere). A bit darker than the blue I saw from Cu + Na2CO3 electrolysis, but still
reminiscent of a basic copper solution. No I don't have any personal experience with controlled demonstration of tetrammonium ion and its exact
color.
So I took yet another portion and, also suspecting lead, added sodium sulfate - surely a definitive test. No precipitate has formed yet; in fact,
it's turned orange again, and is now giving off nonflammable (probably CO2) gas! WTF, I didn't put anything gaseous in there!
Anyway...that's my story...feel free to......well no, piss on yourself if you feel the urge to piss. Anything else constructive, feel free to
post.
Edit: I think I'll add more HCl to that white ppt and see if it needs extra low pH to progress from Cu2O > Cu2Cl2 > 2CuCl2. I'd like
to know what's going on with that sulfate solution though.
Tim
[Edited on 28-4-2005 by 12AX7]
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