Sciencemadness Discussion Board - Can't make CuO

Can't make CuO

bonemachine - 30-10-2002 at 05:55

I have trouble making CuO with electrolysis. Can anyone describe a method?

CUO

the_describer - 30-10-2002 at 12:10

Can you tell me the sense to make CuO by
electrolysis !

bonemachine - 30-10-2002 at 13:55

oxigen produced from cathode will combine with Cu. That's what i thought but it doesn't work.... Now can you help?

madscientist - 30-10-2002 at 15:55

You'll get the blue insoluble Cu(OH)₂. Filter it out, and heat it strongly, to get the black CuO.

bonemachine - 31-10-2002 at 01:26

Is it dangerous to heat it or to handle it with bare hands? Both Cu(OH)2 and Cu?

bonemachine - 31-10-2002 at 03:23

While electolising water with Cu electrodes it turned blue meaning that Cu(OH)2 is present. I added a small amound of H2O2 to spped up the formulation of Cu(OH) and the solution instantly turned dark green to black.

Also i want to ask if it is posible to exist Cu2FeO4? And if it could then electrolysis could work?

vulture - 31-10-2002 at 05:05

I added a small amound of H2O2 to spped up the formulation of Cu(OH)

I'm sorry but that's impossible. From Cu(OH)2 to CuOH is a reduction not an oxidation, thus this is not possible with H2O2.

bonemachine - 31-10-2002 at 06:22

Anyway my finished product is dark green and i have filtered it out. Looks like rust but the electrodes where Cu and the water had a little h2o2. What is that green stuff CuO?

Marvin - 31-10-2002 at 11:08

Why not just add sodium hydroxide solution to copper sulphate solution, and heat the precipitated hydroxide until it goes completely black?

What are you after the CuO for?

bonemachine - 31-10-2002 at 11:49

because i don't have those materials and i prefer electrolysis.

bonemachine - 31-10-2002 at 11:57

Also i forgot.... Could someone tell me the exact color of CuO? I mean it is totaly black or dark green? I am realy not sure why the blue color turned dark green from the adition of h2o2 and if it is CuO.

madscientist - 31-10-2002 at 14:39

CuO is black.

bonemachine - 1-11-2002 at 06:29

Then what could be that green stuff?

raistlin - 12-11-2002 at 17:20

No offense, but I dont think I have ever read a thread on this site that seems as kewlish as this... If you dont know what something is, dont ask stupid questions on here, go look it up on google or dogpile or on whatever! Do sciencemadness a favor and look things up before you talk.

bonemachine - 13-11-2002 at 04:45

I found nothing searching and i ask again what is this green stuff?

BASF - 13-11-2002 at 11:07

you want to get pure CuO, right?

Make a solution of CuSO4, then add a mix of NaOH and H2O2.
Black CuO will precipitate.
-That´s a method out of wet chemical analysis, if i remember right...

bonemachine - 14-11-2002 at 02:02

Ok i tried this method and it works fine thanks!!

raistlin - 14-11-2002 at 17:00

It could be copper chloride, but I dont know.

bonemachine - 15-11-2002 at 03:17

The strange is that it was sky blue at the begining and when i aded 3% H2O2 it turned totaly green.

trinitrotoluene - 24-12-2002 at 21:34

I did the same experenment and I had some pretty poor results. What I got was only a thin layer of CuO. I did work but it was a very slow process. I say 20 times slower then electrolysis of iron to make Fe2O3.

rikkitikkitavi - 25-12-2002 at 01:24

you could try to electrolyse with Cu-anode in a NaCl solution.

A red preciptate of Cu2O will be formed

anode Cu=> Cu+
Cathode 2 H2O => 2OH- + H2

Cu + 2 Cl- => [CuCl2]- (complex)
but this complex reacts with OH-

2[CuCl2]- + 2 OH- => Cu2O + H2O + 4 Cl-

And I belive that heating Cu2O in air to about 600-700 C will form CuO since this is more stable.

/rickard

madscientist - 25-12-2002 at 22:23

I've noticed that electrolyzing a solution of MgSO₄ with a Cu anode, and then allowing it to sit for some time, precipitates Cu(OH)₂.

Another use of CuO

Theoretic - 18-6-2003 at 07:29

You could use CuO for thermite mixtures, as well as CuCl2 and CuSO4.
You can get CuO by electrooxidation of copper in a salt solution (the stronger the better) and then heat the green stuff to get CuO.

vulture - 18-6-2003 at 08:22

Read topics before posting, this has been said before!

Organikum - 22-6-2003 at 09:56

ehem, some humble questions....

from copper sulfate:
- NaOH precipitates Cu(OH)2 ?
- NaOH + H2O2 precipitates CuO ?

but Cu(OH)2 + H2O2 cannot give CuO?

btw. H2O2 is told to be doublefaced - reducing or oxidizing both is possible.

Ramiel - 22-6-2003 at 19:34

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

Organikum - 23-6-2003 at 08:24

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]

vulture - 23-6-2003 at 08:35

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]

thanks

Organikum - 23-6-2003 at 08:56

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?

Organikum - 23-6-2003 at 14:11

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.

chemoleo - 10-9-2003 at 19:04

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?

Electrolytic cupric oxide

chloric1 - 11-10-2003 at 17:09

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! :cool:

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.

CuO?

Ramiel - 13-12-2003 at 06:28

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 CuSO4 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.

unionised - 13-12-2003 at 14:11

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.

washing soda alternate

Mr. Wizard - 13-12-2003 at 17:15

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]

To unionized

Ramiel - 13-12-2003 at 20:37

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.

unionised - 14-12-2003 at 04:25

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]

KABOOOM(pyrojustforfun) - 22-12-2003 at 19:54

unionised: I deeply doubt (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

Electrolytic copper oxide

12AX7 - 4-4-2005 at 03:46

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?

chemoleo - 4-4-2005 at 04:32

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]

12AX7 - 4-4-2005 at 04:52

DOH! I had searched, guess I was too specific.

12AX7 - 6-4-2005 at 22:13

...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

uber luminal - 7-4-2005 at 19:12

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.

Ramiel - 7-4-2005 at 21:09

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..

uber luminal - 8-4-2005 at 10:07

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...

12AX7 - 26-4-2005 at 10:02

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]

chemoleo - 26-4-2005 at 11:14

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]

12AX7 - 26-4-2005 at 11:36

"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

chemoleo - 26-4-2005 at 12:11

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?

12AX7 - 28-4-2005 at 11:53

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]

12AX7 - 12-5-2005 at 14:01

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]

12AX7 - 21-5-2005 at 03:46

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.)

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

12AX7 - 22-5-2005 at 13:39

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

12AX7 - 23-6-2005 at 23:47

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

CopperCarbonate.jpg - 17kB

Pyridinium - 24-6-2005 at 11:28

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]

The_Davster - 24-6-2005 at 11:55

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]

12AX7 - 10-7-2005 at 13:53

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

Lambda - 10-7-2005 at 14:22

Maybe a little Cu+ in your Cu++ solution Tim.

[Edited on 10-7-2005 by Lambda]

12AX7 - 10-7-2005 at 14:56

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

chemoleo - 10-7-2005 at 15:32

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.

Lambda - 10-7-2005 at 15:38

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]

The_Davster - 10-7-2005 at 15:53

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.

chemoleo - 10-7-2005 at 16:03

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]

The_Davster - 10-7-2005 at 16:21

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]

12AX7 - 10-7-2005 at 17:21

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+ CuO2-2 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

The_Davster - 10-7-2005 at 18:05

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.

12AX7 - 10-7-2005 at 18:19

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]

kine - 15-12-2005 at 09:33

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.

neutrino - 15-12-2005 at 09:50

Was the solution still acidic when you found this precipitate?

12AX7 - 15-12-2005 at 11:24

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

The_Davster - 15-12-2005 at 20:46

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.

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. :mad:

[Edited on 16-12-2005 by rogue chemist]

Darkblade48 - 15-12-2005 at 23:51

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?

The_Davster - 18-12-2005 at 22:31

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.

12AX7 - 19-12-2005 at 00:34

Na2SO4 works well, too. The key being heat that decomposes Cu(OH)2 > CuO + H2O.

The_Davster - 19-12-2005 at 00:42

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?

12AX7 - 19-12-2005 at 12:01

If it does, and the salt doesn't break down on boiling, then nevermind.

It should work better than MgSO4 at the very least because Mg(OH)2 will want to plate on the cathode...

Tim

The_Davster - 19-12-2005 at 17:35

Well I did a longer bigger test of producing CuO in KNO3 solution. About 9 copper wires 4" long were suspended over a 600mL beaker filled with KNO3 solution as the anodes. The same number of the same lenght of copper wires were used as the cathodes. 12V(max 7A) was applied for aproximatly 4h-5h.
for the first 2h, Cu(OH)2 was flaking off the anode, but around the fourth hour, the cell had heated up enough to be producing CuO right away. Around the 4.5h mark, ammonia(or at least what smelled to be ammonia) began to be given off

. I must have reduced nitrate to ammonia somehow. Seeing as this was running indoors, I turned off the power. Interesting. Perhaps if the temp was kept lower, ammonia formation could be avoided.

EDIT: The CuO produced this way is not as pure as CuO produced the usual way of NaOH + CuSO4 while heating, there was a bit of Cu2O and Cu particles in there as well, Cu likely from electrode corrosion.

[Edited on 20-12-2005 by rogue chemist]

darkflame89 - 19-12-2005 at 23:32

The ammonia would come from the cathode side. At the temp. that your cell encountered during electrolysis, nascent hydrogen produced from the cathode would reduce nitrates to ammonia. I only wonder if the electrolysis setup was run long enough, might the copper(II) oxide redissolve to form the ammonia complex. Provided that the setup is stirred that is.

Magpie - 25-5-2008 at 20:01

Although my current work with CuO is not electrolytic I thought this a good thread in which to continue a discussion of its preparation.

My original goal was to prepare cupric acetate monohydrate, Cu(OAC)2*H2O. However, as often occurs when you don't have a proven method, problems popped up along the way. My approach was to use CuO as the precursor for the cupric acetate. My preparation of CuO would be with the readily available root killer, CuSO4*5H2O.

Experiment 1
CuSO4*5H2O was dissolved in deionized water and warmed to around 50C. A stoichiometric amount of NaOH crystals was then added. This first formed the light blue ppt of Cu(OH)2 which upon heating to near boiling produced very dark chocolate brown suspended solids, which I took as CuO. This was Buchner filtered, washed, dried, and weighed. The weight indicated a CuO yield over 100%. Upon reading some reference for making verdigris I suspected sulfate contamination of my product. Dissolution of a small sample in 6N HCl and adding a few drops of BaCl2 confirmed the presence of sulfate.

Experiment 2
Several references specified the addition of NH4OH until the initial light blue Cu(OH)2 turns to a deep blue from the complex [Cu(NH3)4]++, prior to the addition of the stoichiometric amount of NaOH. By doing this an almost perfectly black ppt of CuO was formed and the test for sulfate was negative. Also, the yield was just slightly less than theoretical 100%.

I wanted to share these results and give enough detail so that those more versed in copper chemistry could comment. Can anyone tell me exactly what chemistry is taking place here? I.e., why did I get contaminated CuO wo/ammonia addition and apparently good CuO w/ammonia addition?

12AX7 - 25-5-2008 at 20:35

Hmm... ya know, I have:
Cu(2+) + SO4(2-) <---> CuSO4(aq) pKf = -2.36

So maybe you have a 0.1~1% sulfate impurity complexed in. CuSO4 is soluble, so I don't really know how the complex would precipitate. If it is a complex, displacing it with a stronger ligand (NH3) would get rid of it, but that begs the question, how does CuO precipitate from that solution? I thought Cu(NH3)4(OH)2 was soluble. Does the solution remain blue at all?

Heating it to 500-800C should cook off most anything anyway (including CuCl2, incidentially), although some CuO may decompose to Cu2O, which is more stable at very high temperatures (by a hundred C, near the melting point).

Tim

Magpie - 25-5-2008 at 21:36

Quote:

Does the solution remain blue at all?

After addition of the NaOH the solution goes to light blue as is Cu(OH)2. Then as this is heated the suspension goes black as CuO is formed. When Buchner filtered the filtrate shows some clear blue characteristic of [Cu(NH3)4]++. This is a source of a slight loss of Cu, but not much.

Quote:

Heating it to 500-800C should cook off most anything anyway

I wanted to avoid calcination as a reference says that this makes the CuO insoluble except in boiling strong acid. It would be interesting to calcine a sample to determine purity, however.

[Edited on 25-5-2008 by Magpie]

[Edited on 25-5-2008 by Magpie]

[Edited on 25-5-2008 by Magpie]

12AX7 - 26-5-2008 at 00:11

So only a small amount of ammonia is needed, enough to produce a catalytic amount of Cu(NH3)4 you might say?

Hmm, that implies Cu(NH3)4 (or some n-ammine species) coexisting with Cu(2+) or CuOH+ species, but not Cu(OH)2 as that would precipitate. So I guess the question is, what pH can it reach, without precipitating, but while retaining some NH3 in solution? I guess it can't be too horrible, I have CuOH+ <---> Cu(OH)2(aq) pKa = 8.7, not far from NH3. It should suffice to add any ammonium salt before adding NaOH, no?

Tim

ShadowWarrior4444 - 26-5-2008 at 09:22

Preparing CuO by electrolytic processes is as simple as electrolyzing a copper metal anode using an NaOH electrolyte. I have recently tested this as part of another project--when given 13.8v at 25A the anode will be quickly eroded with the evolution of a black insoluble powder.

>99% pure copper can be had in the form of uncoated high-gauge copper wire at most local hardware stores (in the electrical wire section.) Copper nails are also sold, however they are likely an alloy.

[Edited on 5-26-2008 by ShadowWarrior4444]

12AX7 - 26-5-2008 at 21:11

What strength NaOH? Cu(2+) is soluble in alkali.

Electrolysis of a Na2SO4 solution with copper anode gives Cu(OH)2, which decomposes on heating to CuO, especially convienient if the cell is run hot. NaCl complexes Cu(I) giving Cu2O which would need to be roasted to CuO. A nitrate or chlorate solution probably serves as well, but the cathode would need to be protected to prevent reduction. Perchlorate probably works as well as sulfate (say, I have a bunch of sodium perchlorate, I can easily try that).

Tim

ShadowWarrior4444 - 27-5-2008 at 13:01

Quote:

Originally posted by 12AX7
What strength NaOH? Cu(2+) is soluble in alkali.

Electrolysis of a Na2SO4 solution with copper anode gives Cu(OH)2, which decomposes on heating to CuO, especially convienient if the cell is run hot. NaCl complexes Cu(I) giving Cu2O which would need to be roasted to CuO. A nitrate or chlorate solution probably serves as well, but the cathode would need to be protected to prevent reduction. Perchlorate probably works as well as sulfate (say, I have a bunch of sodium perchlorate, I can easily try that).

Tim

The NaOH concentration need not be very high. It would simply be a matter of powering up the cell, then adding small portions of electrolyte until a black precipitate is visible from the anode.

My favorite method for producing copper hydroxide is to use an MgSO4 (Epsom salt) electrolyte; though, it should be noted that running the cell hot will not appreciably decompose the hydroxide. Dry copper hydroxide will need to be roasted in air at 185C to produce CuO, and while wet copper hydroxide will decompose over time, it will not be satisfactorily fast nor will it provide a relatively pure product.