Sciencemadness Discussion Board - Cu(OH)2 reaction

Cu(OH)2 reaction

guy - 27-10-2005 at 14:08

I am trying to produce lots of copper hydroxide. I have tried electrolysis but that is quite slow. I am attempting another reaction that should produce Cu(OH)2 but its not. It is

Cu + 2OH- --> Cu(OH)2 + 2e- E=+0.36
2e- + ClO- + 2H2O --> Cl- + 2OH- E=+0.89
------------------------------
Cu + ClO- + H2O --> Cu(OH)2 + Cl- E= +1.25V

This reaction should work right?

neutrino - 27-10-2005 at 17:11

I wouldn't expect that reaction to go too far; both the reactant and product are solid.

Do you have access to any copper salts? These would be a much better starting point than copper metal.

Darkblade48 - 27-10-2005 at 17:12

Quote:

Originally posted by neutrino
Do you have access to any copper salts? These would be a much better starting point than copper metal.

Dissolve any copper salt into solution and add a solution of (dilute) NaOH to it, and the copper hydroxide should precipitate right out.

from metal

chloric1 - 27-10-2005 at 18:04

Well if you must use metal, it would be better to put metal in ammonia solution with an air pump like the bubblers for fishtanks. When you have a bluish purple solution put a low temp heater and heat to 40 or 50 degrees C to drive off ammonia.

You see many metal hydroxides are quite asorbant and they can take in considerable chloride ions as an impurity. The beauty of this process is no anions are used and the Cu(OH)2 will be as pure as your metal and aqua ammonia. :cool:

guy - 27-10-2005 at 20:13

i put ammonia copper and hydrogen peroxide and LOTS of O2 gas came out. the solution went yellow
wat is going on?

12AX7 - 28-10-2005 at 02:02

Ooo, NH3 and oxidizers can do some fun things...

Tim

Darkblade48 - 28-10-2005 at 04:12

Quote:

Originally posted by guy
i put ammonia copper and hydrogen peroxide and LOTS of O2 gas came out. the solution went yellow

Went yellow? That's odd. The peroxide should release Cu+2 ions into the water which should complex with the ammonia (Assuming it's in excess) to form the complex [Cu(NH3)]+2 which is a deep blue colour.

chloric1 - 28-10-2005 at 04:32

Darkblade-I dont think its that simple. Copper is a catalyst for the decomposition of hydrogen peroxide. Also, if there is still metal left then it is in excess favoring cuprous complex formation.

Buy the way, in case you didn't know, electrolysis of NaCl or KCl with copper
electrodes gives yellow cuprous oxide.

Did this myself WAY the hell back in 1990 or 91. So young and naive, I thought I was goig to get chlorates!

The_Davster - 28-10-2005 at 04:48

I got the same result as Chloric1, I plugged a chlorate cell mock-up in with the wrong polarity last night, (Cu cathode, Pt anode), and yellow Cu2O was prouced in large volume right away. Oops...

guy - 28-10-2005 at 14:07

Okay, after a while the solution turned dark blue, but every time hydrogen peroxide is added it turns yellow (but no ppt). could it be [Cu(NH3)4]- but I think those are clear.

12AX7 - 28-10-2005 at 14:21

Hum, peroxide complex then? Possible higher oxidation state? Just guessing here.

Tim

The_Davster - 28-10-2005 at 21:59

You know, the origional reaction you proposed may actually work, I have been messing with a chlorate cell recently(Pt anode, Cu cathode). Upon running it for a few minutes, then turning it off and going to to bed and the next day a large ammount of Cu(OH)2 had precipitated and was forming stalacites off the Cu cathode, with small areas of a yellow stalacites(Cu2O) and CuO. I imagine the reaction you proposed above is the cause of this. Also, later into the cell run, even turning the current off for a second and Cu oxides and hydroxides would form. So perhaps a hot hypochlorite solution and copper metal would work.

chloric1 - 29-10-2005 at 14:39

Well, an additional note if I may. I like the idea of hypochlorite oxidation. I know electrochemical machineing uses unbeleivable current densities to erode metal into specified shapes. Maybe a strong NaOH solution with equimolar concentration of hypochlorite and a battery charger or simular would be great. Any powdered copper metal or cuprous oxide would be oxidized swiftly. But be careful as this may be favorable for oxide formation instead.

12AX7 - 29-10-2005 at 15:18

A basic solution will probably result in CuO2- and resulting plating out of metal rather than oxides (unless you want a metal deposit that is). For example, I got a deep blue solution (and no precipitate) from a sodium carbonate electrolyte (see CuO thread).

Tim

guy - 6-11-2005 at 17:57

tetraamine copper hydroxide
I want to make use of the tetraamine copper hydroxide to copper sulfate.

Is it possible to have the reaction

[Cu(NH3)4](OH)2 + MgSO4 ---> [Cu(NH3)4]SO4 + Mg(OH)2

filter ppt then heat the Cu(NH3)4SO4 to get CuSO4.

PS: What is the pH of tetraamine copper hydroxide? I think it should be really high since its so soluble.

[Edited on 11/7/2005 by guy]

The_Davster - 7-11-2005 at 01:38

Your origional rxn does not work. A length of copper wire in 10%NaOCl since my last post did nothing. Only where the liquid met the air was their blue copper salts forming on the wire, above the 'water' line.

chloric1 - 7-11-2005 at 09:29

add 3% hydrogen peroxide and you have a different story. OF coarse I believe you get the black oxide instead.

I figured it out

guy - 8-11-2005 at 15:20

Quote:

I wouldn't expect that reaction to go too far; both the reactant and product are solid.

Its pretty simple, why didnt I see it before. All I needed to do was add something to react with the Cu(OH)2 to make the reaction shift left. The reaction will proceed much faster if I added and acid such as acetic. Then I can add OH- to precipitate the rest out.

Darkblade48 - 8-11-2005 at 16:31

Quote:

Originally posted by guy
Its pretty simple, why didnt I see it before. All I needed to do was add something to react with the Cu(OH)2 to make the reaction shift left. The reaction will proceed much faster if I added and acid such as acetic. Then I can add OH- to precipitate the rest out.

Err, yes, adding an acid will then react with the Cu(OH)2 that you made, and make the equation make more Cu(OH)2 (I think you mean shift the reaction to the right, not to the left), but then, adding an acid to a solution that has OCl- isn't exactly wise, as you might liberate some chlorine gas.

guy - 8-11-2005 at 19:56

Yes I agree that it could be dangerous and irritating. Maybe this is a better idea: Adding excess Cl- to dissolve the copper hydroxide. Is this possible? and how would I reconvert these (CuCl4)2- back to Cu(OH)2?

12AX7 - 8-11-2005 at 21:44

If you have copper oxychloride, you have to boil with a hydroxide, which decomposes the Cu(OH,Cl)2 structure to CuO + RCl(aq), where R is Na, K, etc. Alternately, you can pyrolyze (mind the irritant fumes) to CuO + CuCl2 and leach out soluble CuCl2 with water.

You could dissolve the chloride in water (and with acid if you have oxychloride; use nitric) and add a lot of lead or silver nitrate, so as to precipitate the chloride ligands as PbCl2 or AgCl. Not really efficient, but it would work.

You could dissolve in strong sulfuric acid (no more than 50% or so) and boil out the HCl (mind the fumes), leaving copper sulfate.

Tim

[Edited on 11-9-2005 by 12AX7]

guy - 12-12-2005 at 17:39

I want to alter the Ammonia + Oxygen + copper method to produce copper sulfate. Instead of ammonia, I subsitute this with ammonium sulfate. This should produce copper hydroxide, which reacts with acidic ammonium sulfate to produce copper tetraamine sulfate. Is this going to work, because I don't want to waste my money going to buy an air pump and not have it work.

It would be nice to have someone try this and post their result. Thanks

Theoretical Reaction:
2Cu + O2 + 2H2O ------> 2Cu(OH)2
2OH- + 4NH4+ + Cu(OH)2 -----> [Cu(NH3)4]2+ + 4H2O

[Edited on 12/13/2005 by guy]

[Edited on 12/14/2005 by guy]

Darkblade48 - 12-12-2005 at 20:27

I would suppose that would work, but the first reaction, 2Cu + O2 + 2H2O ------> 2Cu(OH)2 will take place very slowly, and probably will not occur to any reasonable extent.

If you really want to get some Cu(OH)2, you can easily do the following:

1) Make a solution of MgSO4
2) Using copper electrodes, run a 9V (DC) current through it
3) Cu(OH)2 will precipitate out

chemoleo - 12-12-2005 at 20:53

Actually, I do think it would work, but I'd not just add ammonium sulphate, but an equimolar amount of H2SO4 - just to thorougly lower the pH, and then add the H2O2.
Be careful though, possibly this could catalyse the rapid decomposition of H2O2.
Without the extra acid I should think this'd be very slow.

guy - 12-12-2005 at 21:50

Yes i suppose adding sulfuric acid will speed up the reaction, but I have no acids other than Phosphoric and acetic. I was hoping that this reaction would be good yielding since it is very similar to the extraction of gold process.
4Au + 8CN- + O2 + 4H2O ---> 4[Au(CN)2]- + 4OH-

And copper is even more easily oxidized than gold, and the hydroxide is much more soluble.

[Edited on 12/13/2005 by guy]

chemoleo - 12-12-2005 at 22:05

Well then, use ammonium acetate, acetic acid, H2O2 and copper, or bubble with air and leave out the H2O2. THat should work too.

Darkblade48 - 12-12-2005 at 22:07

Quote:

Originally posted by guy
I was hoping that this reaction would be good yielding since it is very similar to the extraction of gold process.
4Au + 8CN- + O2 + 4H2O ---> 4[Au(CN)2]- + 4OH-

Ah, but in this reaction involves gold complexing with the cyanide anion, which removes gold from the left side of the equation. This pushes the equilbrium to the right, so that gold (solid) continually gets "dissolved".

guy - 12-12-2005 at 22:58

Quote:

Originally posted by Darkblade48
Ah, but in this reaction involves gold complexing with the cyanide anion, which removes gold from the left side of the equation. This pushes the equilibrium to the right, so that gold (solid) continually gets "dissolved".

Yes, so does this reaction, as copper complexes with ammonia.

2Cu + O2 + 2H2O + 8NH3 <------> 2[Cu(NH3)4]2+ + 4OH-
*************************************
Okay, this reaction would go a lot faster some ammonia to start with. Eventually this reaction should start becoming more and more basic and start to form more ammonia on its own.

I tried to do some math to calculate the amount of NH3 needed.

Cu(OH)2 <----> Cu2+ + 2OH- K1=1.6 x 10^-19
Cu2+ + 4NH3 <------> [Cu(NH3)4]2+ K2=1.2 x 10^12

K1 x K2 = 1.92 x 10^-7

Say I want to make 2 moles of [Cu(NH3)4]2+ in 1L water.

[2][2]2
------------ = 1.92 x 10^-7
[NH3]4

[NH3] = 80.3 M

80.3 + 4(2) = 88.3 M needed.

Therefore I need 5.83kg of (NH4)2SO4.

IS THIS RIGHT??
******
The efficiency of this reaction would be greatly increased if the Ksp of Cu(OH)2 is increased. I have tried searching for solubility graphs of Cu(OH)2 but got nothing. This information would be very useful. If the solubility can be increased, less (NH4)2SO4 will need to be used, and less contamination.

[Edited on 12/14/2005 by guy]

[Edited on 12/14/2005 by guy]

guy - 14-12-2005 at 19:18

Using Activated carbon

This site talks about the use of Carbon in Pulp (CIP) in the process of gold extraction. Could this be applicable for my tetraaminecopper sulfate process also?

guy - 15-12-2005 at 20:43

No, wrong calculations!!

I forgot to add the K of NH4+ + OH- --> NH3. Now the Equilibrium constant is 3.703 x 10^3. Very large, so I will be getting some good yield!

I'll post back if I ever get around to doing this.

The final reaction:

2NH3 + (NH4)2SO4 + Cu(OH)2 -------> [Cu(NH3)4]SO4 + H2O

{[Cu(NH3)4]2+}
-------------------------------------- = 3.703 x 103
[NH3]2[NH4+]2

Does anyone have any ideas on how to design this thing? I'm thinking a bucket with the solution (ammonia added in small intervals) and copper (elevated a few inches) and airstones connected to an airpump.

[Edited on 12/16/2005 by guy]

Design.bmp - 208kB

guy - 20-12-2005 at 15:35

Yes! I finally did it and I made a whole bunch of Cu(NH3)4SO4, which can be heated to drive off the ammonia. Should I post the whole procedure in the Prepublication forum? First I'm gonna to take some pictures.

Darkblade48 - 20-12-2005 at 18:28

Quote:

Originally posted by guy
Yes! I finally did it and I made a whole bunch of Cu(NH3)4SO4, which can be heated to drive off the ammonia. Should I post the whole procedure in the Prepublication forum? First I'm gonna to take some pictures.

I was under the impression that heating the complex salt would force the ammonia out, thereby destroying your product.

guy - 20-12-2005 at 22:35

Why would it force the ammonia out? It would first form tetraamine copper hydroxide which reacts with the Ammonium sulfate to form ammonia and water and tetraamine copper sulfate.

[Edited on 12/21/2005 by guy]

Picture 003.jpg - 662kB

Darkblade48 - 21-12-2005 at 00:04

I guess I should have worded it a little more carefully. I'm under the impression that the ammonia would just evaporate off, and might do so more readily under heat. The sample of the complex salt that I have slowly seems to be losing its dark blue colour, and everytime I open the jar, the odor of ammonia escapes from the bottle.

darkflame89 - 21-12-2005 at 02:39

Hey, guy, how many days did you take to dissolve your copper?

In the same vein, would a similar reaction involving the chloride complex of copper ion work?

Cu + 4Cl- + O2 + 2H20------------>[Cu(Cl)4]2- + 4OH-

woelen - 21-12-2005 at 10:59

No, with chloride it does not work. You instead get insoluble basic copper chloride, which in fact is copper hydroxide/chloride, something like Cu(OH)Cl.

guy - 21-12-2005 at 14:45

It took me three days to dissolve all the copper. I think it took a little longer than it should have because for one day I didn't put the copper directly over the air bubbler. Chloride might not work well becuase if the formation constant for [CuCl4]2- is much smaller than for forming tetraaminecuprate. And anyways, ammonia and ammonium sulfate are way cheaper than salt. And It wont have any contaminiants in the end once you heat it and drive off the ammonia.

chemoleo - 21-12-2005 at 14:54

VERY neat!

I wouldn't have thought it'd work that well! So you didnt use any H2SO4 to acidify?

So let me get this straight, all you had in there was ammonium sulphate and ammonia. How did you ensure that the ammonia didnt evaporate, because you must have put many cubic liters of air through there! Did you refill with ammonia now and then? The stink must have been enormous!
I can't see an easy way to trap the ammonia either... so you can't do that inside I suppose.

Still, very nice.

Let us know how easy it is to decompose the tetraamine sulphhate to normal copper sulphate.

guy - 21-12-2005 at 14:58

Well, to trap the ammonia i had to design it the pump would have to stay inside the container so it would keep recirculating the air and ammonia. I'll post the pic.
I did not use any acid since I have none.

HAHA. Anyways maybe I could use this copper sulfate somehow to make sulfuric acid.

[Edited on 12/21/2005 by guy]

Picture 001.jpg - 797kB

guy - 21-12-2005 at 15:00

Heres the inside

Picture 002.jpg - 745kB

guy - 21-12-2005 at 19:00

NO! What happened?!

When I evaporated the solution, copper HYDROXIDE precipitated out. How is this possible? Maybe I measured my reactants wrong. Does ammonium sulfate (from fertillizer) have hydrates?

These are my reagents.

10g Cu
54.5 g NH3 (from 10% solution) so 5.45 g of NH3
20.78 g (NH4)2SO4

Could this be a possible explanation?
2Cu + O2 + 4NH3 + 4NH4(+) <-----> 2[Cu(NH3)4](2+) + 2H2O

As I heat the solution in an attempt to collect the product, the remaing ammonia in the solution is first driven off. This shifts the equilibrium back left, forming once again copper hydroxide. :mad:

It seems like the only way to separate is to precipitate it using 100% ethanol or isopropyl. There is a lot of water so should I first try and remove some of the water with anhydrous CaCl2? Man, Im soo close yet so far away :mad:

[Edited on 12/22/2005 by guy]

[Edited on 12/22/2005 by guy]

Darkblade48 - 21-12-2005 at 23:21

I figured your product would decompose, that's why I suggested against heating it (the decomposition temperature of the complex is supposedly 120 C.

You can use methanol as well to precipitate out the salt (at least that's what I used, it's cheaper for me to get MeOH rather than IPrOH and EtOH in high purities).

Depending on how much water you have, you may or may not find that trying to get rid of the water with anhydrous CaCl2 is effective. From the looks of it though, you might have to use a lot of CaCl2 to suck up all that water.

darkflame89 - 21-12-2005 at 23:37

[Cu(NH3)4](2+) + 2OH(-) <---> Cu(OH)2 + NH3 (g)

Maybe the above reaction is favored when the temperature is higher. Reminds me of the temperature senstive cobalt chloride complex. There has to be some hydroxide ions inside the solution leftover. Unless the reagents are used in stoichiometric amounts and the reaction goes to completion where everything is used up, the solution is not going to be neutral. Perhaps you should test the solution to see if its neutral?

Else, you might have to filter out the copper(II) hydroxide, add ammonia to redissolve it, and add magnesium sulphate to swap the ions. This is under the assumption that you do not have sulphuric acid right?

Related website worth taking a look:Etching with air regenerated Cupric chloride

[Edited on 22-12-2005 by darkflame89]

guy - 27-12-2005 at 18:11

Arg! Its impossible to decompose the tetraamine copper sulfate. I washed completley dry with denatured alcohol. Then I heated it and it still turned to copper hydroxide. Why?? :mad:

chemoleo - 27-12-2005 at 18:17

Hmm, I feared this might happen - it is because you essentially eliminate ammonium sulphate, and its decomposition products.
How do you know it's copper hydroxide btw? It decomposes actually under heat to form black CuO. Is that what you get?

Or are you getting just dehydrated CuSO4, which is white, and dissolves in H2O to form the blue CuSO4 solution?

PS you heat it at what temperature? Please be more specific! Weigh it before and after heating! That should tell you what you got.

[Edited on 28-12-2005 by chemoleo]

guy - 27-12-2005 at 19:18

YAY!

Thanks for chemoleo for pointing out if it turns to CuO. It didnt turn to CuO. This time I heated it much longer and it turned to a white solid which is dehydrated CuSO4!!

It figured out why before it turned to copper hydroxide when I put it in water. I didnt heat it completly so some of the remaining tetraamine sulfate turned to copper hydroxide when I added it in excess water.

So I guess this way works for making CuSO4.

How long does it take for the anhydrous copper sulfate to dissolve? It looks like its taking a while.

[Edited on 12/28/2005 by guy]

chemoleo - 27-12-2005 at 19:38

Hehe that's good to hear.

Yes, I recently played with analytical anhydrous CuSO4, and I was quite surprised how long it takes to dissolve.

Please, do something for me.

Take your dried, alcohol precipitated tetraamine copper sulphate. Weigh whatever quantity you like.

Then heat it, until it almost starts to glow, and all colour has gone. Then WEIGH it again please.

You can do it even in that very same crucible as long as you determine the weight of the crucible.

Furthermore - how do you know it's CuSO4? Does EVERYTHING dissolve, or is there a residue, of whatever nature? Does the product dissolve ENTIRELY in H2O, with no remainder?

That's what I want to know!!!! It's good stuff, seriously, but you have to learn to describe what you did!

[Edited on 28-12-2005 by chemoleo]

guy - 27-12-2005 at 23:20

So far I have heated it up until it is a white solid. First it was 20g of tetraamine copper sulfate (The measurements are not accurate since I only have a kitchen scale). Then after heating to a white solid, the mass is approximately 10g.

I have to wait 'til tommorow to see if the product dissolves entirely. There seems to be some CuO residue in the product.

[Edited on 12/28/2005 by guy]

guy - 1-1-2006 at 19:08

Final Product

After dissolving the anhydrous salt and recrystallizing it, I finally obtained the copper sulfate pentahydrate crystals.

Copper Sulfate.jpg - 752kB