| Pages:
1
2 |
guy
National Hazard
Posts: 982
Registered: 14-4-2004
Location: California, USA
Member Is Offline
Mood: Catalytic!
|
|
Cu(OH)2 reaction
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
International Hazard
Posts: 1583
Registered: 20-8-2004
Location: USA
Member Is Offline
Mood: oscillating
|
|
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
Hazard to Others
Posts: 411
Registered: 27-3-2005
Location: Canada
Member Is Offline
Mood: No Mood
|
|
| 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.
|
|
|
chloric1
International Hazard
Posts: 1070
Registered: 8-10-2003
Location: GroupVII of the periodic table
Member Is Offline
Mood: Stoichiometrically Balanced
|
|
from metal
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.
Fellow molecular manipulator
|
|
|
guy
National Hazard
Posts: 982
Registered: 14-4-2004
Location: California, USA
Member Is Offline
Mood: Catalytic!
|
|
i put ammonia copper and hydrogen peroxide and LOTS of O2 gas came out. the solution went yellow
wat is going on?
|
|
|
12AX7
Post Harlot
Posts: 4803
Registered: 8-3-2005
Location: oscillating
Member Is Offline
Mood: informative
|
|
Ooo, NH3 and oxidizers can do some fun things...
Tim
|
|
|
Darkblade48
Hazard to Others
Posts: 411
Registered: 27-3-2005
Location: Canada
Member Is Offline
Mood: No Mood
|
|
| 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
International Hazard
Posts: 1070
Registered: 8-10-2003
Location: GroupVII of the periodic table
Member Is Offline
Mood: Stoichiometrically Balanced
|
|
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!
Fellow molecular manipulator
|
|
|
The_Davster
A pnictogen
Posts: 2861
Registered: 18-11-2003
Member Is Offline
Mood: .
|
|
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
National Hazard
Posts: 982
Registered: 14-4-2004
Location: California, USA
Member Is Offline
Mood: Catalytic!
|
|
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
Post Harlot
Posts: 4803
Registered: 8-3-2005
Location: oscillating
Member Is Offline
Mood: informative
|
|
Hum, peroxide complex then? Possible higher oxidation state? Just guessing here.
Tim
|
|
|
The_Davster
A pnictogen
Posts: 2861
Registered: 18-11-2003
Member Is Offline
Mood: .
|
|
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
International Hazard
Posts: 1070
Registered: 8-10-2003
Location: GroupVII of the periodic table
Member Is Offline
Mood: Stoichiometrically Balanced
|
|
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.
Fellow molecular manipulator
|
|
|
12AX7
Post Harlot
Posts: 4803
Registered: 8-3-2005
Location: oscillating
Member Is Offline
Mood: informative
|
|
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
National Hazard
Posts: 982
Registered: 14-4-2004
Location: California, USA
Member Is Offline
Mood: Catalytic!
|
|
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
A pnictogen
Posts: 2861
Registered: 18-11-2003
Member Is Offline
Mood: .
|
|
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
International Hazard
Posts: 1070
Registered: 8-10-2003
Location: GroupVII of the periodic table
Member Is Offline
Mood: Stoichiometrically Balanced
|
|
add 3% hydrogen peroxide and you have a different story. OF coarse I believe you get the black oxide instead.
Fellow molecular manipulator
|
|
|
guy
National Hazard
Posts: 982
Registered: 14-4-2004
Location: California, USA
Member Is Offline
Mood: Catalytic!
|
|
I figured it out
| 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
Hazard to Others
Posts: 411
Registered: 27-3-2005
Location: Canada
Member Is Offline
Mood: No Mood
|
|
| 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
National Hazard
Posts: 982
Registered: 14-4-2004
Location: California, USA
Member Is Offline
Mood: Catalytic!
|
|
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
Post Harlot
Posts: 4803
Registered: 8-3-2005
Location: oscillating
Member Is Offline
Mood: informative
|
|
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
National Hazard
Posts: 982
Registered: 14-4-2004
Location: California, USA
Member Is Offline
Mood: Catalytic!
|
|
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<sup>+</sup> + Cu(OH)2 -----> [Cu(NH3)4]<sup>2+</sup> + 4H2O
[Edited on 12/13/2005 by guy]
[Edited on 12/14/2005 by guy]
|
|
|
Darkblade48
Hazard to Others
Posts: 411
Registered: 27-3-2005
Location: Canada
Member Is Offline
Mood: No Mood
|
|
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
Biochemicus Energeticus
Posts: 3005
Registered: 23-7-2003
Location: England Germany
Member Is Offline
Mood: crystalline
|
|
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.
Never Stop to Begin, and Never Begin to Stop...
Tolerance is good. But not with the intolerant! (Wilhelm Busch)
|
|
|
guy
National Hazard
Posts: 982
Registered: 14-4-2004
Location: California, USA
Member Is Offline
Mood: Catalytic!
|
|
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<sup>-</sup> + O2 + 4H2O ---> 4[Au(CN)2]<sup>-</sup> + 4OH<sup>-</sup>
And copper is even more easily oxidized than gold, and the hydroxide is much more soluble.
[Edited on 12/13/2005 by guy]
|
|
|
| Pages:
1
2 |
![[*]](images/xpblue/default_icon.gif){kind=icon}
