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Author: Subject: Separation of a US nickel
Sedit
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[*] posted on 3-11-2011 at 05:35


Quote: Originally posted by blogfast25  


IF (I'm still not wholly convinced) it's possible to separate the two with ammonia, then using the right amount of ammonia would be crucial.



This is the assumption I have come to conclude as well in that the Copper has such a high affinity for the Ammonia that it allows the hydroxide to precipitate out. Like I said before I never expected it to be this complicated I just seen a patent for precipitating Nickle onto carbon and reducing it for generating a means of Activated Nickle for hydrogenolysis reactions, they used Ammonia and it seemed right to me at the time to just complex the Copper and precipitate the Nickle. Experimentation confirmed it right away and after washing the fine precipitate with NH3 solution and a few times with water conversion to the HCl salt proved without a doubt what I had was Nickle as I obtained large needles of dark green crystals.

It was later that I found Nickle also is supposed to complex although I have about as much trouble believing this as you do believing that Ammonia will work. I have trouble with this because as I washed the precipitate with Ammonia I would do so until there was no more blue color from Copper forming... This is the main point that bugs me. Why would the precipitated oxide/hydroxide not complex well while the other salts will? The Nitrate seems to complex well, the HCl seems to complex well however the sulfate does not.

Its a very simple experiment and I would urge many here to repeat it because I am just not qualified enough to really do more then just report my results and allow everyone to figure out why.


Main points-
-- Use the Sulfate salt to avoid issues
-- Use Dilute Ammonia so that it precipitates before it complexes. I feel this may be crucial.





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[*] posted on 3-11-2011 at 05:51


Quote: Originally posted by Sedit  
Why would the precipitated oxide/hydroxide not complex well while the other salts will? The Nitrate seems to complex well, the HCl seems to complex well however the sulfate does not.



Aha, interesting point. The solubility product of Ni(OH)2 is 2.0 x 10<sup>-15</sup>. This is in all likelihood in competition with the complexation constant of 5.5 x 10<sup>8</sup> and this would explain why previously precipitated Ni(OH)2 cannot be solubilised by means of ammonia, or at least in your specific conditions.

I’ll try to put a bit of mathematical meat on this assertion a bit later on…

Update:

Well, well, quite an interesting but slightly ambiguous result.

For simplicity’s sake I’ll abbreviate the reactions to:

M2+ + 2 OH- === > M(OH)2
Ks = [M2+] x [OH-]<sup>2</sup> … (1)

And

M2+ + n A === > MA<sub>n</sub><sup>2+</sup>
Kf = [MA<sub>n</sub><sup>2+</sup>] / ( [M2+] x [A]<sup>n</sup>;), with n = 4 for Cu and n = 6 for Ni … (2)

For a weak base, we can approximate that [OH-] = √(Kb x Cb) … (3), with Kb the base constant (10<sup>-4.75</sup> for NH3) and Cb the formal concentration of the base. Because of poor protonation (weak base!), also [NH3] ≈ Cb.

Insert (3) into (1) and extract [M2+] = Ks / (Kb x Cb) … (4)

From (2) extract [M2+] = [MA<sub>n</sub><sup>2+</sup>] / ( Kf x Cb<sup>n</sup> ) … (5).

Since as (4) = (5), we now have a simple equation for

[MA<sub>n</sub><sup>2+</sup>] = ((Ks x Kf) / Kb ) x Cb<sup>n-1</sup>

The formula clearly shows the ‘conflict’ between Ks and Kf.

Assume now that we have a mixed precipitate of both hydroxides, washed (no extra OH-present) and we add ammonia to it until the actual final concentration of [NH3] ≈ 1 mol/l, then we can calculate the value of [MA<sub>n</sub><sup>2+</sup>] for both cases (the Ks for Cu(OH)2 is 2.6 x 10<sup>-19</sup>;) and obtain values of 0.16 mol/l for the copper complex and 0.056 mol/l for the nickel complex. A significant difference, but, assuming these values are roughly correct, not a great separation resolution…

I've never actually heard or read about a Cu/Ni ammonia based separation, which adds to my skepticism. If resolution is poor, that would explain why it isn't used in practice. Unlike the separation based on their sulphides which is 100 % in the right conditions.

I will have a go at this...

[Edited on 3-11-2011 by blogfast25]




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[*] posted on 5-11-2011 at 12:50


Blogfast, I should have known better than to take the shortcut with the concentrated ammonia. Thanks for taking the time to run the numbers. :)


I'm still confused with writing the chemical equations of the electrochemical separation of the metals. Can somebody help clear this up or provide a 'template'?

Another area of confusion for me is whether the insoluble matter which forms on the cathode is Cu or Cu+ (or a mixture). I've been running with the latter but would like to be sure.

Can the approximate efficiency of the cell be calculated if voltage, amperage (avg), time, and total dissolved anode mass are known? The hardest part for me lies in finding what the 'theoretical 100%' should be.

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[*] posted on 5-11-2011 at 13:31


Tanker:

Not so fast: I'm not 100 % sure of the numbers because there's quite a bit of simplifying going on in my little model, so I'm running a test. I've just prepared an equimolar mix of Cu(OH)2 and Ni(OH)2 for testing with 1.3 M NH3 tomorrow.

As regards the equations during the electrochemical separation, it's basically Cu(s) (from coin) === > Cu2+(aq) + 2e (oxidation of copper) and Cu2+(aq) + 2e === > Cu(s) (reduction of copper cations).

The theoretical yield is basically [url=http://en.wikipedia.org/wiki/Faraday's_laws_of_electrolysis]Faraday's Laws of electrolysis[/url] . In short, using Faraday's constant, the valence of the cation, the time and current applied during that time, the theoretical amount of opper deposited can be calculated.

Not sure where this $&%^&*^&*&ing syntax error comes from: here's the url to Wiki's Faraday's Law of electrolysis:

http://en.wikipedia.org/wiki/Faraday's_laws_of_electrolysis

What forms on the cathode is pure copper: Cu+ isn't a substance because it is not electrically neutral. If Cu+ was involved, Cu2SO4 would deposit: that's not the case here.



[Edited on 5-11-2011 by blogfast25]




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[*] posted on 6-11-2011 at 08:12


Here’s the moist equimolar mix of Cu(OH)2 and Ni(OH)2, I made about 0.04 mol (of each) for further use, made from neutralising equal (molar) amounts of NiCl2 and CuSO4, and washing profusely:




Then a few ml of NH3 1.5 M were dispensed into a clean test tube and a small pinch of the mixed hydroxides was lowered into it with a glass rod. It dissolved instantly and completely, resulting in that typical clear, deep blue of cupper (II) tetrammonia complex cations. I kept adding pinch after pinch and it just kept dissolving completely. Unless some inter-copper/nickel species is being formed, this is clear evidence that separating Cu2+ and Ni2+ by means of selective complexation with ammonia is basically impossible.

A tube with mixed complexes, cupper (II) tetrammonium and nickel (II) hexammonia:






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[*] posted on 6-11-2011 at 10:12


Funny blog, I just performed the same experiment with the same results.

I precipitated my sulfates from the coins using NaOH and dissolved the preciptate completely in NH3 solution forming the blue solution.

However keep in mind that the experiments of mine that did precipitate a GREEN precipitate was done by adding the NH3 solution slowly to the salt solution.

Its taking forever for these coins to dissolve into the H2SO4 but I should be ready for better test tube runs soon enough.





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[*] posted on 6-11-2011 at 13:18


Quote: Originally posted by Sedit  
However keep in mind that the experiments of mine that did precipitate a GREEN precipitate was done by adding the NH3 solution slowly to the salt solution.



Did the precipitate persist? Did it eventually dissolve? If 'yes' and 'no', then you should replicate your experiment and note the experimental conditions accurately here.




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[*] posted on 6-11-2011 at 13:50


Well after filtering it I washed it several times with NH3 solution and it never dissolved, the blue color just got lighter and lighter, washed with clean H2O and then dissolved the precipitate in HCl and it yielded large dark green needles on drying. I reduced this with Al foil and yielded a magnetic black precipitate....

It's bugging the hell out of me honestly.

It HAS to be a dilution issue but only experimentation will tell. I'm considering just reducing my sulfate solution down so I can get on the ball with it.





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[*] posted on 6-11-2011 at 14:00


Blogfast: If you still have the test tube with the mixed metal complexes, maybe you could slowly add distilled/DI water and see if anything precipitates??

I was shocked when I applied Faraday's Laws to my setup and came within 200mg! It had me pulling my hair out for a little while because I screwed up the total run time of the cell. Looking back at my notes, I found the mistake and recalculated.

Good stuff - will help tremendously with this and future electrolysis experiments, etc. Thanks for the nudge.

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[*] posted on 6-11-2011 at 14:14


Quote: Originally posted by m1tanker78  
Blogfast: If you still have the test tube with the mixed metal complexes, maybe you could slowly add distilled/DI water and see if anything precipitates??

Tank


I'll try that tomorrow... Perhaps at very low NH3 concentrations some separation could be possible. I'm not hopeful though.




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[*] posted on 6-11-2011 at 14:37


It will not precipitate, try converting some of the hydroxide to a dilute HCl solution to generate the mixed salts. Dissolve the salts in H2O and drip in dilute NH3, then you will almost surely see the precipitate I was seeing. I have seen the blue and I have seen the green, I was just not careful enough in any of my experiments to determine why I got varying results in different instances.

Perhaps referencing back to the patent that gave me the idea would shed some light. Give me a minute and I will link it in here.


[edit]
Ah good the pictures of the precipitate are still there as well I thought they where lost. It was the HCl salt I was using. Since the experiments where over a year old my brain has a tendency to forget details.

http://www.sciencemadness.org/talk/viewthread.php?tid=10527&...

Looking back in that thread I found it cool that I noticed the same think Tank did with Copper selectively precipitating from the solution and had forgot all about it.

[Edited on 6-11-2011 by Sedit]





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[*] posted on 7-11-2011 at 07:02


I think you're confusing the colour of the mixed hydroxide with that of Ni(OH)2.

I added some 1 M HCl to the solution and the blue/green precipitate of Cu(OH)2/Ni(OH02 reappeared (I will add a picture later on). The (paper) pH of the supernatant liquid was only about 5.

This is 'game over', as far as I'm concerned.

The precipitate obtained by adding HCl to the mixed complexes solution:



Highly suggestive of mixed Cu/Ni hydroxides!

[Edited on 7-11-2011 by blogfast25]

[Edited on 7-11-2011 by blogfast25]




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[*] posted on 8-11-2011 at 09:10


Update:

On request by Sedit, I added quite a lot of NH4Cl to the 1.5 M ammonia, prior to adding the mixed Ni/Cu hydroxides. It made no difference: the mixed hydroxides dissolved all the same an no green precipitate or green insoluble matter was observed.




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[*] posted on 8-11-2011 at 10:25


Thanks for running the test blog. I wish I could say that I am mistaking the precipitate I obtained for just mixed hydroxide but I am not. Washing of that precipitate with Ammonium hydroxide gave me a pure light green precipitate with no blue and on addition of HCl formed dark green needles of nickle chloride.

I wish I could explain it but as of yet I can't. Its a very humbling adventure to perform an experiment repeatedly and gain various results with no real idea why but I will get to the bottom of it sooner or later. Now I fully understand POKs frustration after making the Potassium and no one was able to reproduce his results at first.





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[*] posted on 8-11-2011 at 14:56


Quote: Originally posted by Sedit  
Thanks for running the test blog. I wish I could say that I am mistaking the precipitate I obtained for just mixed hydroxide but I am not. Washing of that precipitate with Ammonium hydroxide gave me a pure light green precipitate with no blue and on addition of HCl formed dark green needles of nickle chloride.

I wish I could explain it but as of yet I can't. Its a very humbling adventure to perform an experiment repeatedly and gain various results with no real idea why but I will get to the bottom of it sooner or later. Now I fully understand POKs frustration after making the Potassium and no one was able to reproduce his results at first.


Please do get to the bottom of it. I'm interested in your explanation...




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[*] posted on 8-11-2011 at 15:08


Quote: Originally posted by blogfast25  
Update:

On request by Sedit, I added quite a lot of NH4Cl to the 1.5 M ammonia....


Was this a typo? I thought Sedit had used dilute ammonia to wash the precipitate. I could try to reproduce this for a '3rd opinion' but I'd need to add the copper back into the nickel(II) chloride solution to simulate the approximate Cu/Ni ratio of US nickels.

I *did* observe a precipitate when I added 18M ammonia to a nickel chloride solution before. That's the one where the precipitate shrunk down to a compact pellet when air-dried (upthread somewhere). I'm pretty sure that iron was the culprit but never got around to testing it.

If all goes as planned, I'll be testing the electrolytic separation of the metals soon with a different acid - probably H2SO4 (drain cleaner) this time. I'm also tempted to redo a previous [undocumented] experiment where I used ammonium sulfate electrolyte.

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[*] posted on 8-11-2011 at 16:58


I did, I asked Blogfast to add some Ammonium chloride to his blue solution to see if there was some sort of common ion effect at play.

As tank is confirming sometimes you WILL get a precipitate, sometimes it is the mixed hydroxides giving a blue color, SOME times its the desired green Nickle precipitate. Its such an odd thing that there is some factor at play that I just can't think of.




Added a file so I don't lose it, It deals with the solubility of Copper(II) in an ammonia chloride solution as Ammonia concentration rises. There is a drastic increase as excess Ammonia enters the system.

[Edited on 9-11-2011 by Sedit]

Attachment: 961-978-1-PB.pdf (1.4MB)
This file has been downloaded 788 times






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[*] posted on 9-11-2011 at 18:56


I decided to continue with the previous electrolysis experiment (to separate Cu and Ni) rather than trying a different electrolyte. After about 2H, the anode began to disintegrate.

I added more HCl and upped the voltage to 1.5V. Current flow began at > 6A but steadily fell to < 3A as the rest of the anode dissolved away. The final mass of the [fragments of] anode was 32.9g. This means that 51.3g-worth of nickels dissolved. The total combined run time was 20:15.

If my reasoning is correct, there should be an equivalent of close to 13 grams of Ni metal that was dissolved. In spite of increasing to 1.5V, no detectable amount of Ni was deposited at the cathode (I kept a close eye on that!). The extra acid helped form a firmer (but still easily removable) deposit of spongy copper. This means that less, if any, drops back into the electrolyte at scraping time.

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[*] posted on 13-11-2011 at 14:56


Sedit, not sure if this could help but I came across this while searching for something else. This source claims the same thing you do (scroll to the very bottom).

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[*] posted on 17-11-2011 at 12:23


Some results are in, as noted by myself in the Cu++ Ni++ separation thread and by tank here electrochemical separation holds high potential. I took a solution of the Sulfate salts and used 5v 500ma with two carbon electrodes over night and got a large amount of Copper precipitated out with no notice of Nickel in the mix.

This suggest that simply dissolving the Coins in HCl/H2O2 and subjecting the solution to electrolysis will achieve the desired separation. This avoids the need to make the electrodes that Tank produced although I do think what hes doing is a great idea since it saves on reagents and you don't get the carbon contamination that needs to be filtered like I have.

I have no idea how complete this will run but if in doubt a combination of the two methods can be employed to clean any remaining copper from the mix since the Copper amine complex is 5x more likely to form if I read blogfast numbers correctly. I see now where he was going wrong. He is using to much ammonia solution. You only need to use just enough to neutralize the solution and no more. After the precipitate has formed then and only then should a slight excess of ammonia be used. There may be a small amount of loss of Nickle this way but it would ensure the removal of all the Copper from the mix.


There you have it folks, The separation can be achieved on the cheep. If I had more time and money(I'm all out of HCl and H2SO4 takes to damn long to dissolve the coins) I would produce a full writeup on a large scale as I do wish to obtain a good amount of Nickle salts right now to play around with Nickle salts as reducing agents.





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[*] posted on 17-11-2011 at 12:51


Quote: Originally posted by Sedit  
I have no idea how complete this will run but if in doubt a combination of the two methods can be employed to clean any remaining copper from the mix since the Copper amine complex is 5x more likely to form if I read blogfast numbers correctly. I see now where he was going wrong. He is using to much ammonia solution. You only need to use just enough to neutralize the solution and no more. After the precipitate has formed then and only then should a slight excess of ammonia be used. There may be a small amount of loss of Nickle this way but it would ensure the removal of all the Copper from the mix.


Hmmm… for one, I didn’t do it that way. I started from fresh precipitate (1:1 molar ratio Ni:Cu), then added it to ammonia. It dissolved and I kept adding precipitate and it kept dissolving.

I will in the coming days carefully dissolve some of the precipitate in the minimum quantity of HCl, then slowly and in small aliquots add weak ammonia (about 0.6 M) and see what happens. My guess, also based on the numbers, is that the amount needed to get the Cu<sup>2+</sup> complexed will be enough to complex at least part of the Ni<sup>2+</sup> complexed. And partial separation isn’t of much use. Which is why I believe the method isn’t in industrial use…




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[*] posted on 17-11-2011 at 13:03


Sedit, we must me 'tracking'.. :D

I reduced out much of the remaining copper (I hope - will confirm later) overnight by using a bare carbon rod anode. I'm also working on a blot test to roughly estimate copper/nickel in solution (infancy stage). I settled on 1.55V @ ~ 200mA. This seemed to be a good compromise so I wouldn't have to babysit the cell - i.e. get some SLEEP! The compromise also avoided shredding of the anode and excess O2 evolution which is waste, anyway. The spongy copper deposit becomes extremely fragile as the Cu concentration nears extinction. Filtration took care of the copper that flaked off and settled out. The solution is beautifully colored but I'm going to boil it down so I can kick the nickel out.

Sedit if you can regulate the voltage on your ps you can probably just stack the coins in the electrolyte and use a simple gravity contact to dissolve the alloy relatively quickly. At 500mA you can expect to dissolve and remove ~ 0.6g of copper per hour if my fuzzy math is correct.

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[*] posted on 17-11-2011 at 13:58


Quote: Originally posted by blogfast25  


Hmmm… for one, I didn’t do it that way. I started from fresh precipitate (1:1 molar ratio Ni:Cu), then added it to ammonia. It dissolved and I kept adding precipitate and it kept dissolving.

I will in the coming days carefully dissolve some of the precipitate in the minimum quantity of HCl, then slowly and in small aliquots add weak ammonia (about 0.6 M) and see what happens. My guess, also based on the numbers, is that the amount needed to get the Cu<sup>2+</sup> complexed will be enough to complex at least part of the Ni<sup>2+</sup> complexed. And partial separation isn’t of much use. Which is why I believe the method isn’t in industrial use…


Drip the Ammonia in slowly and allow it to run down the side of the test tube, it should show you the precipitate as it forms. I agree there will be some losses in Nickle but if it completely removes the copper as I suspect that will be acceptable.

I am finishing running the cell right now and the amount of Copper is greatly diminishing with still no evidence of Nickle so I suspect that this may be a highly effective way to remove all the copper. Since I used H2SO4 to dissolve the coins I do not want to boil down the solution so I am going to go with the Ammonia anyway to precipitate the Nickle out.


I just took a small sample of my mix and added NaOH to it. LOL, after cleaning up my mess I repeated it much slower:D. Anyway, the hydroxide precipitate was green:) I seen no strong evidence of Copper hydroxide which I find very encouraging meaning a final cleanse with Ammonia hydroxide may be a moot point and unneeded.



PS: I made a mistake, my power supply was 5v 825ma.

[Edited on 17-11-2011 by Sedit]





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[*] posted on 17-11-2011 at 15:35


Sedit, 5V and no nickel deposit?? Is your voltage dropping under load?

Quote:
I have no idea how complete this will run..

An old plater once told me that the last 1-2% is a PITA. I tend to believe him even though he didn't have any math or graphs to back up his life work. LOL

We also should realistically consider that most platers dummy plate (basically what we're doing here) overnight. IOW, they're on a schedule and some don't even bother maintaining the existing bath. I believe that at a certain point (arbitrarily ~ 1% w/w Cu/Ni - maybe less), there is an equilibrium between Cu deposited and Cu re-oxidized (dissolved). Like it or not, if Cu sponge is left to sit in an acidic solution, it will semi-promptly dissolve especially with air exposure. To sum it up, when the reduction curve meets the oxidation curve, it's time to stop. I just want to stress that even when copper stops depositing on the cathode, that doesn't mean it's completely gone from the bath although it should be damned close.

I'm still curious about the ammonia route and look forward to Blogfast's observations when he tries to replicate.

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[*] posted on 17-11-2011 at 17:32


I don't know the exact numbers Tank my multimeter needs batteries sorry. I can't say there is truly no Nickle but I do know the color coming out is highly suggestive of pure Cu as there is no darkness to it as observed when Nickle is in the mix.

I will replicate the Ammonia route soon enough as I will use it to wash the remaining few percent of Cu out of the solution. I will neutralize the solution left

I have to say though there seems little remains judging by the color of the precipitate acquired from NaOH. It was a pure light green precipitate with no signs of blue to it that Copper hydroxide gives.

A final wash with Ammonia should be more then sufficient to remove that last of the Copper as a complex I believe.

I honestly feel we are very very close to cracking this one once and for all Tank. ;)





Knowledge is useless to useless people...

"I see a lot of patterns in our behavior as a nation that parallel a lot of other historical processes. The fall of Rome, the fall of Germany — the fall of the ruling country, the people who think they can do whatever they want without anybody else's consent. I've seen this story before."~Maynard James Keenan
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