Sciencemadness Discussion Board - Separation of a US nickel

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 10-15. This is in all likelihood in competition with the complexation constant of 5.5 x 108 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-]2 … (1)

And

M2+ + n A === > MAn2+
Kf = [MAn2+] / ( [M2+] x [A]n

, 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 (10-4.75 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+] = [MAn2+] / ( Kf x Cbn ) … (5).

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

[MAn2+] = ((Ks x Kf) / Kb ) x Cbn-1

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 [MAn2+] for both cases (the Ks for Cu(OH)2 is 2.6 x 10-19

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]

m1tanker78 - 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.

Tank

blogfast25 - 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]

blogfast25 - 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:

Sedit - 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.

blogfast25 - 6-11-2011 at 13:18

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.

Sedit - 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.

m1tanker78 - 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.

Tank

blogfast25 - 6-11-2011 at 14:14

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.

Sedit - 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]

blogfast25 - 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]

blogfast25 - 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.

Sedit - 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.

blogfast25 - 8-11-2011 at 14:56

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...

m1tanker78 - 8-11-2011 at 15:08

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.

Tank

Sedit - 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]

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m1tanker78 - 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.

Tank

m1tanker78 - 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).

Tank

Sedit - 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.

blogfast25 - 17-11-2011 at 12:51

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 Cu2+ complexed will be enough to complex at least part of the Ni2+ complexed. And partial separation isn’t of much use. Which is why I believe the method isn’t in industrial use…

m1tanker78 - 17-11-2011 at 13:03

Sedit, we must me 'tracking'..

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.

Tank

Sedit - 17-11-2011 at 13:58

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

. 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]

m1tanker78 - 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.

Tank

Sedit - 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.

Sedit - 18-11-2011 at 15:34

First of all allow me to apologize for the size of these photos,it exceeds forum limits so I was forced to remove the JPG ending, just download them and rename it with the jpg ending. , if someone can at the very lest show me how to resize them inline I will edit this but I am forced to present these pictures because like any good scientist It is my duty to back up claims made by myself and where as I applaud Blogfast for his valiant effort I knew there where minor and crucial alterations being made effecting the out come of the reaction.

Here are two photos of freshly prepared solutions from the coins made by dissolving them in H2O2 and HCl.

I slowly dripped in Ammonium hydroxide solution so that the reaction can be observed as two layers where the boundary was where the reaction was taking place. You can clearly see two very useful bits of information in them. You can clearly see the green precipitate falling out of the solution and up top the copper amine complex.

This my friend is case closed

, My mind can now rest relatively peacefully tonight. Now on to the next challenge of using this as nothing more then a final wash to remove trace Cu left over from the electro-deposition.

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Would right more but am in a hurry, enjoy :cool:

[Edited on 19-11-2011 by Sedit]

m1tanker78 - 18-11-2011 at 17:32

if someone can at the very lest show me how to resize them inline I will edit this[...]

Sedit, here are your images posted inline - rotated and resized:

Very nice demo Sedit! Create an account with Picasa, Photobucket, etc. You can then post inline images using the appropriate tags.

[~img]link to your image goes here[~/img]

The '~' is so it will display properly as text. Remove it from each tag when posting pics!

Tank

Sedit - 18-11-2011 at 23:55

I would rather just try to download and image editor that is virus free. and stable on my piece of shit computer but until then I am forced to do what I just did. Thank you for getting the pictures right. I just feel like a kewl talking out my ass the whole time I have been unable to present my work in the form of a pictorial.

I will make this short because quite frankly, I indulged in spirits a little to much in the past few hours and typing is very very hard,

You can see the green precipitate that people are having trouble achieving clear as day so Blogfast declaration of game set match, checkmate has been visually dis-proven and that was my main goal right now to ensure that it is known that my claim is valid.

Ok its bed time fellows good night, I hope you all gained some incite as to why we are all getting various results.

cyanureeves - 19-11-2011 at 06:22

is it morning yet? nice work gentlemen now if only the blue is indeed just copper and the green just nickel. maybe this only works with the hcl method,either way its darn good.i quit the ammonia method because i electrolyzed in sulfuric acid solution and not hcl. but i think sodium hydroxide would've dropped both the copper and nickel from my solution in form of hydroxides and further heating the precipitates with sodium hydroxide will separate the copper from nickel when taken back to solution as the nickel will remain as insoluble hydroxide.nickel in its hydroxide should give me nickel oxide if i roast it in open air. will a coin dissolved in either hcl and peroxide or sulfuric acid be the same as electrolyzing the coin in the solutions?because using current is just much quicker.its awesome how nickel can make things so pretty and shiny and super hard.

[Edited on 19-11-2011 by cyanureeves]

Sedit - 19-11-2011 at 09:32

LOL, uggg, yeh its morning..... Funny how EtOH is a couple hours of fun then 24 hours of side effects

. Beer taste so much worse in my mouth the second day.

My experiments in the past suggest that the green precipitate is nothing but the Nickle. Unless Copper chlorides can fit themselves into the crystal structure of Nickle chloride then its highly improbable that there is copper contamination in the green precipitate.

You can see in the photos a transition from what appears to be the mixed hydroxides to the pure green nickle but then again that could be nothing more then a trick of light as the green precipitate attempts to shine through the dark blue solution.

blogfast25 - 19-11-2011 at 09:37

@Sedit:

Perhaps the remainder of the spirits were still talking but please restrain from putting words into my mouth: I never wrote “game set match, checkmate”, I wrote “game over”, not as a triumphant expression but as one of: ‘I’ve had enough of this’. Not once have I doubted your word, but I was unable to replicate your results and doubted (and still do) any feasibility of actual, practical separation this way.

Using a slightly acidic (this is important!) solution of CuCl2/NiCl2, both approx. 0.2 M, I’ve managed to produce your green precipitate, even though the following crap photo doesn’t do it justice:

With the naked eye, the bottom of the interface between added ammonia solution and Cu/Ni solution can be seen to be a green precipitate. I added very slowly the weak ammonia (0.5 to 1 M NH3) to the Cu/Ni solution and at some point, yes, this strange interface, including Ni(OH)2 precipitate and copper ammine complex, occurs.

With hindsight, explaining what is going on here is not difficult and doesn’t contradict what I wrote. This is a highly non-homogenous, non-equilibrium situation: at the interface the pH is just so that Ni(OH)2 precipitates, yet there is just enough free NH3 to complex the Cu.

On shaking everything dissolved and neither Ni(OH)2 nor the copper ammine complex survives because resulting solution is acidic: neither NH3 nor OH- can survive in appreciable amounts.

To make this somehow the mechanism for effective separation between Cu2+ and Ni2+ is a pipedream, IMHO: it would require such accurate control of all concentrations and pH, as well as heterogeneous conditions to make it work.

I don’t believe separation between various ammine complexes (Zn, Co, Cu, Ni…) on the basis of differential complexation and differences in solubility of their hydroxides is actually viable, nor have I ever come across descriptions of such methods.

The model I put together applies (like most chemical models) in homogenous equilibrium conditions.

It's also no big surprise that we didn't obtain the same results at first: I was doing something very differently from you...

[Edited on 19-11-2011 by blogfast25]

Sedit - 19-11-2011 at 09:57

I'm just busting your balls blogfast,

I do honestly appreciate your efforts without a doubt. Many times people will discredit something and never once head to the lab to see for themselves and the fact you did quickly earned my respect without a doubt. You seen a claim you did not feel was valid and at the speed of lightning setup to reproduce the experiment. That's not just a good scientist that's someone with a love and a passion for what they do ::End Kissing ass::

I noticed the same as well, once shaking happened the precipitate vanished. This is why it is so had to duplicate these results. The scale I was running it in in the past was much higher leaving alot more room for error when it comes to concentration. If the concentration of NH3 to chloride solution is just right the Nickle hydroxide will precipitate out. Any more and its complexed. Any less and it stays as the chloride salt.

I also question the practicality of my method but for different reasons. On a larger scale its very simple to achieve the stable precipitate however... It is very hard to filter and given it should have multiple washes to ensure purity this is a huge hurdle. Because of this I feel it be best you use it as a final wash of the electrochemical method and nothing more so that if your going to filter it lets just do it once and do it right because it reminds me of trying to filter clay.... it just clogs up everything and hardly even settles making decantation a pain in the ass as well.

blogfast25 - 19-11-2011 at 10:15

Ok, no hard feelings!

It is quite remarkable, these ‘interface’ conditions, very peculiar…

Sedit - 19-11-2011 at 11:12

Yes the interface is just clear evidence of how concentration is highly effecting the outcome of the reaction.

m1tanker78 - 19-11-2011 at 16:22

will a coin dissolved in either hcl and peroxide or sulfuric acid be the same as electrolyzing the coin in the solutions?

Not really. A well-controlled electrolysis will remove a significant amount of Cu. Further electrolysis with a carbon anode will remove the vast majority of remaining Cu.

Simply dissolving the alloy in acid/peroxide will conserve the metal ratio so they must be separated by some other means. Note that Sedit and I observed that some Copper Sulfate will crystallize from H2SO4/H2O2 so that's one soft exception.

Tank

Sedit - 19-11-2011 at 16:26

Yeh but dissolving them in H2SO4 takes so frigging long there is no way I would bother with H2SO4 and H2O2, its just a waste of time and reagents when HCl and H2O2 can handle them over night for the most part.

m1tanker78 - 19-11-2011 at 17:17

@Sedit, true but it can be sped up some by adding a bit of nitrate. Watch the fumes (from a distance).

For those interested in precipitating nickel with aluminum:

A word of caution about the grade of aluminum used. Most aluminum foils will give mixed or bad results. I set up a simple experiment to demonstrate this. I used approx 10mL of electro-refined Nickel(II) Chloride in each shot glass. Food grade Al foil was added to one while pure Al drill shavings were added to the other. The precipitates were rinsed equally well and left in water to photograph.

Foil-precipitated nickel on left. Pure Al-precipitated nickel on right:

Weird/beautiful magnetic nickel flower (still under water).

I'm glad I didn't precipitate the whole lot of electrolyte with Al foil. Both precipitates were about equal in their strong attraction to a magnet. I can't help wondering what component of the alloyed Al gives the reddish color.

I'll be conducting a few experiments with ascorbic acid. I would love to finally be able to say with some certainty how much copper remains in the electrolyte after refining. If anyone can save me a little time on this, I'd love to read your suggestions.

Thanks,

Tank

blogfast25 - 20-11-2011 at 06:46

I'm glad I didn't precipitate the whole lot of electrolyte with Al foil. Both precipitates were about equal in their strong attraction to a magnet. I can't help wondering what component of the alloyed Al gives the reddish color.
Tank

Colloidal copper, most likely.

Sedit - 21-11-2011 at 19:18

Oh my god do I have great news for all of you.

Check this picture out and please resize and inline if feeling incline to do so.

Check it out. I diluted the Copper/Nickle dissolved in HCl a bit, at this time it was a brown solution that looked green when in trace amounts like a drop.

So I carefully added Ammonium hydroxide and low and behold I achieved this what you see in the picture. Prior to the Dark blue copper amine complex forming all of the Nickle precipitates out as the hydroxide. Notice the green solution on top and the large quantity of green precipitate on the bottom.

Tell me, do you seen anything that suggest that this is the mixed hydroxides? The key seems to be to loss a small amount of nickle by keeping the solution slightly acidic.

This is a MAJOR break through.

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m1tanker78 - 22-11-2011 at 06:06

Sedit, I don't think I follow. Did you use electro-refined nickel chloride or the mixed chlorides? If the latter, where's the copper? If you decant the supernatant and rinse several times then allow the ppt to air dry, I'll bet you'll see results similar to mine (fluffy ppt dries to a tiny pellet). The residue(s) left on the walls of the container as the ppt dries could provide some clues as well.

I had very similar results when I used strong ammonia. But then, the chloride solution was highly acidic IIRC.

Tank

blogfast25 - 22-11-2011 at 06:30

Tank, can you publish these photos above? I can't see them (no recognised file extension)

m1tanker78 - 22-11-2011 at 06:42

Tank, can you publish these photos above? I can't see them (no recognised file extension)

Sedit's images:

Tank

[Edited on 11-22-2011 by m1tanker78]

blogfast25 - 22-11-2011 at 07:06

The key seems to be to loss a small amount of nickle by keeping the solution slightly acidic.

Hmmm… that just doesn’t make any sense.

The Ks for Ni(OH)2 is 5.5 x 10-16

So, [Ni2+] x [OH-]2 = 5.5 x 10-16 … Eq.1

Also, at all times:

[H3O+] x [OH-] = 10-14 = Kw … Eq.2

Isolate [H3O+] from Eq.2 and insert in Eq.1, then make explicit in (solve for) [H3O+], and take the negative logarithm:

pH = - log [H3O+] = 6.37 - ½ log [Ni2+] ( = - ½ log (Kw2/Ks

- ½ log [Ni2+])

This is the minimum pH needed to keep a given concentration of Ni2+ in solution without any precipitation to occur:

[Ni2+] = 1 M … pH < 6.37
[Ni2+] = 0.1 M … pH < 6.87
[Ni2+] = 0.01 M … pH < 7.37
[Ni2+] = 0.001 M … pH < 7.87

For Cu(OH)2, Ks = 4.8 x 10-20, repeat the above and get:

pH = - log [H3O+] = 4.34 - ½ log [Cu2+] ( = - ½ log (Kw2/Ks

No surprise there: Cu would start precipitating first, from lower pH values because Cu(OH)2 is much less soluble. This is quite similar to the separation of Cu and Ni based on difference in solubility between their sulphides.

Complexation in these conditions can be neglected because [NH3] ≈ 0

Sedit, there’s gotta be something wrong there.

Edit:

Oookaaayy: long calculation but in these photos there's no copper whatsoever, if you ask me!

m1tanker78 - 22-11-2011 at 07:21

Oookaaayy: long calculation but in these photos there's no copper whatsoever, if you ask me!

My thought exactly. If he pre-refined the chloride then I guess it makes sense. If not, he has some copper hunting to do!

EDIT:

http://en.wikipedia.org/wiki/Dicopper_chloride_trihydroxide

Quote:

Cu2(OH)3Cl can be prepared by hydrolysis of a CuCl2 solution at pH 4 ~7. A variety of bases such as sodium carbonate, ammonium, calcium, or sodium hydroxide may be used (eq. 3).[1]

2CuCl2 + 3 NaOH → Cu2(OH)3Cl + 3 NaCl (eq.3)

Cu2(OH)3Cl can also be prepared by the reaction of a hot CuCl2 solution with freshly precipitated CuO (eq. 4).

CuCl2 + 3 CuO + 3 H2O → 2 Cu2(OH)3Cl (eq.4)

If sufficient chloride ions are present in solution, hydrolysis of CuSO4 with alkali also produces Cu2(OH)3Cl (eq. 5).

2 CuSO4 + 3 NaOH + NaCl → 2 Cu2(OH)3Cl + 2 Na2SO4 (eq.5)

It's insoluble to boot. I wonder if this compound has been overlooked?

Tank

[Edited on 11-22-2011 by m1tanker78]

blogfast25 - 22-11-2011 at 08:27

Tank:

The one thing that could 'mask' the copper is a lot of chloride: chlorocuprate anions (CuCl42-

then form and these are emerald green. But it needs quite a lot of Cl-. My solutions (above) are in fact mixed CuCl2/NiCl2 and they're blue. The complexation constant of the chlorocuprate complex is fairly small.

No, Cu2(OH)3Cl hasn't really been overlooked: it's just that in borderline conditions, Cu2+ + 2 OH- === > Cu(OH)2 is a bit of an oversimplification: the hydrolysis of Cu2+ happens in stages. At higher pH the Cu2(OH)3Cl reverts to Cu(OH)2 'proper'.

Imagine it like this:

Cu(H2O)n2+ + H2O === > Cu(OH)(H2O)n-1+ + H3O+

And again. With Cl- these subspecies can then form hydroxy chlorides, like the one you mentioned and these are often insoluble. Fe3+ does it too, more so even due to the higher charge and stronger associated central electrical field, which leads to de-protonation of the water cloud...

I'm afraid this is not a day of 'break throughs'

There remains one possibility: if he's saturated a Cu2+/Ni2+ solution with chloride, then most copper will be present as green chlorocuprate complex anions, indistiguishable from Ni2+. That would offer slight protection against hydrolysis. But add NaOH to a solution containing chlorocuprate and Cu(OH)2 drops out all the same: the complex isn't strong enough and Cu(OH)2 (or copper (II) hydroxy chlorides) too insoluble. Perhaps with weak NH3? Doesn't sound convincing to me.

I'm gonna check this out... Is his green precipitate in this case a copper hydroxy chloride?

[Edited on 22-11-2011 by blogfast25]

Sedit - 22-11-2011 at 09:56

I really don't know what to make of it yet honestly. It shocked me to see so much precipitate without the appearance of the amine complex.

The filtrate did have a slight blue tinge to it but that could have been a play on light I'm not sure.

I washed the precipitate with cold clean water a couple times then just added HCl to gain a dark green solution I'm boiling down. There is a slight red tinge to the reduced product after the addition of Aluminum so there is some chance that Copper made its way in. This could just be an artifact of poor washing however or it could be something more only further experimentation will tell.

I'm going to go do some more work and see where it gets me, be back in a few hours with hopefully more pictures to show.

[edit]

Quote:

My thought exactly. If he pre-refined the chloride then I guess it makes sense. If not, he has some copper hunting to do!

The solution I filtered off of the precipitate I slowly added more Ammonia to expecting to recover a small amount of extra hydroxide... instead it was just a rapid copper amine complex...

Don't stress it, im as confused as yall are at the moment. This was not what I was expecting at the start of the experiment.

[Edited on 22-11-2011 by Sedit]

blogfast25 - 22-11-2011 at 10:45

I dispensed about 20 ml of 36 % HCl in a beaker and started dissolving the 1:1 Cu:Ni mixed hydroxide into it. At first the solution was pretty green, due to the excess chloride, then it slowly became more blue. And at some point during the additions a blue-green, sandy (not gelatinous) precipitate formed. Stirring didn’t dissolve it. This is almost certainly Cu2(OH)3Cl (or Cu(OH)2.Cu(OH)Cl, rewritten) It looks very much like ‘Verdigris’, basic copper carbonate (CuCO3.Cu(OH)2). Adding a little more HCl dissolved it (hence no photo).

Side note: if the pH at which this hydroxy chloride forms is low enough, separation between Cu2+ and Ni2+ should be possible that way. Will check that tomorrow.

Some more mixed hydroxide was the added until the blueish solution below was obtained:

In that I dissolved a good dollop of NaCl and the green of chlorocuprate anions reappeared:

To this I added weak ammonia in a test tube:

This is the same phenomenon as before: interface lack of homogeneity/equilibrium. That precipitate is gelatinous and greener than the earlier observed precipitate.

But it should be possible to precipitate Cu2(OH)3Cl from this solution with careful addition of NaOH…

[Edited on 22-11-2011 by blogfast25]

m1tanker78 - 22-11-2011 at 17:11

I'm gonna check this out... Is his green precipitate in this case a copper hydroxy chloride?

He mentioned this...

Quote:

The key seems to be to loss a small amount of nickle by keeping the solution slightly acidic.

...which is what aroused my suspicion that this could be the hydroxy chloride of copper (among other things). I skimmed through your equations and your experimental findings (will read more judiciously later). Your last experiment seems to reinforce the fact that the ppt only exists in a narrow [slightly acidic] pH range.

Maybe this helps to explain one of my previous observations that I dug up from pg. 3 of this thread. I plated nickel from what was supposed to be the copper-ammonia complex. I later discovered that by changing the current density, an inferior copper deposit could be obtained as well. (edit) In retrospect, it could have been that Ni2+ was depleted - not so much to do with changing the current density.

Contrary to what has been proposed and reported before, the experiment clearly demonstrates that nickel ions are present, possibly in abundance, in the ammonia complex. The proportion of copper ions (if any) of the same is unclear as of yet.
[...]

Tank

[Edited on 11-23-2011 by m1tanker78]

Sedit - 22-11-2011 at 20:29

Its copper precipitate of some sorts even though at first sight it had me convinced it was no doubt some form of Nickle precipitate. I established this by turning the precipitate into its chloride and then adding Aluminum.

As the hydrogen rapidly gassed off I held a flame over it and it gave burst of blue flames where as to the best of my knowledge Nickle should give no color in the flame test.

Sorry for getting people hopes up but you can see from the photos where my excitement came from,that precipitate looks exactly like Nickle hydroxide and even though the reduced product after my little flame experiment proved there was some Nickle in there it also showed heavy Copper contamination.

I have the chloride solution drying as we speak so when its dry and crystallized it will make it much easier to determine the true contents of the large amount of green precipitate.

I vow that this is not over yet and I will conquer this one but curve balls and the unexpected seem to occur around every corner

PS: Blogfast when I post my pictures and it ask what program to use to open them just use Internet explorer it will open them with ease. Either that or just add a JPG extension to the file.

[Edited on 23-11-2011 by Sedit]

What a sick sick world we live in with this hobby.

The wise chemist thought to himself " I know, I will just sprinkle in some NaOH and the precipitated hydroxide from that should tell me once and for all if its copper of some kind or Nickle of some kind...."

So the chemist did just that expecting to see either blue or green precipitate.... Instead the chemist found himself with an opaque brown solution. :mad:

As if this whole set of nonsense couldn't get much stranger to the chemist this happens.

Pictures in the morning its late now. ::Sigh::

[Edited on 23-11-2011 by Sedit]

blogfast25 - 23-11-2011 at 09:59

Sedit, thanks for the tip on reading your pic files.

I added 1 M NaOH dropwise to the green solution (CuCl2/NiCl2/NaCl, pH about 3) above, from a burette. The first precipitate kept redissolving on stirring but from about pH = 4 (paper) permanent precipitation occurred. It took quite a bit of patience to carefully get to pH = 5 but precipitate kept on forming. Due to a mishap I then added slightly too much and stopped: It looked like this by then:

This was filtered and the filtrate was clear and colourless. The pH meter gave pH = 7.9 and not surprisingly no further precipitation occurred on adding more NaOH: the nickel had precipitated too, showing that as a separation method, if at all possible, this would require very accurate pH control, probably buffering. (And beware of possibly irreversible co-precipitation of Ni(OH)2!)

This precipitate was profusely washed with hot DIW and the last wash water then tested for chloride with lead acetate solution (silver nitrate would be better but I haven’t got any). No precipitation was observed (see left tube, bottom pic). This is the filter cake after washing, slurried in a bit of DIW:

1 M NaOH was then added to a slurry of the filtercake, it turned blue and more fluffy, presumably because of conversion of the copper hydroxychloride to copper hydroxide, liberating the chloride as NaCl (pH was now >> 7). It was heated slightly:

This was filtered and the first filtrate intercepted, then acidified with acetic acid to pH ≈ 4 (paper) and strong PbAc2 added. Turbidity developed immediately, see right hand tube:

To be honest I had expected a stronger precipitation of lead chloride, as this would be stronger evidence that the original filter cake contained chemically bound chloride. But pure Cu2(OH)3Cl only contains about 17 % Cl and the precipitate also contains Ni(OH)2 of course and there’s lot of water too. I may repeat this experiment in more controlled conditions, and without Ni present, with dried Cu2(OH)3Cl.

All the green precipitates we've seen in the above posts are likely to be copper hydroxychloride. Tank was right: we really did overlook the role chloride seems to play here...

[Edited on 23-11-2011 by blogfast25]

Sedit - 23-11-2011 at 13:11

Ok Pictures coming soon but the kids just got home from school and I have to take care of the important things in life before getting pictures...

I dried the Chloride made from that green precipitate you see in the large jar above leaving me with a green powder.

I dissolved this in a minimum amount of water and added NaOH prills like I stated last night only to yeild a brown precipitate that is the color of MnO. This I don't understand at all.

Ok, today with my experiment I again dissolved the green powder in a minimum amount of water as before and added a NaOH solution this time. This time I got a blue precipitate then with a few more drops of NaOH solution the precipitate turned green like that of before. But here's where something odd took place... The solution above the precipitate took on a blue color exactly like that of the Ammonia complex. I feel this is due to Ammonia being released from Ammonium Chloride that was still in the mix but there is the possibility that the precipitate itself IS a Nickle amine complex (or Copper) and that's why things are so weird with this reaction.

I suspected the Cl- ion played a roll in the past hence the reason I asked Blogfast to add NH4Cl in hopes that would produce some sort of results but it didn't.

I'm starting to wounder if the Nickle amine complex is not very soluble in a solution of the Copper amine complex.

I'm just musing at the moment since I just performed the experiment and got alot of other things to do right this minute so keep the flames at a minimum. Later tonight I will get the pictures and further this discussion.

blogfast25 - 23-11-2011 at 14:10

I dissolved this in a minimum amount of water and added NaOH prills like I stated last night only to yeild a brown precipitate that is the color of MnO. This I don't understand at all.

[…]

The solution above the precipitate took on a blue color exactly like that of the Ammonia complex. I feel this is due to Ammonia being released from Ammonium Chloride that was still in the mix but there is the possibility that the precipitate itself IS a Nickle amine complex (or Copper) and that's why things are so weird with this reaction.

[…]

I'm starting to wounder if the Nickle amine complex is not very soluble in a solution of the Copper amine complex.

Adding prills of NaOH leads to local overheating due to the high solvation heat of NaOH: Cu(OH)2 loses water to form (brown black) CuO easily.

Second point: could be residual ammonia (did you smell any?) but more likely it’s cuprate anions (Cu(OH)42- from Cu(OH)2 + 2 NaOH); copper is slightly amphoteric and Cu(OH)2 dissolves somewhat in strong NaOH to a cobalt blue cuprate, almost indistinguishable from the corresponding ammine complex.

Third point: no. Remember how a 50/50 Cu(OH)2/Ni(OH)2 mix dissolved effortlessly in 1.5 M NH3. The nickel ammine complex’ solubility is unaffected by the copper ammine complex.

Try using fairly dilute, cool alkaline solutions, like 1 M (40 g NaOH per liter).

Sedit - 23-11-2011 at 14:55

I believe you are correct and that is Copper oxides due to over heating, the blue precipitate in the background is the one where I use dilute NaOH.

The color of the solution cleared up so I will repeat the experiment in an attempt to make sure it stays so I can photograph it.

Its odd as you can see the precipitate is blue but when I first started to add the NaOH solution the precipitate was green and only turned blue after further addition of NaOH.

Im going to have to make a standard NaOH solution so I can determine the exact concentration at any point in time as this is important. Here are the photos. I will perform more experiments later. Its dinner time now

WTF, I cant attach files, is it showing up on yalls computers because I don't see it on mine

[Edited on 23-11-2011 by Sedit]

[Edited on 23-11-2011 by Sedit]

[Edited on 24-11-2011 by Sedit]

blogfast25 - 24-11-2011 at 06:24

As far as I’m concerned the mystery is now solved, and it really is ‘game over’.

I distinctly remember when I first precipitated the mixed hydroxide that it was this verdigris green/blue colour and that on adding the requisite small excess of NaOH the stuff turned blue. This is likely to happen anytime there’s sufficient chloride present:

2 CuCl2(aq) + 3 NaOH(aq) === > Cu2(OH)3Cl(s) + 3 NaCl(aq)

On adding some more NaOH the transition from verdigris to blue occurs:

Cu2(OH)3Cl(s) + NaOH(aq) === > 2 Cu(OH)2(s) + NaCl(aq)

As I wrote above, I feel all the green precipitates including the test tube ones we’ve seen are in fact the copper hydroxychloride, probably with co-precipitated Ni(OH)2.nH2O in it. No real separation was ever achieved, not even at the interface.

The copper hydoxychloride appears to start precipitating a bit earlier (from about pH = 4) than straight copper hydroxide but its precipitation range (4 to 7 acc. Wiki, in accordance with my observations) overlaps too much with that of nickel hydroxyde (which starts from about 6.5) to be of any real use as a separation method for Cu2+ and Ni2+. The only conditions in which it might be possible to separate out any nickel is if it is present in really low concentrations (< 0.001 M) to begin with.

It’s been fun and educational…

No, I can't see any pics.

m1tanker78 - 24-11-2011 at 06:49

It’s been fun and educational…

Yes it has!

Sedit, I don't see any pics either.

Tank

Sedit - 24-11-2011 at 11:57

Trying once again to attach the photo...

Somethings wrong folks, can someone else attempt to attach a file of somekind here to see if its the forum or something odd on my end.

[Edited on 24-11-2011 by Sedit]

m1tanker78 - 24-11-2011 at 20:21

Trying once again to attach the photo...

I'm telling ya, create a Photobucket or similar account. Upload pics. Some allow you to edit right in the browser. Copy the url it gives you. Paste it here (between the img tags). Presto.

If you're hell-bent on uploading the pics directly then search the net for an open source program that will allow you to resize and crop - at minimum.

Has anyone tried precipitating copper from the mixed chlorides with ascorbic acid? I tried yesterday to add AscA solution to dilute-ish Ni/Cu chloride but nothing precipitated.

I've read a few accounts of people using AscA to precipitate colloidal copper from a Cu2+ solution but no mention of other metals (nickel in this case). What am I doing wrong?

Tank

Sedit - 27-11-2011 at 21:46

Photobucket and the likes are the reason forums look like shit as the years go by because the images vanish, I do not like my photos vanishing because my memory is very very bad. So bad that I can't even remember why I have been trying to upload this photo I am about to attach without rereading the last page or so of this thread but here it is and now I got a 30 day trial version of Adobe photoshop (could use a crack if yall got one

) I will post as many pictures as I can.

PS. The acidic precipitate (The green shit in the big jar higher up in this thread)was purely Copper complex of some sorts with no magnetic ability after reduction with Zinc or Aluminum, If I find that there is no Nickle at all in the mix or at lest a very low percentage then that might be a viable option for removing the Cu from the mother liquor.

Sorry for the hackish post but I just had to get this one out there as I have been away and trying to fix up my computer at the same time.

Formatik - 28-12-2011 at 00:11

I don't know if this method has been mentioned yet, but I was reading some time ago in Die Bestimmungsmethoden des Nickels und Kobalts und ihre Trennung von den anderen Elementen, which mentioned a reference of Brunck who suggests to separate copper from nickel by precipitating the copper through sodium hydrosulfite (sodium dithionite) in acid solution.

Taking a closer look at the Brunck reference (Ann. Chem. 327 (1903) 244), who says that the separation is quantitative. Brunck points out iron, zinc, nickel and cobalt are not precipitated through sodium hydrosulfite in acidic solution (but copper is). Zinc is said also to show not even the smallest tendency to go into the copper precipitate (not the case when precipitation is done with H2S).

blogfast25 - 28-12-2011 at 09:33

Quote: Originally posted by Formatik

Zinc is said also to show not even the smallest tendency to go into the copper precipitate (not the case when precipitation is done with H2S).

With a solubility product Ksp for ZnS of about 10-25 as opposed to 10-36 for CuS, zinc and copper can almost certainly be separated from each other based on solubility of their sulphides (using H2S saturation) in acidic solutions, in the same way that it's done for Cu/Ni separation by sulphides at pH < 7. The Ksp for NiS is similar to that of ZnS.

Formatik - 29-12-2011 at 09:10

This comment on zinc was Brunck's finding on a partial co-precipitation he apparently observed.

Brunck, Ann. Chem. 327 (1903) 244.

m1tanker78 - 13-5-2012 at 07:49

I scaled up the electrolysis experiment again only this time I used a piece of cupro-nickel I found at a marine scrap yard. I hit a couple of minor snags (fell asleep as always) so this run wasn't as clean as the previous one. By that I mean that I wasn't particularly careful to remove as much copper as possible. I stopped the electrolysis and filtered the cell liquor to remove elemental copper and copper(I) chloride (the other snag!).

I believe I devised a simple qualitative test for copper(II) in the cell liquor. In a tiny test tube, I placed some filtered and dilute cell liquor. I added about the same volume of DMSO and added some granules of NaHSO4. With light heating, a dark precipitate immediately formed which I presume to be CuS.

I don't know if this method would work (be cost effective, etc) for bulk separation but my focus is on a reliable and simiple qualitative test for copper in the nickel(II) chloride cell liquor. FWIW, the test solution smells like DMS, not H2S. After sitting for a bit, some small flat colorless crystals have manifested at the bottom of the test tube.

Could somebody replicate this with the pure chlorides of nickel and copper and perhaps one 50/50?

CuS(?) can be seen in the test tube:

Tank

AJKOER - 16-5-2012 at 10:30

Those following some of the chemistry surrounding Oxalic acid and Oxalates that I have been reporting may find some interesting observations for the separation of Cu/Ni. In particular :

1. Per "A dictionary of chemistry and the allied branches of other sciences", Volume 4, by Henry Watts, page 257, to quote:
"Oxalate Of Copper. Neutral cupric Oxalate, ....(according to Lowe, Jahresb. 1860, p. 213), is a light greenish-blue precipitate, insoluble in water, nearly or quite insoluble in oxalic acid, but easily soluble in the neutral oxalates of ammonium, potassium and sodium. It does not give off the whole of its water even at 120°, but decomposes at a somewhat higher temperature."
Link:
http://books.google.com/books?pg=PA257&id=lYXPAAAAMAAJ#v...

2. Same source, page 262, to quote:
"Oxalate Of Nickel, ... Greenish-white precipitate insoluble in water, soluble in ammonia and in ammoniacal salts. It dissolves also in potash, forming a crystallisable potassio-nickel-oxalate. Neutral oxalate of ammonium dissolves oxalate of nickel, and the solution yields by evaporation green prisms of ammonio-nickel-oxalate. On adding to the aqueous solution of this bait a small quantity of ammonia, a pale green precipitate is 'formed, consisting, according to Winckelblech (Ann. Ch. Pharm. xiii. 278), of oxalate of nickel and nickel-ammonium"

3. Source: "Zinc, cadmium, mercury, bismuth, tin, antimony, arsenic, nickel, cobalt ...", by Carl Schnabel, page 578, to quote:

"The powdered ore is heated with concentrated oxalic acid solution, which leaves nickel undissolved. The residue is reduced to obtain nickel." Link:
http://books.google.com/books?id=DWx9AAAAIAAJ&pg=PA578&a...

Suggested Synthesis to Extract Nickel
Dissolve the Cu/Ni in a concentrated Oxalic acid solution. Add a solution of Sodium oxalate to dissolve the Copper oxalate precipitate (I am expecting this to work as a soluble Nickel-Sodium oxalate double salt apparently does not exist, so Nickel Oxalate will not dissolve). Separate the residue which should be Nickel, some Nickel oxalate and other impurities for further thermal processing.

I am expecting the creation of some Nickel oxalate as apparently Oxalic acid with mild heating dissolves Nickel Oxide and also can cause the dissolution of Nickel in Nickel ferrite. References:

"Dissolution of Nickel Oxide in Oxalic Acid Aqueous Solutions"
Per the abstract:
"The dissolution of nickel oxide (bunsenite) in acid solutions containing oxalic acid has been studied at 70.0°C. The dependencies of the rate of dissolution on total oxalate concentration and on pH have been explained by assuming a mechanism involving the transfer of two different surface complexes, I and II, that predominate in different pH ranges. The rate law is R=k1{I}+k2{II}, where {} denotes surface concentration. The values k1Ns=3.04×10−3 mol Ni m−2 s−1, k2Ns=1.84×10−3 mol Ni m−2 s−1, together with the stability constants K1=675 mol−1 dm3 and K2=60 mol−1 dm3 fit all the results very well. The species formed in more acidic media is both more stable and more reactive."

LINK:
http://www.sciencedirect.com/science/article/pii/S0021979701...

"Dissolution of Nickel Ferrite in Aqueous Solutions Containing Oxalic Acid and Ferrous Salts"
Per the abstract:
"The dissolution of nickel ferrite in oxalic acid and in ferrous oxalate-oxalic acid aqueous solution was studied. Nickel ferrite was synthesized by thermal decomposition of a mixed tartrate; the particles were shown to be coated with a thin ferric oxide layer. Dissolution takes place in two stages, the first one corresponding to the dissolution of the ferric oxide outer layer and the second one being the dissolution of Ni(1.06)Fe(1.96)O(4). The kinetics of dissolution during this first stage is typical of ferric oxides: in oxalic acid, both a ligand-assisted and a redox mechanism operates, whereas in the presence of ferrous ions, redox catalysis leads to a faster dissolution. The rate dependence on both oxalic acid and on ferrous ion is described by the Langmuir-Hinshelwood equation; the best fitting corresponds to K(1)(ads)=25.6 mol(-1) dm(-3) and k(1)(max)=9.17x10(-7) mol m(-2) s(-1) and K(2)(ads)=37.1x10(3) mol(-1) dm(-3) and k(2)(max)=62.3x10(-7) mol m(-2) s(-1), respectively. In the second stage, Langmuir-Hinshelwood kinetics also describes the dissolution of iron and nickel from nickel ferrite, with K(1)(ads)=20.8 mol(-1) dm(3) and K(2)(ads)=1.16x10(5) mol(-1) dm(3). For iron, k(1)(max)=1.02x10(-7) mol of Fe m(-2) s(-1) and k(2)(max)=2.38x10(-7) mol of Fe m(-2) s(-1); for nickel, the rate constants k(1)(max) and k(2)(max) are 2.4 and 1.79 times smaller, respectively. The factor 1.79 agrees nicely with the stoichiometric ratio, whereas the factor 2.4 implies the accumulation of some nickel in the residual particles. The rate of nickel dissolution in oxalic acid is higher than that in bunsenite by a factor of 8, whereas hematite is more reactive by a factor of 9 (in the absence of Fe(II)) and 27 (in the presence of Fe (II)). It may be concluded that oxalic acid operates to dissolve iron, and the ensuing disruption of the solid framework accelerates the release of nickel." Link:
http://www.ncbi.nlm.nih.gov/pubmed/11254279

Note, as I am currently out of stock in H2C2O4, I am not immediately able to test this suggested synthesis.

[Edited on 16-5-2012 by AJKOER]

Formatik - 26-5-2012 at 15:09

Another way to do separation is through aqueous hydrazine. Starting from an acid mixture of copper and nickel nitrate. Diluting this with some water then adding portion-wise aqueous hydrazine. Hydrazine first attacks the Cu ion mainly reducing it to form copper. Then, about after that point it will form a purple lilac explosive complex (nickel nitrate hydrazinate). This residue is hazardous and possibly unstable due to copper contamination. The hydrazine addition should be done just up until that point at which the lilac salt starts appearing or just before this, otherwise it becomes difficult to judge and the nickel will have been used up (a colorless solution obviously means the reaction has gone too far). The brown of the copper makes this difficult to see. But the copper can easily be reduced out, then filter the solution and solutions can be obtained that do not have the faintest scent of hydrazine, and are pure green color.

m1tanker78 - 27-5-2012 at 07:05

Formatik, would you happen to have a reference for this?

Ajoker: I'm oxalic acid-poor as well. It'll have to wait.

Tank

Formatik - 27-5-2012 at 10:20

Formatik, would you happen to have a reference for this?

No reference here this time, it's from my own experimentation. Because the amounts weren't noted, I had to kind of feel out the reaction a few times. It is a quick route for impatient people.

I also doubt the suggested oxalic acid route above will work. The reference: Neueste Erfindungen und Erfahrungen auf den Gebieten der praktischen Technik, Elektrotechnik, der Gewerbe, Industrie, Chemie, der Land und Hauswirthschaft (1895), Vol.21, p. 353 speaks of a solution of nickel oxalate in potassium oxalate. So nickel oxalate will build double salts with alkalis.

AJKOER - 29-5-2012 at 08:03

Quote: Originally posted by Formatik

I also doubt the suggested oxalic acid route above will work. The reference: Neueste Erfindungen und Erfahrungen auf den Gebieten der praktischen Technik, Elektrotechnik, der Gewerbe, Industrie, Chemie, der Land und Hauswirthschaft (1895), Vol.21, p. 353 speaks of a solution of nickel oxalate in potassium oxalate. So nickel oxalate will build double salts with alkalis.

Per my cited source ("A dictionary of chemistry and the allied branches of other sciences", Volume 4, by Henry Watts, page 262) to quote again: "Oxalate Of Nickel, ... Greenish-white precipitate insoluble in water, soluble in ammonia and in ammoniacal salts. It dissolves also in potash, forming a crystallisable potassio-nickel-oxalate. Neutral oxalate of ammonium dissolves oxalate of nickel, and the solution yields by evaporation green prisms of ammonio-nickel-oxalate. " so Nickel oxalate apparently forms double salts with Potassium & Ammonia, which is in agreement with your source on the Potassium double salt.

Interestingly, a Sodium-Nickel double salt is not referenced nor does a general internet search find any such salt. However, there is an instance of Nickel plating in a Sodium oxalate / Oxalate acid medium (see http://www.tandfonline.com/doi/abs/10.1080/03602550500373881... ). If a double salt formation was possible, wouldn't it's formation interfere with the Ni plating process?

So my doubts are leaning more towards the fact that the Sodium double salt happens not to exist.

[Edited on 29-5-2012 by AJKOER]

Formatik - 29-5-2012 at 15:25

No, there's a reference which confirms it won't work since there is such a complex salt. Analytical Chemistry (1921), Vol. 1, p. 385 by F.P. Treadwell states nickel oxalate forms a soluble sodium nickel oxalate complex salt with sodium oxalate. The reference from Watts which was Winckelblech, never attempted to determine any interaction with sodium salts, so that's why it was never noted. It was not comprehensive.

AJKOER - 29-5-2012 at 19:04

Formatik:

You are correct. Found it in an online Googlebook. Source: "Analytical chemistry", Volume 1, by Frederick Pearson Treadwell, page 385 (as you noted).

Link:
http://books.google.com/books?id=ubuaqKkwsQoC&pg=PA385&a...
----------------------------------------------------------

The same source also provided me with another route to investigate. First, here is the authors comment on Copper related reactions on page 277:

"The proper solvent for copper is nitric acid:

3 Cu + 8 HNOs -> 3 Cu++ + 6 N03- + 4 H20 + 2 NO (g)

Bright copper is not dissolved by hydrochloric acid alone, but in the presence of a weak oxidizing agent, e.g., ferric chloride, the solution of the metal is easily effected. Hot hydrobromic acid dissolves it with evolution of hydrogen, forming cuprous hydrobromic acid: 2 Cu + 6 HBr <=> H4[Cu2Br6] + H2 (g)

At the beginning of this reaction the solution usually turns dark violet on account of the formation of the cupric salt of cuprous hydrobromic acid, owing to the copper being somewhat oxidized on the surface. In this case, however, the solution soon becomes colorless, owing to the reduction of the cupric salt by metallic copper. On adding water to the clear solution cuprous bromide is precipitated: [Cu2Br6]-- ==> Cu2Br2 + 4 Br-.

Copper is not attacked by dilute sulfuric acid, but it dissolves in hot concentrated sulfuric acid, forming cupric sulfate with evolution of sulfur dioxide:

Cu + 2 H2S04 -> CuS04 + 2 H20 + S02 (g)

The behavior of copper toward acids can be understood by reference to the electromotive series (p. 41). As copper is below hydrogen in the series it can be oxidized by hydrogen ions only when the concentration of cupric ions is kept very low (cf. p. 43). Hydrobromic acid dissolves copper because a slightly ionized complex is formed. Sulfuric acid dissolves copper by virtue of the oxidizing power of the hexavalent sulfur."

So, hot HBr dissolves Cu owing to the formation of a slightly ionized complex. Interestingly, per another source with respect to Nickel, to quote:

"Due to the corrosive nature of halogenated gases, special attention should be paid to their handling. High purity WF6 and HBr are supplied typically in stainless steel, pure nickel or nickel lined cylinders. Passive films formed on the two surfaces are different."

Link: http://www.smlassociates.com/cormetal.shtml

So HBr appears to be unreactive with respect to Nickel owing to the formation of a passive film.

So assuming the existence of the Cu/Ni in the form of an alloy does not presence any issues (?), an idea that might be interesting to investigate is:

> Place the Cu/Ni coin (or first reduce the coin to a fine powder) in hot HBr and see if the Ni remains largely undissolved. I would also collect the H2 gas, as per the reaction:

2 Cu + 6 HBr <=> H4[Cu2Br6] + H2 (g)

for each mole of Hydrogen, two moles of Copper have been consumed. Knowing the observed volumes of H2 produced and starting weight of the coin and the residual undissolved mass, one can check on the progress of the reaction. Also, from the filtered solution upon adding water, per the author's cited reaction:

[Cu2Br6]-- ==> Cu2Br2 (s)+ 4 Br-

the weight of the cuprous bromide precipitated also speaks to the amount of Copper consumed.

[Edited on 30-5-2012 by AJKOER]

AJKOER - 12-8-2012 at 05:48

Here is an extract from an old recipe from "Hand-book of chemistry", Volume 5, page 357 by Leopold Gmelin that mentions some of the steps noted above to separate out Nickel:

"5. Proust heats roasted copper-nickel with dilute sulphuric acid, and adds carbonate of potash to the filtered solution, to precipitate arseniate of ferric oxide, till iron can no longer be detected in the liquid by ferrocyanide of potassium. The liquid is then filtered again, and sulphuretted hydrogen passed through it, to precipitate arsenic, copper, and bismuth, till it is so far saturated as to retain the odour of the gas after being kept for 24 hours in a closed vessel. The liquid, once more filtered and then evaporated, yields crystals chiefly consisting of sulphate of nickel-oxide and potash, while the cobalt-salt, for the most part, remains in solution; the former is repeatedly dissolved and recrystallized to free it from adhering cobalt-salt, and its solution afterwards treated with carbonate of potash, which precipitates the nickel in the form of carbonate."

Also, per page 355:

"The copper-nickel or the cobalt-speiss is generally roasted in a state of powder (at a gentle heat at first, to prevent it from baking together), whereby the greater part of the arsenic is removed, the nickel oxidated, and a saving of nitric acid thus effected. Since, however, the roasting process leaves a portion of the arsenic combined in the form of arsenic acid with the oxide of nickel, the roasted ore must be several times intimately mixed with charcoal dust and again roasted, as long as vapours of arsenic continue to be evolved. Erdmann moistens the roasted cobalt speiss with water and places it in a cellar till it is converted into hydrate; it is thereby rendered more easily soluble.
1. Laugier dissolves the roasted copper-nickel or the speiss in nitric acid, passes sulphuretted hydrogen through the dilute acid solution till all the arsenic, copper, bismuth, and antimony are precipitated—then falters
—precipitates all the iron, cobalt, and nickel with carbonate of soda— washes the precipitate thoroughly, and treats it first with oxalic acid and then with ammonia, as described on page 319, repeating the solution of the nickel-oxalate in aqueous ammonia, till the liquid which stands above the resulting precipitate no longer exhibits a rose-colour, and is almost wholly free from cobalt."

Link: http://books.google.com/books?pg=PA357&lpg=PA32&dq=h...

[Edited on 12-8-2012 by AJKOER]

cyanureeves - 12-8-2012 at 06:08

this is very similar to what has been done except potassium carbonate is here used and it was also mentioned by someone here that maybe sodium metabisullfate could be used to drop nickel. sulphuretted hydrogen and sodium metabisulfate both can be used to drop gold from auric chloride i think. the one thing that is very different here is the roasting in sulfuric acid, dissolving in sulfuric acid has been done but not roasting. its interesting how in this case hydrogen sulfate is used to drop the copper first and not the nickel.

m1tanker78 - 21-8-2012 at 19:05

Would somebody please demonstrate separation by bubbling H2S gas through a solution of the combined metals. Seems easy enough for the chemist who has the reagents, equipment, knowledge and the desire.

Tank

tetrahedron - 24-10-2012 at 00:58

@tank: this thread's OP is looking to obtain pure nickel salts, but you admit it's hard to achieve:

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

i attempted electrorefining Cu/Ni in a similar way to yours. using a CuSO₄ electrolyte, for the anode i soldered one side of the coins to the tip of some insulated Cu wire individually, then i spray painted the whole side, giving me about 15cm³ total surface area on the front, for 15g Cu/Ni (3:1). the cathode was a Cu wire. i don't remember the current but it was pretty low (around 0.5A). i stopped the experiment (total duration <24h) when i noticed the electrolyte seeping under the clear paint. the electrodeposited Cu looked the same as if a pure Cu anode had been used.

The metal powders are too intertwined and there may have been some coprecipitation too, making good separation difficult. It’s a recipe for very impure Ni and very impure Cu, both of which would then need refining anyway.

That's precisely what I observed. I was trying to get around having to scrape the Copper(I) Oxide (and/or hydride?) off the cathode every hour or so. Aside from that, the electrolytic separation of the metals seems to work well. I'm working on starting another batch as I type...

have you tried boiling these mixed powders in the electrolyte to reduce the last bits of copper in solution by displacing it with nickel, and and get pure NiCl₂? the Cu metal already precipitated should play no role. i had no mixed powders, so i used some fresh coins for the metathesis (this has severe limitations due to the small surface area). soon they turned a rusty color. the powder method should be pretty efficient. in industry, copper impurities in nickel electrorefining are removed with (pure) nickel powder as well.

blogfast25 - 25-10-2012 at 05:51

Would somebody please demonstrate separation by bubbling H2S gas through a solution of the combined metals. Seems easy enough for the chemist who has the reagents, equipment, knowledge and the desire.

Tank

As this is a bit of a ‘classic’ of older, wet analytical chemistry, try and consult older texts on this, for full details.

In practice I think you’d have to buffer the solution to a fairly low pH (say 4), then saturate it with H2S gas: CuS precipitates, Ni2+ stays in solution. Then flush out most of the H2S with air and separate the metals by filtration.

Not difficult in principle but it will require scrubbing out the highly toxic H2S very effectively and decent filtration media. Commercial bleach is a decent H2S scrubber because it oxidises the H2S to elemental sulphur and water. But use at least two wash bottles.

cyanureeves - 25-10-2012 at 16:44

is this H2S the same gas that gas line workers watch out for with beeper detectors? isnt it as poisonous as cyanide? any how does this H2S work only on cupronickel in HCl or would it work with cupronickel in sulfuric acid? how is this poison gas made in the lab?

tetrahedron - 25-10-2012 at 17:16

i'll answer your least loaded question

from iron(II) sulfide and hydrochloric acid

how is this poison gas made in the lab?

Quote:

(the whole quote box is a link..pretty cool huh)

elementcollector1 - 25-10-2012 at 19:19

Reported spam above.

m1tanker78 - 25-10-2012 at 21:15

@tetrahedron:

What type of power supply did you use? Seems like you succeeded in plating out some copper. Did you test the cell liquor afterward (qualitative test for nickel)?

Quote:

have you tried boiling these mixed powders in the electrolyte to reduce the last bits of copper in solution by displacing it with nickel, and and get pure NiCl2?

I haven't but wouldn't be difficult to do. I have a crop of copper powder that's contaminated with nickel powder from my last problem-plagued run. I still have the cell liquor which is saturated nickel chloride contaminated with some copper chloride. I could kick the metals out, rinse well and add in more acid. How acidic should it be? I believe that both metals would dissolve especially being powder form. However, your suggestion might help explain some strange behaviors I observed in the cell. Gears are turning but need WD-40.

Quote:

OP is looking to obtain pure nickel salts, but you admit it's hard to achieve:

I made that comment in the context of maintaining an electroplating bath. They use pure nickel anodes to replenish the bath and remove some other offending metals (copper being one).

@cyanureeves:
You're the man for the task. Just be careful around H2S.

Tank

tetrahedron - 26-10-2012 at 03:32

What type of power supply did you use? Seems like you succeeded in plating out some copper. Did you test the cell liquor afterward (qualitative test for nickel)?

a lab power supply. i put all the data i remember in my previous post. the qualitative test for nickel was colorimetric. the electrolyte turned from blue to green, but not as green as it should be. i'm pretty sure i didn't let enough moles of e^- through, so i shouldn't complain that it didn't go to completion =D

I made that comment in the context of maintaining an electroplating bath. They use pure nickel anodes to replenish the bath and remove some other offending metals (copper being one).

assuming no side reactions at the electrodes, we have (here i use a molar ratio of 3:1 for Cu:Ni for simplicity, although it should be a mass ratio)

at the anode Ni + 3 Cu ---> Ni²⁺ + 3 Cu²⁺ + 8 e^-

at the cathode 4 Cu²⁺ + 8 e^- ---> 4 Cu

i.e. in the electrolyte the total reaction is

4 Cu²⁺ --(8 e^-)-> Ni²⁺ + 3 Cu²⁺

or simply

Cu²⁺ --(8 e^-)-> Ni²⁺

slower than using a pure Ni anode, but the outcome is the same (of course the newly oxidized copper ions won't diffuse to the cathode fast enough to ensure 100% removal; a divided cell with a pure nickel anode is preferable in order to remove most of the remaining Cu²⁺ at the very end).

How acidic should it be? I believe that both metals would dissolve especially being powder form. However, your suggestion might help explain some strange behaviors I observed in the cell.

i used no acid at all since i didn't care how well the Cu plated out..saturated CuSO₄ only

[Edited on 26-10-2012 by tetrahedron]

blogfast25 - 26-10-2012 at 09:49

is this H2S the same gas that gas line workers watch out for with beeper detectors? isnt it as poisonous as cyanide? any how does this H2S work only on cupronickel in HCl or would it work with cupronickel in sulfuric acid? how is this poison gas made in the lab?

When ‘cupronickel’ is dissolved in acid you obtain a solution of Cu(II) and Ni(II) salts.

When saturating an acidic solution of Cu2+ and Ni2+ with H2S (yes, it is more toxic than HCN –find a related thread by ‘Fluke’ on H2S poisoning on this forum) the exceedingly low solubility product (Ks) of CuS is reached and CuS precipitates. The solubility product of NiS is much higher and isn’t reached and thus NiS doesn’t precipitate (although NiS is insoluble, the concentration of S2-is very low in such solutions).

H2S is made from FeS + acid or Al2S3 + water.

If you have no prior experience working with significant amounts of H2S then stay away from it: H2S poisoning is extremely nasty.

[Edited on 26-10-2012 by blogfast25]

Hexavalent - 26-10-2012 at 10:40

find a related thread by ‘Fluke’ on H2S poisoning on this forum

I believe his name is 'Klute', and his thread can be found here;

http://www.sciencemadness.org/whisper/viewthread.php?tid=104...

[Edited on 26-10-2012 by Hexavalent]

blogfast25 - 26-10-2012 at 11:13

Quote: Originally posted by Hexavalent

find a related thread by ‘Fluke’ on H2S poisoning on this forum

I believe his name is 'Klute', and his thread can be found here;

http://www.sciencemadness.org/whisper/viewthread.php?tid=104...

[Edited on 26-10-2012 by Hexavalent]

Oooops - 'Klute', indeed. My bad. Thanks for looking up the thread, hexa.

cyanureeves - 26-10-2012 at 14:38

brrrrrrrr! i'm not going anywhere near H2S for anything but it would be so much fun to see nickel separated whether in solution or solid.

blogfast25 - 27-10-2012 at 06:00

brrrrrrrr! i'm not going anywhere near H2S for anything but it would be so much fun to see nickel separated whether in solution or solid.

The H2S method is (was) really intended for quantitative lab separations on a gram (or so) scale. Analytical work, not 'production'...

Go upthread for some interesting work on separating quite large amounts of Ni and Cu in solution.

cyanureeves - 27-10-2012 at 06:51

i have followed this since it began almost and it turned into a few different ways of separating the nickel. it is almost like the hydrazine thread and i think tanker actually reduced it to the point of being able to pick up the nickel left over with a magnet.i would like to see this procedure condensed because at certain points the process was stopped to answer questions about certain procedures. multiple people were working on this thing and some were looking for nickel some for just copper and as i was following i took bits from one or the other member.by the time i was finished i ended up with a 1974 nickel again, heads tails and all.

m1tanker78 - 27-10-2012 at 09:33

The H2S method is (was) really intended for quantitative lab separations on a gram (or so) scale. Analytical work, not 'production'...

I agree. My aim is to quantify the impurities of the metals after bulk separation by electrolysis. I'd love to put the icing on the cake, so to speak.