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Author: Subject: Separation of a US nickel
tetrahedron
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[*] posted on 25-10-2012 at 17:16


i'll answer your least loaded question
Quote: Originally posted by cyanureeves  
how is this poison gas made in the lab?

Quote:
from iron(II) sulfide and hydrochloric acid

(the whole quote box is a link..pretty cool huh)
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[*] posted on 25-10-2012 at 19:19


Reported spam above.



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[*] posted on 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. :D

Tank





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[*] posted on 26-10-2012 at 03:32


Quote: Originally posted by m1tanker78  
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
Quote: Originally posted by m1tanker78  
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).
Quote: Originally posted by m1tanker78  
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]
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[*] posted on 26-10-2012 at 09:49


Quote: Originally posted by cyanureeves  
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]




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[*] posted on 26-10-2012 at 10:40


Quote: Originally posted by blogfast25  
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]




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[*] posted on 26-10-2012 at 11:13


Quote: Originally posted by Hexavalent  
Quote: Originally posted by blogfast25  
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.




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[*] posted on 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.
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[*] posted on 27-10-2012 at 06:00


Quote: Originally posted by cyanureeves  
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.




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[*] posted on 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.
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[*] posted on 27-10-2012 at 09:33


Quote: Originally posted by blogfast25  

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.

Quote: Originally posted by cyanureeves  

[...]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

Which procedure? If you're referring to the electrolytic procedure, my eye tells me I achieved nearly complete separation. Simple tests (magnet, colors) so far are in agreement but I need some solid reinforcement with hard data instead of relying on educated guesses. The nickel powder I recover from this resembles iron filings in that it's strongly attracted to a magnet.

I'll condense the procedure and start a separate thread so that it's easier to reference and/or critique.

Tank




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[*] posted on 27-10-2012 at 11:38


sounds great to me.
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[*] posted on 15-2-2013 at 04:45


It was mentioned earlier, but can SO2 be bubbled through the solution to get the copper out of solution? Anyone tried that?
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[*] posted on 15-2-2013 at 06:48


The cell potential for Cu2+/Cu (cathode, reduction)///SO2/SO4(2-) (anode, oxidation) is +0.14 V, so 'theoretically' the reduction of cuprous ions with sulphur dioxide in aqueous solution is possible. But 0.14 V is a small value... Compare it to the reduction of silver where the cell potential is +0.60 V.

[Edited on 15-2-2013 by blogfast25]




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[*] posted on 15-2-2013 at 11:54


Quote: Originally posted by Mixell  
You can use dimethylglyoxime to separate the nickel from the copper (Ni(dmgH)2 precipitates), although its not very affordable, may be only on small/demonstration scale.

UPDATE- Just read on the German wiki that it forms a complex with copper too, but my question below still stands.

I'm planning to form this complex myself, as I understand, it is used as a solution in ethanol, can acetone replace it?
Also, do I need any special conditions to dissociate the dimethylglyoxime into dmgH- and H+ like a basic environment? Or a dilute solution at neutral pH will also give good yield?

[Edited on 31-7-2011 by Mixell]


I'm really late to this party, but this should work, because the nickel complex is insoluble, while the copper complex is soluble in water. You just have to make sure you add enough DMG to react with both.

Vogel's Quantitative Inorganic Analysis gives the procedure- the H2DMG is dissolved in ethanol (acetone would probably work). Dissolve your nickel-containing sample in distilled water with a small amount of hydrochloric acid (for 0.3 g of (NH4)2Ni(SO4)2 hexahydrate, use 200 mL water and 5 mL 6 M HCl. Heat to 70 - 80 oC, add a slight excess of the DMG solution (5 mL for every 10 mg of Ni present), and then add ammonia until precipitation occurs, and then "in slight excess". Allow to stir at the same temperature for half an hour (this decomposes any excess H2DMG), then filter out the precipitate. If iron or chromium is present, then add 5 g of citric or tartaric acid to prevent the precipitation of the Fe(DMG) complex or Cr(DMG) complex.

You want to add a slight excess of ammonia because you want the solution basic enough that the DMG is present as the anion, but not so much that you get the ammonia coordinating to the nickel.

ETA: The above procedure is not recommended for samples containing more than 30-40 mg of nickel, due to the bulk of the precipitate. For larger samples, Vogel recommends making the solution (for practice, 0.33 g of nickel ammonium sulphate in 100 mL water) acidic (1.85) with HCl, then adding 10 g urea and 30 mL DMG solution. Upon heating with a steam bath, the urea breaks down to give ammonia, which neutralizes the acid. Since this slowly raises the pH (precipitation starts within 10 minutes, and is complete after 90 minutes), you get a compact crystalline precipitate instead of a flocculent one.

[Edited on 15-2-2013 by DraconicAcid]




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[*] posted on 16-2-2013 at 05:35


DA:

The DMG method is a really expensive way. Most of us don't have access to it. Great for wet analysis, not so great to separate significant quantities...




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[*] posted on 28-9-2013 at 23:39


just a quick update
Quote: Originally posted by tetrahedron  
pure nickel anode is preferable in order to remove most of the remaining Cu2+ at the very end

this worked. using flat Ni electrodes allowed for the removal of the remaining Cu2+ from the solution of the mixed sulphates. the current ranged between 1.5-1.8 V at 0.7-1.9 A/dm2 max, and was stopped when a dulling of the pink Cu-plated cathode was observed. the anode was left in solution overnight, but no new Cu deposited on it.

i now have a process for the electrolytic separation of cupronickel:

1. electrolysis of CuSO4 with Cu cathode (e.g. electrical wire) and a 'flat' Cu/Ni anode made by soldering a bunch of nickels to individual wires, and insulating the solder and any exposed wire with spray paint or similar (just spray the whole backside and let dry thoroughly); the Cu plates out preferentially, leaving behind a high Ni2+ concentration (possibly the anode needs to be changed a few times before the Cu2+ is anywhere near depleted)

2. electrolysis of the previous product with a flat Ni anode and a cathode of the same size; the remaining Cu2+ and solder ions will plate out completely (hopefully, since the electrode potentials of Pb/Sn are very close to that of Ni), leaving behind pure nickel sulfate solution, which can be stored, crystallized, or used for Ni plating.
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[*] posted on 29-9-2013 at 22:55


A few years back I was trying to do what the OP was, and I tried several different methods of separation. Fortunately I have access to an EDS system, so I could determine exactly what I had.

One method was in exploiting the different solubilities of nickel and copper hydroxide. Neither is very soluble, but I wanted to try it. The problem with this method was that in the presence of other hydroxides, copper hydroxide becomes much more soluble than it is by itself. After adding DI water, stirring under heat, and cooling to precipitate the excess copper hydroxide, I was only getting something like 90% nickel.

I tried using hydrogen sulfide to remove the copper. That worked, but it was a mess, and stank up the lab...even in an efficient fume hood.

Electrolysis was slow, and could not achieve good separation. The reduction potentials of nickel and copper are close enough together, that once the nickel becomes concentrated in solution, it plates out with the copper. At this point there is still a lot of copper in solution. You can minimize this by using a very low current density, but this takes a VERY long time.

Using stochiometric amounts of potassium iodide, I was able to achieve very good separation. Cupric iodide is unstable, and cuprous iodide is very insoluble. The drawback is the expense of iodide salts.

The way the big boys do it, is via nickel tetracarbonyl, aka "creeping death". I decided not to try that.

After spending a frustrating amount of time to isolate a few bits of impure nickel salts, I discovered a nearby pottery store that specialized in chemical sales. It was like walking into a candy store. In one long isle, there were piles of little paper baggies of various chemicals. I stocked up on carbonates/oxides of nickel, copper, lithium, barium, strontium, manganese, and several other things. The stuff was all relatively cheap, which was a big surprise. Every time I go back there I wear rubber gloves, and I seal my purchase in a large ziploc bag. I then wipe off the bag with a wet towel, and cautiously dispose of the gloves in the trash. It freaks out the poor girl at the sales desk.

I bought a 4" square piece of electrolytic nickel on eBay, and several years later I'm still using pieces of it for electroplating experiments. It was well worth the $50.

The experience that I gained trying to isolate the two elements served me well, but in the end I had to just go buy it.

[Edited on 30-9-2013 by WGTR]
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[*] posted on 30-9-2013 at 14:48


you were getting 90% nickel out of copper/nickel hydroxides simply by adding DI water and heating then rapid cooling?the copper hydroxide precipitates and the nickel hydroxide wont?90% is not bad at all but why cant it be further heated and precipitated?
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[*] posted on 30-9-2013 at 17:29


Something like that. The heating and cooling was slight, and was only to speed up the equilibrium. It's not completely necessary. Very efficient stirring is necessary, however.

Solubilities according to Wikipedia:

1. Nickel (II) hydroxide: 1.4 x 10^(-3) mol/L

2. Copper (II) hydroxide: Ksp = 2.20 x 10^(-20). From this:
2.20 x 10^(-20) = [Cu2+] [OH-]^2
2.20 x 10^(-20) = 4x^3
x = 1.77 x 10^(-7) = mol/L Cu(OH)2

On the surface it looks like this would be an effective way of separating nickel from copper, but it's more difficult than that, especially given the low solubility of both compounds.

Since you asked, I'm going to run back to the lab to see if I can find my old samples, along with the EDS spectrums. I think I still have them.

[Edited on 1-10-2013 by WGTR]
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[*] posted on 30-9-2013 at 19:59


I discovered that I ran the experiment back in 2009, so it's somewhat miraculous that:

A. I still had the samples, and
B. I was able to find them in all of my junk.

I need to correct some of what I said before. It turns out that I wasn't remembering everything very clearly.

I've attached two files, one with the samples and spectrum for the iodide run, and the other one for the hydroxide runs.

The precursor salts were prepared by reacting the metal carbonates of a US nickel (25% Ni/ 75% Cu) with various acids.

In the iodide experiment, the mixed nitrate salts were reacted with potassium iodide. There is a handwritten note on the spectrum that shows how this works. The extra iodine is shown as a gas, because I removed it with gentle heat. Surprisingly, the separation wasn't as good as I remembered. Perhaps I wasn't using the right amount of KI. I included an image of both the copper(I) iodide and the nickel carbonate. It should be obvious which is which.

In the attachment for the hydroxide experiment, I included three spectrums along with their associated samples (on filter paper). First of all, I apologize, but I scanned them in the reverse order. So the last spectrum is actually the first, and so on.

In this experiment I precipitated the hydroxides from the mixed sulfate salts. I mentioned heating and cooling the solution earlier, but had forgotten why I did it. When heating a stirred solution of copper/nickel hydroxides, the hydroxides become more soluble (up to the point that the hydroxides irreversibly decompose to oxides).

So, I'll explain the spectrums, now. The first run (on the third page) is a spectrum from the leftover solids in a stirred warm solution of the mixed hydroxides. The spectrum is fairly meaningless, because I didn't record how much hydroxide I put into this first solution. That's why I didn't bother quantifying the results. It does show, however, that I added enough water to bring all of the nickel into solution, leaving fairly pure copper hydroxide behind.

Next, I took the clear warm supernatant, and let it cool to room temperature. The resulting precipitate was redissolved in a smaller volume of stirred, warm DI water; and the excess solids were analyzed. This is the second spectrum (second page). I started to see nickel turn up in this sample (6%), so I decided that I was approaching equilibrium in the solution.

I then took the warm supernatant, and let it cool to room temperature. The resulting precipitate was analyzed (spectrum on first page), and it was noted that the nickel concentration was 97%.

Keep in mind that I did this over four years ago, and my lab practices weren't as good as they are now. This is why the results seem a little haphazard. Also, note that the solubilities of the hydroxides are very low. The small bits on the filter paper should be an indication of how much material I was working with. I also didn't bother reheating, and precipitating again, to see if I could get even better purity.

Attachment: Iodide_precipitates.PDF (122kB)
This file has been downloaded 198 times

Attachment: Hydroxide_precipitates.PDF (336kB)
This file has been downloaded 290 times



[Edited on 1-10-2013 by WGTR]
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[*] posted on 3-10-2013 at 18:34


so really converting the copper/nickel sulfate solution to the hydroxides then dissolving in hot water and stirring,then cooling.repeat process with precipitate and more nickel than copper will eventually drop. take it once the nickel and copper are in their oxides it will be almost impossible to separate.
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[*] posted on 4-10-2013 at 03:05


Run a stream of acytelene through the solution of copper/nickel nitrate then filter the solution.
Be warned this makes copper acetylide, a sensitive explosive.

[Edited on 4-10-2013 by bismuthate]

[Edited on 4-10-2013 by bismuthate]

[Edited on 5-10-2013 by bismuthate]




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[*] posted on 4-10-2013 at 16:44


I'd like to try separating the mixed hydroxides again before I say much more about my previous results. This topic comes up now and then, and it seems important to some people, so I want to be sure that the experiment can be duplicated.

I may not be able to get to it right away, though. It's been a long week in the lab, and I'm pretty tired.
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[*] posted on 4-10-2013 at 19:21


no problemo WGTR. i cant believe its been two years already since this thread was started.
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