Sciencemadness Discussion Board - how extrac nickel

how extrac nickel

plante1999 - 6-2-2011 at 14:37

how i can extract nickel from copper nickel alloy?

thanks!!

Sedit - 6-2-2011 at 14:38

Tweezers, really little tweezers.

plante1999 - 6-2-2011 at 14:47

....i see. i want to extrat it chemicaly to make nickel sulfate.

bbartlog - 6-2-2011 at 17:02

Not all that easy. There are already threads here on separating nickel and copper in solution; so you could dissolve the cupronickel in HCl (with oxidizer, i.e. bubbled air or H2O2) and use one of the methods discussed there (many of these revolve around reducing the copper to precipitate it as a cuprous salt).
There is also the Mond process which involves using carbon monoxide to selectively remove the nickel as nickel tetracarbonyl, but I imagine it would be hideously dangerous to do at home, and anyway I'm not sure it would work so well on a substrate that was mostly copper.

Sedit - 6-2-2011 at 17:11

OK ok, you can get some of the copper out atlest appearent from my experiment with it where I left the sulfate of US nickle coins out for sometime, (months) after along time I was left with a green solution and on the bottom was large(1inch and up) copper sulfate crystals all along the bottom.

I know that the mother liquid was now more concentrated in Nickle sulfate now but am unsure how much more concentrated.

I also tried (aq)NH3 which may have merit but I never fully investigated it. Copper forms a complex and Nickle precipitates as the hydroxide/oxide? However Nickle can also form a complex but since copper does not precipitate in this manner after a few washes I was able to clean the precipitate and convert it to the chloride using HCl and it proved to be Nickle chloride. I call the experiment incomplete because I never obtained data on yeilds or correct concentration of ammonia solution in order to get the best purity and yeilds.

blogfast25 - 7-2-2011 at 09:15

There are countless threads on this.

1. Dissolve coins in dilute NITRIC acid (HCl is useless for Cu AND Ni). Takes minutes, depending on acid strength.

2. Separate Ni from Cu on the basis of the difference in solubility of CuS and NiS (solubility products respectively 6 E-39 and 3 E-19). You need to adjust the pH of the solution to the right value, then saturate the solution with H2S. Copper drops out as black CuS, Ni2+ stays in solution. Filter, acidify filtrate and boil to get rid of sulphide. DO NOT KILL yourself with H2S poisoning.

Which pH exactly you need can be calculated and the calculation is available on this forum. Search and yee shall find! Look also for a Ni/Cd separation thread: same principle...

Edit: a quick calculation showed that at pH < 2 the copper will still precipitate as CuS but the nickel stays in solution as Ni2+. At that pH the concentration of S2- is simply too low for the solubility constant of NiS to be reached...

[Edited on 7-2-2011 by blogfast25]

Wizzard - 7-2-2011 at 10:01

Easiest way is the Mond process. BBart is right, but it's not that dangerous.

1. Granulate the alloy as small as you can.
2. Seal the alloy and carbon monoxide (generated by Zn+CaCO3->ZnO+CaO+CO in a sealed glass vessel with heat) in a long glass vessel (for a temperature gradient). Consider a lower-case H, with your alloy in one leg, the CO source in the other and the heated end up top. Seal off the CO source leg after the vessel is pressurized, and you will want to get at least some kind of vacuum in the chamber to suck out the atmosphere, but I dont think it will tamper with your reaction much (especially if you're just making nickel sulfate, and not seeking pure nickel).
3. Heat one side to at least 75*C, and the other side (decomposition) to 250*C.

The nickel and only the nickel will be deposited on the hot side- Just be careful when you open the vessel, do it outside! CO is toxic!

blogfast25 - 7-2-2011 at 10:15

Quote: Originally posted by Wizzard

Easiest way is the Mond process. BBart is right, but it's not that dangerous.

1. Granulate the alloy as small as you can.
2. Seal the alloy and carbon monoxide (generated by Zn+CaCO3->ZnO+CaO+CO in a sealed glass vessel with heat) in a long glass vessel (for a temperature gradient). Consider a lower-case H, with your alloy in one leg, the CO source in the other and the heated end up top. Seal off the CO source leg after the vessel is pressurized, and you will want to get at least some kind of vacuum in the chamber to suck out the atmosphere, but I dont think it will tamper with your reaction much (especially if you're just making nickel sulfate, and not seeking pure nickel).
3. Heat one side to at least 75*C, and the other side (decomposition) to 250*C.

The nickel and only the nickel will be deposited on the hot side- Just be careful when you open the vessel, do it outside! CO is toxic!

Got a convenient source of CO, do you Wizzard? ;-)

ScienceSquirrel - 7-2-2011 at 10:21

Nickel carbonyl is far more toxic and nastier than carbon monoxide, making it at home with improvised equipment is just plain stupidity.

blogfast25 - 7-2-2011 at 12:13

If you’re looking for a cheap source of Ni, use scrap Nichrome wire. Cr and Ni are relatively easy to separate, on account of Cr’s amphoterism. Dissolve in nitric, precipitate hydroxides with NaOH (or Na2CO3) and then leach with copious amounts of strong, hot NaOH. Chromite gets lixiviated out, Ni(OH)2 stays insoluble. There may be easier ways to do this, but this one doesn’t take much…

Wizzard - 7-2-2011 at 13:47

Just snagged from Wikipedia, calcium carbonate and elemental Zinc

And a small amount out in the open can't be all THAT toxic- The directions provided (sealed in glass under vacuum), with a small vessel, all it would take is some professional caution.

I think sealing it under glass and then pulling a vacuum (and then heating said vessel) would require more caution- I dont think anybody recommends doing these things in their living room, so I would assume the whole project would be lab/garage do-able.

Also, nickel tertracarbonyl decomposes in air, in 60±5 seconds

Ref: http://adsabs.harvard.edu/abs/1980Sci...208.1029S

ScienceSquirrel - 7-2-2011 at 14:03

It is very toxic, Wikipedia has a well referenced article;

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

bbartlog - 7-2-2011 at 14:26

Quote:

HCl is useless for Cu AND Ni

Eh, it depends on the timescales you are operating on. You're right that if you want to dissolve the metal within minutes or a few hours, HCl is not so good (maybe with 30% H2O2 it would work, haven't tried it). But over a period of days, bubbled air plus HCl can work to dissolve either copper or cupronickel. I agree that nitric acid is preferable if you have it; it just happens that HCl is cheap and OTC where I am, whereas I haven't found an OTC source for nitric acid. Which is why I've used it.

I like the sulfide precipitation method you propose; I actually don't remember seeing that one before.

Quote:

Easiest way is the Mond process

uh. Granulate cupronickel? With what, a mortar and pestle? How is that easy? But OK, let's say you use a bench grinder and get yourself some cupronickel dust (sounds like a PITA on any larger scale but w/e). Next step is to somehow create a sealed glass vessel in the shape of an 'H', preferably with vacuum attachment, carbon monoxide source and stopcock. Then we get to control temperature in two places carefully. I should note that nickel tetracarbonyl can decompose explosively, so if you screw up somehow you could end up shattering your glass apparatus and suddenly exposing yourself to an exceedingly toxic gas.

Quote:

be careful when you open the vessel, do it outside! CO is toxic!

It's not the CO you have to worry about, here. Short-term exposure limit for CO is 400ppm (for a 15 minute exposure, according to the EPA - so this is still sort of a safe level). By comparison, exposure to 3ppm of nickel tetracarbonyl for half an hour hits the LD50, and a single good lungful at 400ppm would likely kill you. Conservatively I would say it's a thousand times more toxic than carbon monoxide. And as an added bonus, sublethal exposures can still leave you with heavy metal poisoning, whereas most sublethal CO poisoning will not result in serious long-term effects.
I'm not particularly cautious with chemicals, but really - leave this one alone unless you have professional-level skills and equipment. I might add that many professionals also leave it be, see for example: http://pipeline.corante.com/archives/2004/03/28/thing_i_wont...

not_important - 7-2-2011 at 14:46

Simple granulating of the alloy isn't very effective, the metals need to be very fine powders else the copper will shield much of the nickle (remember that most nickle-copper coins contain more copper than nickle).

Nickle carbonyl is extremely toxic, and not to be messed with lightly. Deposits nickle within the body, it's toxic by itself, and even seconds is enough exposure to really hurt you. And note that the decomposition time you referenced is in a carbon monoxide free state, and the Mond process needs an excess of CO for practical rates; if it leaks it'll hang around for longer than that minute.

Retard-3000 - 7-2-2011 at 14:58

Could he not react the alloy with nitric acid to form nickel nitrate and copper nitrate then use a more reactive metal then Ni and Cu such as Aluminium or Magnesium to displace the nickel and copper?

blogfast25 - 8-2-2011 at 09:57

Bbart:

Yes, if you have days to spare, HCl is great for Cu (LOL)! Ni and Cu are very similar there BTW, see a thread I got on ‘best solvent for silver electroplated plated nickel’. Nitric is the solvent of choice for both, whether you have it or not.

The sulphides separation methods were once standard textbook material for wet analytical chemistry. Used for separating ‘more soluble’ sulphides (like NiS and ZnS) from 'more insoluble ones' (like CuS, HgS, CdS, Sb2S3 etc). By regulating the pH, you regulate the concentration of S2-, thereby controlling which solubility products get exceeded and which not. Smelly but simple and VERY effective (the degree of separation is extremely high).

Still: nichrome wire, anyone?

[Edited on 8-2-2011 by blogfast25]

kmno4 - 8-2-2011 at 12:45

Quote: Originally posted by plante1999

....i see. i want to extrat it chemicaly to make nickel sulfate.

It can be done by electrochemical method. Cu/Ni alloy as anode, electrolyte - H2SO4aq + some dissolved alloy, Cu or SS as cathode.
You will get NiSO4 in solution and Cu on cathode.
Very elegant, very simple and economic way.
(I have impression that I have already seen discussions about separation of Cu-Ni in this forum)

blogfast25 - 8-2-2011 at 12:50

Quote: Originally posted by kmno4

Quote: Originally posted by plante1999

....i see. i want to extrat it chemicaly to make nickel sulfate.

Voltage across electrodes, approximately?

kmno4 - 8-2-2011 at 14:01

It depends, but less than 5 V if everything is OK.
It is more important not to exceed some current density.
At room temp. it is about 1 A/dm2.
In another case Cu discharging at cathode is faster than dissolution of alloy and Cu(2+) ions dissapears from solution.
At this moment voltage rises (at constant current) and cathodic Cu becomes very spongy and tends to grow in direction of anode.
This "critical" current density can be moved to higher values when solution is warm (50-80 C) + stirring.
Stirring is very important factor and can move this c.c.d to 5-10 A/dm2 or higher.
Dissolution of monel-like alloys is not very fast process but electrolysis can be conducting at low current densities without special control even for a week.

blogfast25 - 9-2-2011 at 08:20

Fairly large cathode should be an advantage to keep current density low...

cyanureeves - 13-2-2011 at 13:07

i am making nickel sulfate right now using two cupronickel coins in water/sulfuric acid solution and everything looks pretty good so far. i intend to filter off nickel sulfate solution and ultimately make nickel acetate using vinegar.i dont know if sodium hydroxide or sodium carbonate would be better to add to the nickel sulfate before making the acetate.i would really like to make nickel acetate from nickel oxide but i can only make the black nickel oxide when i burn it with a torch. i have nickel sulfate but it is getting expensive to buy and nickel oxide(the yellow green one) is twice as expensive.is the green nickel oxide pretty hard to make?is the black nickel oxide i make by burning nickel carbonate in fact nickel oxide?nickel acetate is mainly what i want.

blogfast25 - 13-2-2011 at 14:15

What electrolysis set up are you using (cathode, anode, voltage applied, current obtained, distance between electrodes...)?

Use NaOH and you get Ni(OH)2, use Na2CO3 and obtain Ni2CO3. Both dissolve well in vinegar to NiAc2.

Sedit - 13-2-2011 at 15:23

I have some good news, atlest it seems that way to me. I decided to mix up some HCl and H2O2 and place some Stainless steel flatware in it since this is a nickle source that does not contain copper. It has been sitting for a while and of course its not dark green.

I added some to a flask and neutralized it with an ammonia solution. To my delight I got the precipitate I got from the experiments with US coins yet this time there was no blue complex which is suppose to form with Nickle. That seems to suggest that seperation of the two is as simple as salting the ore into its sulfates or HCl salts then mixing it with a solution of Ammonia hydroxide. The copper will stay in solution and you can recover your nickle as a mixture of oxides/hydroxides.

There is one problem and thats the fineness of the precipitate. Gravity filtering is just a pain in the ass and will take forever and it does not settle fast at all so decanting is also a messy operation that takes to long.

Any suggestions on how to clump up the precipitate?

nickel

cyanureeves - 13-2-2011 at 15:43

Quote: Originally posted by blogfast25

What electrolysis set up are you using (cathode, anode, voltage applied, current obtained, distance between electrodes...)?

Use NaOH and you get Ni(OH)2, use Na2CO3 and obtain Ni2CO3. Both dissolve well in vinegar to NiAc2.

i am using 4.8 volts, distance between electrodes is about 4". the copper is all over the cathode and i just keep scraping it off. the solution is the familiar green so far at about 4hrs of electrolyzing. i've done the carbonate before but did not plate with it because i was not sure if it was really acetate. i will be using the hydroxide this time then.both anode and cathodes are cupronickel mexican coins and i'm getting a bonus metal precipitate right below the annode. i believe i read mexican coins also have zinc. thank you.

blogfast25 - 14-2-2011 at 08:27

Quote: Originally posted by Sedit

I have some good news, atlest it seems that way to me. I decided to mix up some HCl and H2O2 and place some Stainless steel flatware in it since this is a nickle source that does not contain copper. It has been sitting for a while and of course its not dark green.

I added some to a flask and neutralized it with an ammonia solution. To my delight I got the precipitate I got from the experiments with US coins yet this time there was no blue complex which is suppose to form with Nickle. That seems to suggest that seperation of the two is as simple as salting the ore into its sulfates or HCl salts then mixing it with a solution of Ammonia hydroxide. The copper will stay in solution and you can recover your nickle as a mixture of oxides/hydroxides.

There is one problem and thats the fineness of the precipitate. Gravity filtering is just a pain in the ass and will take forever and it does not settle fast at all so decanting is also a messy operation that takes to long.

Any suggestions on how to clump up the precipitate?

“Stainless steel flatware”? What do you mean?

Strange indeed that no Ni(NH3)6 (2+) formed… You sure about this?

Freshly precipitated hydroxides can often be compacted by means of gentle simmering because they lose water: M(OH)m.n H2O (s) === > M(OH)m.n-x H2O (s) + x H2O (l)

Sedit - 14-2-2011 at 11:36

Quote: Originally posted by blogfast25

Stainless steel flatware”? What do you mean?

Strange indeed that no Ni(NH3)6 (2+) formed… You sure about this?

Freshly precipitated hydroxides can often be compacted by means of gentle simmering because they lose water: M(OH)m.n H2O (s) === > M(OH)m.n-x H2O (s) + x H2O (l)

Wikis stainless steel artical shows flatware, as in forks and spoons ect.., as "typical composition of 18% chromium and 10% nickel, commonly known as 18/10 stainless".

Im not sure about it blogfast but what I am sure about is that there was absolutely no color to the solution at all other then a grimmy oxidation product of the Fe dropping out of solution making it a dull orange until it settled. Nickle complexes like copper was described in the other thread on the topic to turn a blue color just like copper. There was no blue color when excluding the copper from the equation. This time it was just green precipitate and clear supernatant fluid.

This is inline with other experiments I performed when seperating the Copper before. After a complete washing of the formed precipitate I proceeded to add HCl and formed what by all means looked like Nickle chloride crystals. I put them back in ammonia solution expecting to see some kind of a blue color yet I saw nothing that backed up the old text describing the complex.

The only explinations I can think of, and bear with me because im just musing here, is that the ammonia salt does not allow the complexes formation somehow... The precipitated hydroxide does form a complex but its not soluble so its of little concern, or... It just does not exist at all.

woelen - 14-2-2011 at 12:23

[Ni(NH3)6](2+) is a well-known complex of nickel. It is formed when ammonia is added in large excess amount to a solution of a nickel salt. It has a nice blue/purple color, somewhat like the [Cu(NH3)4](2+) complex, but the color is less intense and it is a little bit more purplish.

Sedit - 14-2-2011 at 12:44

I have seen that slimy green a number of times, thats the precipitate I always see yet never caught the complex at all. Its interesting to say the lest. I see you have mentioned a large exess is needed to form the complex and the fact im using a dilute solution may have alot to do with it.

All in all however that would mean that since copper forms quickly and easily with NH3 and Nickle needs excess it should be no problem to extract the nickle from the copper using ammonia hydroxide because the Cu will want the amine much more then the nickle and when using the dilute solution the precipitate should stay in solution where as the nickle will precipitate.

This is inline with what I have been seeing and when I isolated the nickle and converted it to its chloride I got deep green needles formed from it on drying. It looked almost exactly like this web image I managed to pull up

[edit...NM I can not get the photo to link for some reason....]

The only issue now is getting the precipitate to clump to the point where it can be seperated because it is very fine and generally hard to filter without serious looses. Im going to try placing it on a warm plate later and see if the convection will clump it up for me.

[Edited on 14-2-2011 by Sedit]

smuv - 14-2-2011 at 12:44

Much flatware today is made with nickel free alloys. In order to determine if it is 18/8 (304) you should test to see if it is magnetic. 304 stainless is generally non-magnetic but can be faintly magnetic in areas where it has been worked (read: sharp bends). On the other hand 4xx series stainless have no nickel and are by far more popular for cheep flatware.

Sedit - 14-2-2011 at 12:47

In that case I have what Im pretty sure is some scrap Nichrome wire that I can always run a test with that but I don't feel the results are going to be much different. Ill see if they have the piece labled for simplicity.

Sedit - 14-2-2011 at 15:51

Sorry to double post but I thought I would bump this real quick since I figured out the last piece of the extraction puzzle.

The basic process for say US Nickle coins would be to dissolve them in a suitable acid Nitric being the fasted but I have found H2O2+HCl to work albeit in a day or so.

Next you would dry your resulting solution to get crystals or just reduce your solution enough to carry on to the next step which would be basification with ammonia hydroxide to precipitate the Nickle hydroxides and trapping the Cu ions in solution as an amine complex. Once precipitation stops on the addition of more Ammonia then one should stop.

Allow the precipitate to settle and wash it a couple times with distilled H2O. Each time allowing it to settle a bit to make working with it a bit simple.

The final piece of the puzzle is something I have used before but it just never struck me to use it here and that is to just pour the sample into an unglazed ceramic dish. Terracotta pots would work here assuming the bottoms have been sealed. This will dry your preciptate in relatively short order compaired to filtering.

Mix your precipitate with HCl or H2SO4 and dry to aquire the desired Nickle salt.

not_important - 14-2-2011 at 16:04

You do know that nickle also forms ammonia complexes similar to those of copper, don't you? With just enough ammonia a deep blue complex is formed, a large excess of NH3 shifts this to a lavender hued complex.

Sedit - 14-2-2011 at 16:42

Yup, thats what we have been discussing yet it takes an excess of NH3 to form the complex and from experiments its showing that the Copper is much more hungery for that ammonia then the nickle is. I recall that in the other thread you are the one that brought this complex to my attention some time back and after some experimenting with it I can honestly say that as long as the NH3 is not over done one can simply extract Nickle from a mixture of Copper and Nickle using Ammonia hydroxide.

Its a simple enough experiment to try out and I would suggest to anyone to give it a shot. If nothing else to prove me wrong or point out something im missing. One thing that still has thrown me off however is that after getting the precipitate and cleaning it addition of alot of Ammonia hydroxide does not produce the blue color and I don't understand that one bit.

plante1999 - 14-2-2011 at 17:14

wath pourcentage of H2O2 chould i use for 30%HCl?

Sedit - 14-2-2011 at 18:01

I use 35% H2O2 for Muratic acid. Performing an electrolysis with the Nickle bearing material as the electrode in HCl would also more then likely speed up the process a bit without having to use H2O2 to oxidise the HCl.

I also want to add that I tested my flatware to see if there is Nickle present by taking the solution and reducing it with Aluminum foil. It was very slow at first to displace anything only producing a small amount of silvery metal on the surface of the Aluminum so I added a bit of HCl to speed things along which worked well. A magnet showed indeed that there is Nickle present yet with Ammonia hydroxide there was absolutely NO blue color or any color for that matter. This shows that the complex does not form without some effort in a dilute solution.

The only other metals present in this test are Iron and Chromium is Wiki is to be trusted. I realize I will have to do a photo assay of the process to show everyone and I will indeed do that soon enough in the Cu++/Ni++ seperation thread.

I have to ask Plante, are you non american or lazy? Could you please try to spell a little better or atlest show some effort.

plante1999 - 14-2-2011 at 18:18

Quote: Originally posted by Sedit

. I realize I will have to do a photo assay of the process to show everyone and I will indeed do that soon enough in the Cu++/Ni++ seperation thread.

I have to ask Plante, are you non american or lazy? Could you please try to spell a little better or atlest show some effort.

1-if you do thant it will help me alot

2-sorry for my not very well spelled english but i make very high effort to writh wath i think in english,im a French Canadian and my commune language is french, because i think french forum is done wrong and are not familiar with chemistry (sorry for those like French forum) i prefer to go to much more specialised english forum like sciencemadness , i will try to make less mistake.

thanks!

cyanureeves - 15-2-2011 at 16:58

sedit. i dried my nickel carbonate i made from the sulfate and some copper was in the drying slurry because i could see bright blue swirls on the surface of the light green mass although not much.i added ammonia hydroxide to some of the dried carbonate and sure enough i got a blue that rivals prussian blue in color.i now have a precipitate of a light green color on the bottom of the blue liquid and no matter how much water i keep adding it still precipitates as a light green carbonate color.would this be my nickel hydroxide or just cleaner nickel carbonate?i dont plan on going any further with this nickel other than into a pot for an attempt to acquire an oxide to convert to acetate,and onto my home made lab stand as plating.all this time i thought only silver electrolyzed in ammonia and absent of oxygen would make the bright blue color.

Sedit - 15-2-2011 at 17:33

I think thats your hydroxides or oxides I can't remember which and the blue was your Copper complexing into solution. Test for yourself by drying the precipitate and washing it. After that try to add it to Ammonia solution and see if you can still get that blue color. I personally have not obtained it yet although i'm going to keep trying to get Nickle blue just to understand a bit better but I have on a few occasions added HCl to the precipitate and dryed it to yeild what is with out a doubt fairly pure NiCl<sub>2</sub>

cyanureeves - 15-2-2011 at 19:03

sedit. thanks.it would be cool if indeed it was pure nickel on the bottom because i didnt even use much ammonium hydroxide or maybe i just made a hell of a strong hydroxide with my ammonium nitrate/sodium hydroxide.i dont understand much about your complex and i'm eager to just bake any nickel salt in search of the oxide because it can plate out of a cold solution.sorry for seeming selfish but perhaps i'll only set you back in your search if i do this test, only to find that i still have some impurity with an affinity for ammonia and turns blue.

Random - 16-2-2011 at 10:23

I think I posted one thread earlier and I remember that one metal of these two can be precipitated with salicylic acid.

kmno4 - 16-2-2011 at 15:38

Separation Cu/Ni from mixure of their sulfates is extremely easy.
Simple electrolysis with C anode will remove practically all Cu.
Fe contamination is easy to remove as "Fe(OH)3".
Precipitating of Ni carbonate (to convert it to another salt) should be done after Cu-Fe remove. Conversion of contaminated Ni sulfate to contaminated Ni carbonate is rather losing time...
BTW: ammonia-nickel complexes are blue.

Sedit - 17-2-2011 at 09:37

I think better then US currency in order to get Nickle the flatware or Nichrome wire would be a much better solution.

I decided to take my entire dark green "Flatware solution" and toss in Aluminum and after a hell storm of bubbles because of the excess acid, I was greeted with a seemingly large amount of grey/black magnetic powder precipitate which im, assuming to be elemental Nickle, and a bluish solution. Sure there is alot of contamination in it since I never precipitated the Fe as the hydroxide but this could work as a means of Ni for the determined. Due to the fact that the solution is still blue I believe that the Chrome did not precipitate out like the Nickle did.

Random - 17-2-2011 at 13:06

Quote: Originally posted by kmno4

Separation Cu/Ni from mixure of their sulfates is extremely easy.
Simple electrolysis with C anode will remove practically all Cu.
Fe contamination is easy to remove as "Fe(OH)3".
Precipitating of Ni carbonate (to convert it to another salt) should be done after Cu-Fe remove. Conversion of contaminated Ni sulfate to contaminated Ni carbonate is rather losing time...
BTW: ammonia-nickel complexes are blue.

Can this maybe be done with electrolysis of copper/nickel acetate mixture? I should try this. What will happen with nickel though?

Do you mean copper cathode on which will copper be deposited and C anode?

cyanureeves - 18-2-2011 at 07:30

sedit i decided to go ahead with the ammonia thing and i got the royal blue color solution on adding ammonium hydroxide to the nickel carbonate.if i use just enough to get the blue solution and stop there will be nice green crystals on the bottom.when added more ammonium hydroxide to the remaining dried green crystals i got a carribean blue.a chart i read stated that royal blue is the nickel complex and the sky blue was copper.i dont know why nickel would beat out copper first if this is my case,but remember i used a high copper content coin.im sure i will get the lilac color if i add more ammonia to the royal blue solution.anyway i made nickel carbonate from nickel sulfate and roasted it on a s.steal pot and i finally got what looks like the oxide(yellow-lime green) this time i used a propane burner instead of oxy-acetylene.i was going beyond melting and wiki said the salt decomposes before melting.i'm pure nickel rich right now but one day i'll come back to this thread to dig for nickel metal i'm sure.

Random - 18-2-2011 at 09:46

I think I figured out the best way to do it. Cuprous chloride is insoluble in water, we must then reduce copper (ii) chloride. But HCl with air oxidizes it again into copper (ii) chloride. That means we need to have solution that won't contain any more HCl or very little and then filter cuprous chloride fast!

1. Dissolve cupronickel coins in HCl (maybe add very little of some nitrate solution to speed the reaction a little bit).
2. Add baking soda to destroy excess HCl until the carbonates will start precipitating.
3. Add just enough HCl to dissolve the resulting carbonates.
4. Make a solution of potassium metabisulfite and mix it with cupronickel chlorides.
5. Resulting SO2 and metabisulfites reduce copper (II) ions to copper(I) ions. Cuprous chloride precipitates.
6. Filter that very fast and you have solution of nickel, sodium, so4, so3 and chloride ions.
7. Precipitate nickel carbonate using baking soda .

bbartlog - 18-2-2011 at 10:10

One flaw in your proposed method is that sodium chloride solubilizes cuprous chloride. So while only 1.5g of CuCl will dissolve in 100ml of water (which may be an acceptable level of impurity, I don't know; but it's still not a very clean separation), a solution with 20g of NaCl dissolved per 100ml of water will dissolve a little over 7g of CuCl.
Some other chlorides also increase the solubility of CuCl (KCl and HCl among them).

Random - 18-2-2011 at 10:31

Quote: Originally posted by bbartlog

One flaw in your proposed method is that sodium chloride solubilizes cuprous chloride. So while only 1.5g of CuCl will dissolve in 100ml of water (which may be an acceptable level of impurity, I don't know; but it's still not a very clean separation), a solution with 20g of NaCl dissolved per 100ml of water will dissolve a little over 7g of CuCl.
Some other chlorides also increase the solubility of CuCl (KCl and HCl among them).

Actually, solubility of CuCl is 0.0062 g/100 mL (20 °C). It is almost completely insoluble. Well then, the method can be adjusted to avoid NaCl.

1. When you have cupronickel chlorides solution, add baking soda and precipitate copper and nickel carbonates.
2. Filter and dry them.
3. Add HCl drop by drop to dissolve the carbonates, be careful not to use excess HCl.

Now, add metabisulfite solution and proceed as above, nickel carbonate should really be enough pure now.

I think this is perfected, only maybe nickel carbonyl method can produce it more pure.

I also read that nickel sulphide is soluble, while copper sulphide isn't so you can separate these two with sodium sulphide or H2S. But metabisulfite way seems way easier and a lot less toxic.

[Edited on 18-2-2011 by Random]

cyanureeves - 18-2-2011 at 13:38

well random you should be able to plate out the copper because my carbonate now aceatae is contaminated with copper and i now have a pink pair of vice grips and dikes.i used a nickel anode and hopefully i will plate out all the copper because it bubbles great without even hooking up the juice. but i cheated on the vinegar by adding a bit of kodak bath stop acetic aacid.maybe if you use a s.steal anode or welding rod but no more copper for anode.i never could plate copper on carbon steal.darn copper! nickel plating is so begginer friendly too.

Random - 19-2-2011 at 11:36

Though, it's a lot easier if you separate them chemically, though ss anode should get oxidized too.

I actually again made mistake, maybe dissolving cupronickel carbonates in vinegar and then treating them with metabisulfite would work better.

Just, does someone know solubiility data for copper (I) acetate and can nickel (II) acetate be reduced too with SO2 to form insoluble compound? What is also the solubility data for copper (II) and copper (I) sulfites and nickel sulfites?