Sciencemadness Discussion Board

Separation of Copper/Nickel Alloy

kilowatt - 17-4-2008 at 17:38

I really don't know where to start here, to separate copper and nickel.

I have considered distilling out nickel (II) chloride while at the same time decomposing copper (II) chloride to insoluble copper (I), both of which happen at a temperature around 1000°C, but I'm not even sure if this would work at all as I do not know how the nickel chloride hydrates will break down. Even if it did work in principle, it would involve the release of water and chlorine at 1000°C, which will hopelessly attack any obtainable material, period. In short, there is no probably no way something like this would work.

I have considered fractional crystallization based on the alloy's phase diagram, but this would require tens of steps of melting at subsequently higher and carefully controlled temperatures to achieve decent separation, and this would be hopelessly energy intensive and time consuming. Might as well forget that, too.

I have considered powdering the alloy and using carbon monoxide to remove the nickel as nickel carbonyl in either a liquid or gaseous state, but this would probably only work with high nickel concentration, and leave quite a bit of nickel in the alloy. It would likely be hopelessly time consuming to remove most of the nickel, if possible at all. Again, probably hopeless.

I have considered reducing the mixed chloride solution with SO2 to precipitate insoluble copper(I) chloride according to the reaction:
2CuCl2 + SO2 + 2H2O --> 2CuCl + 2HCl + H2SO4
This looks like the most promising method, or rather the only one that has any promise, but I'm not sure if all the copper could be precipitated leaving pure nickel salt. Would H2SO4 and HCl not react with CuCl before it can be removed, putting some of the copper back into solution?

I'm out of ideas.

UnintentionalChaos - 17-4-2008 at 18:52

There is a reasonable gap in the electromotive series between nickel and copper. You could suspend lead metal (or tin but you'd have to keep it highly acidic to prevent hydrolysis) in the mixed salts (nitrate would be best by a long shot) until you see no further reduction of copper (a few more days after this point never hurt). Then treat with an alkali sulfate to remove the lead, use an alkali hydroxide to get Ni(OH)2 and dissolve in acid of choice followed by recrystallization.

I've isolated a small quantity of crude nickel carbonate from US quarters this way although it is a lot of work for the yield. Pottery supplier may be a much better alternative if you have access.

[Edited on 4-17-08 by UnintentionalChaos]

12AX7 - 17-4-2008 at 18:56

On your first idea, CuCl2 evaporates substantially over 400C or so; most chlorides are volatile. Besides being a surprisingly effective irritant, this would probably carry nickel with it and leave residues of CuO/NiO.

In alkaline solution, NiOOH can be produced with oxidation, but ah, copper doesn't go into solution while alkaline. Perhaps the copper tetrammine complex is stable under oxidation?

Cu(OH)2 and Ni(OH)2 have similar solubilities, no good for seperation. I don't know about Cu(OH,Cl)2 vs. Ni(OH,Cl)2, which has the second problem of: which oxychloride is produced?

CuCl is a possibility. The chloride solution could be cooked with excess metal, reducing copper to CuCl2- and keeping Ni in solution. An excess of water hydrolyzes the complex to CuCl which mostly precipitates. A lot of evaporation and a second pass, or fractional crystallization, may suffice to further purify the product.

If you had a lot of dimethylglyoxime, you could precipitate the nickel quantitatively, then -- somehow -- free the nickel from the complex and recycle the DMG.

Tim

[Edited on 4-17-2008 by 12AX7]

pantone159 - 17-4-2008 at 19:48

There is a thread on this board regarding the production of (insoluble) copper (II) 'aspirinate'. Ni(II) does not form a similar precipitate. Could that be used?

http://www.sciencemadness.org/talk/viewthread.php?tid=9920&a...

kilowatt - 17-4-2008 at 20:26

That copper (II) aspirinate looks really promising, fast and easy, and high yield as long as nickel does not likewise form a precipitate. I would have never imagined that aspirin could be so useful for metallurgy. It looks like it would be easy to re-precipitate the aspirin by acidification with HCl, giving copper chloride, and allowing the same amount of aspirin to be used almost infinitely many times.

CyrusGrey - 17-4-2008 at 21:02

Copper dissolves in ammonia slowly, or more quickly if you use an oxidizer like hydrogen peroxide.

I don't think nickel forms any ammonia complexes?

pantone159 - 17-4-2008 at 21:14

Quote:
Originally posted by CyrusGrey
I don't think nickel forms any ammonia complexes?


Ni(II) does indeed form ammonia complexes.

kilowatt - 17-4-2008 at 21:16

It would be pretty hard to leech all the copper out with ammonia I think, much as it would be pretty hard to leech all the nickel out with carbon monoxide. Maybe a combination of the two would work to eat away at the alloy at slightly elevated temperature though, going into a gaseous carbonyl phase and an aqueous ammonia complex phase. It would be really slow in any case I think. The aspirin method seems simple and fast and does not involve any incredibly toxic things like nickel carbonyl.

Mr. Wizard - 17-4-2008 at 22:36

In a previous post somebody had produced 'copper aspirinate'. Which makes a very good precipitate.
http://www.sciencemadness.org/talk/viewthread.php?tid=9920&a...
There was no precipitate with nickel sulfate. It would seem an easy task to dissolve the metals as a sulfate with an electrolytic cell, using the metal alloy as an anode (+) and sulfuric acid or some other soluble sulfate as the electrolyte. Epsom Salt comes to mind.

kilowatt - 17-4-2008 at 22:47

Quote:

It would seem an easy task to dissolve the metals as a sulfate with an electrolytic cell, using the metal alloy as an anode (+) and sulfuric acid or some other soluble sulfate as the electrolyte.

It would seem a very easy task, but in my experience it is somewhat less easy than one would think. I have found it works best if you reverse the polarity every few minutes between a pair of electrodes made of the metal, and even then it is extremely slow (we're talking days or weeks at around 4V and several amps to dissolve a reasonable amount) and results in a lot of unreacted spongy metal powder which settles out, eventually shorting the electrodes. Really the easiest and fastest way to make copper or nickel sulfate is to first dissolve the metal down in hydrochloric acid with H2O2, then add sulfuric acid to the resultant chloride with the liberation of HCl.

12AX7 - 18-4-2008 at 01:55

Use NaCl electrolyte, forming Cu2O and Ni(OH)2.

Hmm, I wonder if there's a way to use that already. Cu(I) stays in solution, at least until the pH rises. Eh, I think that pretty much reduces to the method I suggested above, i.e., cook with metal to reduce copper to Cu(I).

Tim

-jeffB - 18-4-2008 at 05:27

Quote:
Originally posted by UnintentionalChaos
I've isolated a small quantity of crude nickel carbonate from US quarters this way although it is a lot of work for the yield. Pottery supplier may be a much better alternative if you have access.


Whoa, that's an expensive way to get at nickel! US quarters, dimes, and half-dollars are made from an alloy of 75% Cu, 25% Ni clad over a solid copper core -- it works out to about 8% total Ni. (The weights of these three coins are proportional to their face value, so the cost per unit weight is the same whether you use dimes, quarters, or halves.)

Nickels are made of solid 75/25 Cu/Ni, so they contain a lot more nickel relative to their value -- in fact, according to http://www.coinflation.com each five-cent piece bears about seven cents worth of metal. (Nickel goes for a bit more than three times the price of copper, so that works out to about 3.5 cents worth of each metal.)

Coin compositions and weights here:

http://www.usmint.gov/about_the_mint/index.cfm?flash=yes&...

Each nickel contains 1.25g of Ni, 3.75g of Cu. By contrast, a quarter contains less than .5g of Ni.

Finally, remember that there's currently a US law against "melting or treating" pennies and nickels. Otherwise, metal providers and consumers would surely be buying their copper and nickel from the US Mint in the form of nickels at an effective 30% discount. I obviously don't intend to promote illegal activity here, and I'm sure that's why @UnintentionalChaos used quarters, which as far as I can tell are still fair game. :D

UnintentionalChaos - 18-4-2008 at 07:27

I said a small quantity of nickel carbonate...just enough for a few micro experiments. I took out the quarters when they ran out of cupronickel cladding since the dissolution without heating or strong oxidizers is very slow (I was actually using sulfuric + H2O2 ---> carbonates ---> acetates ----> treat with lead ---> treat with sulfate ----> treat with carbonate, but the idea is the same) I only used $1.50 in quarters...an advantage in that quarters have higher surface area than nickels. I really dont think the government is going to come after you for dissolving a few nickels. They only really seem to have a problem when you are attempting to resell it at a higher value. Coin mutilation beyond that law isnt illegal unless you're trying to make it more valuable than its face value.

[Edited on 4-18-08 by UnintentionalChaos]

Mr. Wizard - 18-4-2008 at 07:35

A little off topic comment about nickels and pennies. They are the only current coins worth their weight or their value, hence they will disappear soon. As a few previous posters have noted the value of the 75Cu/25Ni Nickel is over the stated 5 cents, The pre 1982 Copper penny is worth about 3x it's face value, and even the post 1983 Zinc penny is worth more as Zinc than as a penny. The 'fix' will either be to make plated slugs of steel or round everything off to not require pennies.

I bow to Kilowatt's experience in actually doing the electrolysis. I have not tried it. Yanked back to the real world by real experience.;)

LSD25 - 18-4-2008 at 08:47

If you come up with a good way of separating those two that does not involve froth flotation, I know of at least one major international corporation that will be excited. Of course, you could probably dry it out, form the citrate of the nickel, heat and remove the copper oxides from the metallic nickel. That would be one way to get to the end, I don't believe that copper is quite so easily reduced as nickel, is it? Otherwise, isn't one of 'em magnetic?

kilowatt - 18-4-2008 at 09:30

Quote:

Otherwise, isn't one of 'em magnetic?

Only in metallic form, below its austenitic temperature. I'm afraid that doesn't help much.

not_important - 18-4-2008 at 20:51

Rather than using an electrolytic cell to dissolve both metals, use it to electrorefine one of them.

It don't remember which is easier, believe it's copper, but the idea is to set up the cell to electroplate from the mixed metal electrode to one that will be one of them in fairly pure form. Metals more electropositive than the target go into solution, and thus usually require addition of makeup acid to match the dissolution rate, more electronegative metals don't dissolve - generally falling to the bottom. Because of that the anode should be mounted below the cathode.

The voltage used is quite low, lower than would be used for inert anode electroplating. A volt or two is typical, too high of a voltage results in loose deposits and other metals plating with the target one. High current speeds things up.

The Mond process (carbon monoxide) would remove most of the nickel, provided you used the prober procedure to prepare the metals powders. An interesting process, but as already stated a rather toxic one.

The reduction of Cu(II) to Cu(I) and allowing it to precipitate as the chloride looks pretty practical. If you were to convert the bulk alloy into fairly small particles, filings or small shot for example, you could partially dissolve much of the nickel by using HCl + H2O2 or HCl + Cl2, adding more alloy to reduce the copper and dissolve more nickel, using a little SO2 to force the reduction to completion, and diluting to precipitate CuCl. You can't get all the nickel out of the alloy with that method, as the copper plating back out onto the nickel will protect some of it.

The copper acetylsalicylate approach sounds interesting, but consider that Copper (II) acetylsalicylic acid hydrolysis easily and copper(II) salicylate is soluble.

Copper sulfate decomposes to the oxide at about 650 C, nickel sulfate at about 850 C. If there is not an interaction between the two, this could work to separate the two be heating the sulfates to 700-750 C, cooling, and leeching out the nickel sulfate.

Nicodem - 19-4-2008 at 00:49

This topic of Cu&Ni salts separation was already discussed in another thread. At that time I proposed to dissolve the alloy in nitric acid, evaporate to dryness under vacuum and triturate the solids with acetone to remove the hydrated copper nitrate which is well soluble in it. The remaining crude hydrated nickel nitrate should be further purified, depending on how pure you need it. However, up to day I still have not checked if perhaps nickel nitrate is too soluble in acetone and since my Ni(NO3)2*6H2O is stored in some box somewhere in the basement I'm not particularly motivated in doing a test. However, a home page of a Chinese chemical seller says:
Quote:
Nickel nitrate [2008-04-07]
Molecular Formula: Ni(NO3)2•6H2O

Molecular Weight: 290.81

Physicochemical Properties: Nickel nitrate hexahydrate is a green monoclinic system crystal. Specific Density: 2.05, Melting Point: 56.7°C, Boiling Point: 137°C. When heated at 56°C, it decomposes into trihydrate by losing three water molecules and then converts into anhydrous salt at temperature of 95°C. It easily dissolves in water, and is soluble in alcohol, and slightly soluble in acetone.

Now compared to the relatively high acetone solubility of Cu(NO3)2*2.5H2O, this might give a feasible way of obtaining a crude Ni(NO3)2•6H2O out of the mixture. Recrystallizing it from ethanol or whichever suitable solvent could make it pure enough for most uses.

EDIT: According to Phys. Rev. (Series I) 19, (1904) 156-165, solution of nickel nitrate at 1/10 normality in acetone is possible. They do not report the solubility but they used this as their maximum concentration for conductivity measurements. Unfortunately, as is typical of physicists, they do not mention if it was hydrate or anhydrous, but we are let to assume it is anhydrous. Such solubility is quite a lot since one liter acetone would dissolve at least 18 g nickel nitrate. Still the solubility is surely considerably lower than for hydrated copper nitrate, but losses would be large for purifying a material that contains nickel nitrate as the minor component. There is still the possibility that the hexahydrate has a much lower solubility.

[Edited on 19/4/2008 by Nicodem]

12AX7 - 19-4-2008 at 01:42

That reminds me, copper chloride is somewhat soluble in acetone as well.
Quote:
Originally posted by kilowatt
Only in metallic form, below its austenitic temperature. I'm afraid that doesn't help much.


Nickel doesn't have an austenite phase, it's the same alpha-Ni from room temperature to melting point. You're thinking of steel (iron with <2% carbon). Nickel does have a curie point (around 400C IIRC), much lower than that of iron's, but this is an electronic phase change.

Tim

vulture - 19-4-2008 at 02:28

I can tell from personal experience that NiCl2.6H2O does indeed dissolve into acetone and ethanol to a fair amount. The solution will get a distinct green tint. I could confirm that with dimethylglyoxime if you want.

kilowatt - 19-4-2008 at 13:15

Quote:

Copper sulfate decomposes to the oxide at about 650 C, nickel sulfate at about 850 C. If there is not an interaction between the two, this could work to separate the two be heating the sulfates to 700-750 C, cooling, and leeching out the nickel sulfate.

How volatile are either of these sulfates at these temperatures? Do you have any experience with heating nickel sulfate hexahydrate to such a temperature? I am under the impression that it loses all of its water at just over 100°C, which would be nice because I would hate to have my SO3 (from copper sulfate) come out hydrated. I had not considered this method before because none of the data on nickel sulfate seems to agree on a boiling or decomposition point. It would be the best/easiest though if the anhydrous nickel salt really is stable to 850°C.

Quote:

I can tell from personal experience that NiCl2.6H2O does indeed dissolve into acetone and ethanol to a fair amount. The solution will get a distinct green tint. I could confirm that with dimethylglyoxime if you want.

Knowing the solubility of both copper and nickel chlorides in DMG would be great, even though I don't have any DMG at this time. I may have to experiment with other mildly polar solvents as well.

12AX7 - 19-4-2008 at 19:23

DMG isn't for solvent, it's a well-known nickel test, making a pink precipitate.

Vulture: I'd take your word for it qualitatively, but if you'd take it one step further and do a quantitative test to determine NiCl2 solubility in acetone, that would be great.

Tim