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martin21
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[*] posted on 15-5-2008 at 15:25
Cu++ and Ni2++ separation


Hi!

I have a solution of copper (II) and nickel (II) chloride and I would want to separate both ions. I thought about selective precipitation of the respective sulfides or iodides.
I thought about something else... if a add ammonia to make a Fehling reagent solution than add a reductive sugar, would nickel (II) interfere with the reduction of the copper (II) in copper (I)?

thanks!
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Mr. Wizard
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[*] posted on 15-5-2008 at 16:00


http://www.sciencemadness.org/talk/viewthread.php?tid=10372&...
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[*] posted on 15-5-2008 at 17:20


The sodium salt of acetylsalicylic acid should precipitate the copper selectively. Then just treat with HCl to recover the CuCl2 and acetylsalicylic acid.

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

This exact question has also been answered in previous posts, if I remember correctly--you might want to try the search function.
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martin21
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[*] posted on 15-5-2008 at 19:04


damn I feel stupid... I didn't took the time to search but thanks anyway guys!
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[*] posted on 15-5-2008 at 23:10


Yet another option is to use potassium iodide or sodium iodide. These reduce the copper(II) ions to copper(I) and copper(I)iodide precipitates with a quantitative yield. The nickel(II) ions remain in solution. The copper(I) iodide forms a compact and easily separated precipitate. Recovering the nickel(II) from the solution then can be done by reducing the iodine in solution again (e.g. with bisulfite) and then adding a carbonate or hydroxide and precipitating the nickel(II) are Ni(OH)2 or as a basic carbonate. The basic carbonate is easier to separate.



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[*] posted on 15-5-2008 at 23:13


And produces I2, correct? Could be irritant unless reduced with e.g. H2SO3.

Tim




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[*] posted on 16-5-2008 at 04:23


Yes, iodine is formed as well, but in my post I did mention that, and I suggested reduction with bisulfite ;).



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[*] posted on 16-5-2008 at 10:03


Oh, skimmed over that :-[



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[*] posted on 16-5-2008 at 17:44


Cation exchange resin?

Cation-exchange separation of copper from zinc, nickel and lead :

http://www.springerlink.com/content/h2w64242w72n32rj

I'd get the article for you, but I only have access from 1998-.

IIRC, we separated Cu and Ni on cationic ion exchange resin which was sequentially eluted with HCl of increasing strength (0.1?N HCl then 2M HCl?) (sophomore Analytical Chem. lab, too long ago).

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O3




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[*] posted on 20-5-2008 at 15:32


Simply boiling with Cu-powder should give symproportionation to insoluble CuCl.

CuCl2+Cu---->2CuCl

If you want the CuCl2 back, perhaps it could work to dissolve the CuCl with ammonia then add a carbonate to get CuCO3 which could be reacted with aq. HCl to CUCl2.

[Edited on 20-5-2008 by Jome]
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[*] posted on 14-12-2009 at 18:28


Resurrecting an old thread here:

I've been toying with the idea of an economical source of Nickel salts. Since I have access to Nitric Acid (68%) the plan is to dissolve some scrap Monel Metal which I believe to be 75% Ni, the remainder Cu.
I've seen Cuprous Chloride suggested, and Cuprous Thiocyanate, among others.
On searching through a number of analytical Chemistry books, the preferred method of separation (in late 19th/ early 20th Century) was using H2S. Cupric Sulphide being insoluble in dilute acids.
The scheme I had in mind was to dissolve in Nitric then convert to mixed sulphates via carbonate then use H2S to precipitate the copper. I'm suspicious of the Nitrate's ability to oxidise any sulphide present, hence the conversion to sulphates.
Is anyone aware of any reason that the H2S method would not scale up to kilogram Quantities? (aside from the obvious problem with smell)
I would be interested to hear from any of the guys who posted earlier and perhaps succeeded.

As an aside, one of the books referred to the use of ammonium sulphate to separate the two. Nickel sulphate is evidently insoluble in a saturated solution of Ammonium Sulphate, and crystallises out as a double salt, leaving the copper in solution. If anyone is interested I'll hunt up the reference.
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[*] posted on 15-12-2009 at 00:01


That should also work on cupro-nickel coins, of which I have a large amount of, recently withdrawn from circulation and replaced with smaller coins. They are 75% Cu and 25% Ni. From about 1910, to 1946 in New Zealand and the UK, and until 1963 in Au$tralia, they were of 50% Ag and the rest 37.5% Cu and 12.5% Ni. (I also have a lot of these coins).
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[*] posted on 15-12-2009 at 08:08


American nickels are also 75% copper and 25% nickel, if I recall correctly. If it's really just the nickel you're after I'm not sure this ends up being particularly cheap (though it doesn't seem too awful), but if the copper salts are also of value to you then it's a nice source. Coinage in general is a good source of some metals (for US coinage past and present: zinc, copper, nickel and silver... and gold I suppose). The face value is often not a big markup over the value of the metal, and while you may need to do some purification/separation you at least start with reliable information about the composition.
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[*] posted on 15-12-2009 at 13:35


On a previous post somebody mentioned precipitating copper from the sulfate by using vitamin C, Ascorbic Acid. I tried it in a small 10 cc solution and got a pink metal precipitate. I haven't tried it with a nickel solution, so I don't know. Silver will also precipitate with vitamin C , IIRC.

https://www.sciencemadness.org/whisper/viewthread.php?tid=26...
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[*] posted on 15-12-2009 at 21:50


H2S is possibly the most economical way. If it is not in concentrated nitric acid... that is, if it is a bit dilute, then there is no need to convert to sulphate.

H2S in dilute nitric acid solution will precipitate what are called the heavy metals - silver, lead, copper and mercury. In neutral solution H2S will precipitate nickel, zinc, cadmium, arsenic and probably a few more. This was all part of a systematic scheme for identifying metal cations that I was taught too many years ago.
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[*] posted on 15-12-2009 at 22:10


Convert the metals to a mixed chloride solution via carbonate (or dissolve the alloy in HCl/H2O2, then destroy excess H2O2 by boiling), then add bisulfite solution (or bubble in SO2) to precipitate the copper as CuCl.
Since CuCl is soluble in HCl and chloride solutions ( [CuCl2]- complex formation), and HCl/Cl- is inevitably being generated by the reduction, you have to work in relatively dilute solution to get complete precipitate of copper.
You can use a pure solution of CuCl2 or CuSO4 in dilute HCl with bisulfite to find out how dilute the solution needs to be in order to get quatitative precipitation of CuCl (colorless solution).

Don't wash precipitated CuCl with ordinary tap or distilled water, it'll partially redissolve. Use SO2-saturated water or acidified bisulfite solution.




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[*] posted on 15-12-2009 at 22:19


Just a thought and Im not sure what it means right now but I took the Acidified, with HCl solution obtained from H2O2/HCl, and subjected it to electro reduction. The copper colored precipitate was not attracted to a magnet however when the metals are reduced with Aluminum it will react to a magnet and appears darker then the product from the cell.

Does the Ni not precipitate when subjected to an Electrochemical cell as the free metal like Cu does? I have not put much thought into it at the moment since I just recovered it a few hours ago and im short on time so forgive my ignorance.





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[*] posted on 15-12-2009 at 22:34


@Paddywhacker:
Nice to find someone who admits to some training in the procedure.
Is there any limits on concentration for the copper/nickel solution? (ie, must it be diluted to any degree?)
I guess 500g of Monel metal as mixed nitrates in 2 or 3 litres of dilute nitric acid is what I have in mind. I assume the dilution of Nitric that no longer attacks the metal would be a good starting point?

Would not be real happy about doing 20 litre batches!

On hols for a few weeks, so not sure when I'll be reading any posts. I will be back though.
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[*] posted on 18-12-2009 at 09:06


You don't want it too concentrated as the CuS will make it thick, with poor mixing and reaction completion.

Another thing to consider is the pH. Generally this is around 0.5 to 2, too strong and not all the copper precipitates while too high and nickle drops out as well. As acid is being liberated as the sulfide is formed, the solution becomes more acidic; if too concentrated the pH may drop too low.

There's a good example of the numbers near the end of: http://pages.towson.edu/ladon/solprod.html


Once you've removed the copper, you need to do so for the iron and manganese that make up about 5% of the monel (wasn't he a character in Superboy comics?). Boil the filtrate to drive off H2S, bubbling air through while boiling helps. If you didn't use HNO3 to dissolve it, add a bit of HNO3 or H2O2; you want the iron as Fe(III).

Take about 1/10 of the total solution and precipitate it using NaOH or Na2CO3, try to get a pH of 8 to 9. Wash the precipitate well to remove sodium ions, and keep it wet.

To the remaining 9/10 add aqueous ammonia until a precipitate starts to form. Now add the precipitate from above, stir well, and slowly bring to a boil. Again, it you can bubble air through the solution do so.

The precipitate should turn brownish as the nickle hydroxide or carbonate precipitates Fe(III) and manganese as a mixture of oxidation states, giving a mix of hydrated oxides/hydroxides/basic carbonates; the nickle going into solution. Bubbling air helps push Mn(II) to Mn(III) and mixed Mn(III)-Mn(IV) states, which have much less soluble hydroxides/oxides.

Keep up the stirring for maybe an hour, keeping the solution hot ~80 C. The allow to cool and settle, decant the liquid through a filter. Stirring the precipitate with strong aqueous ammonia will dissolve out some of the excess nickle, this can then be added to the filtrate forming a ppt of Ni(OH)2; acid can be added to dissolve this.

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[*] posted on 22-12-2009 at 15:11


I didn't read all the replies and apologize if this is duplistic.

If you like to get dirty, consider the Lix process, whereby copper amine is extracted using kerosene.

I'd prefer simple electrolytic reduction. Nickel is above hydrogen in the electromotive series, whereas copper is below it. Control of potential ought readily give selective electrodeposition. This also means one can bubble H2 through the solution causing the Cu to precipitate, leaving the Ni in solution.
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[*] posted on 12-1-2010 at 14:48


Thanks you guys.
I found the Solubility product reference from Not_important most useful.


Now I need to pull my finger out and do it.
Gotta make a Kipps apparatus for H2S, then make some strong ammonia soln.
Never ends does it?

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[*] posted on 25-1-2010 at 18:49


A couple of notes on this and a question:

First of all, I don't think that precipitating CuCl is a particularly practical way of removing all the copper from solution. I've tried this approach, starting with a solution of mixed HCl/NiCl2/CuCl2, boiling with copper, and diluting (or cooling) to precipitate the CuCl. While it's an easy way of producing CuCl, it would require a very good handle on the stoichiometry to make sure that you got rid of all the HCl at the same time that you had reduced all the CuCl2 to CuCl, and even then CuCl isn't really insoluble enough for you to be able to say that it's all gone. And that's not even considering the possibility that it would form some sort of complexes with the NiCl2.
For what it's worth I've washed CuCl with distilled water and seen no color change. It's sensitive to oxygen but not so much that a couple of quick rinses will ruin it if you are going to use it right away; as far as I can tell the small portion that oxidizes ends up dissolving and getting rinsed away (the wash turns a very light green while the CuCl stays bone white).

Second, here's an approach I plan to try instead, and I wanted to know whether anyone sees a problem with it:
1) Precipitate both Ni++ and Cu++ as hydroxide (by addition of NaOH)
2) Filter, dry and heat the precipitate to 70C to transform the Cu(OH)2 into CuO (while leaving the nickel hydroxide unchanged)
3) Pulverize and wash the resulting powder with a volume of water, acidified to pH 6 with a small amount of ZnCl2

As far as I can tell looking here: http://www.iapws.jp/Proceedings/Symposium05/264Palmer.pdf and here: http://www.iapws.jp/Proceedings/Symposium08/491Palmer.pdf, CuO is about a thousand times less soluble than Ni(OH)2 in water at pH 6 and room temperature. So I should be able to wash the nickel out of the copper.
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[*] posted on 28-1-2010 at 11:28


This may have allready been mentioned

But i think you might beable to simply use gravity sepperation put the solution in a tall narrow container let it sit for a fair while and then draw off the top and bottom fractions. i have never acctuly tried to sepperate Cu and Ni in this way before but i have observed this sepperation in solution when desolving coins.

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[*] posted on 25-3-2010 at 11:54


I have successfully seperated the Nickle from coins by precipitating it from a mixture of the Chloride salts using ammonia hydroxide keeping the Copper in solution as a complex with the ammonia.

However there is something strange that I can not account for perhaps someone here could explain to me. After precipitation I wash the NiO with H2O but after sitting for a while all the NiO rises to the surface of the water because it has been generating a gas of somekind and rising up. This makes no sense to me. Is it oxidising water? If so to what and what would the REDOX look like. Theres no metallic Nickle so that don't make sense anyway.

What could this gas be and why is it forming?

My only conclusion is its acting simular to froth flotation and the Nickle is complexing with gasses in the H2O effectivly degassing the solution but thats the best theory I could come up with yet.





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[*] posted on 26-3-2010 at 09:07


weird could it be some detergent in the ammonia hydroxide?
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