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Author: Subject: Cu++ and Ni2++ separation
Sedit
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[*] posted on 29-3-2010 at 17:28


Just thought I would upload a couple pictures of the seperation process. I am far from documenting the entire thing which I will produce a very full writeup and post it over at The Vespiary since we need more serious contributions over there.

Here is what it looks like after the addition of the Chloride salts to (aq)Ammonia Hydroxide. This picture does not show the light greenish blue precipitate to well since there is only a little bit of the salt in this picture. This was just a preliminary run to test to see if my theory was right but you can still see a small layer at the bottom of the flask. The deep blue color is from Copperamine complex formed with the ammonia dissolved in the solution.

coppercomplex.jpg - 8kB

The solution is decanted and washed several times with H2O to remove the Cu from the mixture. At first the precipitate is light blue but the more you wash it the greener it becomes until you have something that looks like this.
nickle oxide.jpg - 7kB

Sorry its not the best quality photo but I will work on better versions when I produce the final draft of my writeup which I will hopefully quantify the entire composition of the coins recovering the metallic Copper as well the Nickle salts.

Once you have got it all washed nicely instead of filtering which is a pain in the ass I think your better to just add HCl at this point until you have a transparent green solution and filter the NiCl2 and evaporate.

So what do yall think? Easier then dicking around with Selective precipitation using citrate salts or whatever was mentioned dont you think?





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[*] posted on 29-3-2010 at 21:34


Well, no, as nickel forms complexes with ammonia as well

From A Text-Book Of Inorganic Chemistry Vol-X Metallic Ammines by J.Newton Friend

http://www.archive.org/details/textbookofinorga025530mbp

an example


tetraammino-nickel-sulfate.png - 15kB

and Yes, because I'm sure the copper complex is more stable than the nickel, but I don't know by how much more; I question the completness of separation.



[Edited on 30-3-2010 by not_important]
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[*] posted on 29-3-2010 at 23:00


As far as completeness I can reassure you that there is NO Cu left in the precipitated oxide. I know this because I turned it back to Nickel chloride and precipitated it again showing No! dark blue solution this time. Only the precipitation of a green compound with the properties of NiO. Plus a flame test and a prism assures me there is trace... if any Copper left.

This does however explain why when I first begin to base it with Ammonia hydroxide it forms alot of precipitate but further addition yeilds less product and a darker blue solution. Presumably this is due to excess Ni turning into the amine complex as well but I am unsure at the moment.

In all honesty I doubt the validness of this reference after my personal observations. The nickle complex seems to form slightly if at all.





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[*] posted on 29-3-2010 at 23:18


`twould be useful to be able to peruse these:

http://openlibrary.org/b/OL186958M/formulas_and_stabilities_...

http://openlibrary.org/b/OL20031085M/Metal_ammine_formation_...

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[*] posted on 30-3-2010 at 00:10


I suspect the difference in solubility between Ammonium Nickel Sulfate and any similar copper double salt or complex may be useful for separation, but am having some trouble locating actual solubities for double salts....Simple enough to try.
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[*] posted on 30-3-2010 at 04:28


US Patent 4005031 - Nickel peroxide oxidizing agent
http://www.patentstorm.us/patents/4005031/description.html

This is the patent that gave me the idea on how to perform the seperation. They use Ammonia hydroxide to neutralize a solution of Nickle chloride to precipitate it onto a carbon support as a finely divided oxide. They proceed further into the formation of Nickel peroxide but no where do they make mention of the Ni forming a complex with the NH3. I have read papers describing such a complex and it confuses me as to why they would make no mention of it in the patent.

It makes me wounder if the Chloride salt does not form this complex by precipitating the oxide at such a rate that it does not have the will to form the complex in the Oxide state. I know what im getting is NiO but the qustion is if im losing yeilds while attempting to wash the Copper out.





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[*] posted on 30-3-2010 at 05:01


I think the reason there is no mention of the complex lies in They use Ammonia hydroxide to neutralize a solution, meaning they're not using an excess of NH3. In the same way if you add NH3 (aq) to a solution of a copper salt, watching the pH or calculating the needed amount so as to just react with all the anions, you get Cu(OH)2 ppt; more NH3 is needed to get the complex. And I'm sure the Cu complex is more stable than the Ni, but I don't know how much more.

The reference I gave lists several complexes of NiCl2 with NH3. Like copper and cobalt the anhydrous salr reacts with dry ammonia, and hydrated forms of M[(NH3)6Cl2 can be had from solutions of the halide on treating with ammonia; the Ni compound is obtained in crystals by addition of NH3 and NH4Cl to reduce its solubility, while the copper one will form crystals by concentration in an atmosphere of ammonia.

And NH3 precipitates Ni(OH)2 at ordinary temperatures, heat converts it to NiO while oxidation, either chemical or electrolytic, converts it to NiO(OH) or mixed Ni(III)/Ni(IV) hydrated oxides.





Attachment: BB643.pdf (120kB)
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[*] posted on 30-3-2010 at 10:49


If it is indeed the case that Nickel forms a complex with Ammonia as well then what would be the reason that after isolation of the NiO and converting it back to the chloride I once again precipitated it with more Ammonia hydroxide. This time (even after a couple attempts) there was no blue color at all just more precipitated NiO.

If it does indeed form a complex like stated it should show some signs of a blue solution like what is reported correct?

I am possitive that this is Nickle because I reduced some last night and the powder was attracted to a magnet so theres no doubt at all that Nickle is what im getting. Do you think theres a possibility that the solubility of the Nickle ammonia complex is very low and the precipitate is the complex and not the Oxide like im assuming. I highly doubt it myself but just trying to figure this out because observation is going against all the literature I'v been reading on Nickle salts reactions with ammonia.

I think perhaps next run I am going to wash the precipitate as before and reflux it in H2O for a while and see if I can detect the presence of NH3 in the off gas.





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[*] posted on 31-3-2010 at 06:44


You shouldn't be getting NiO as a precipitate with NH3 unless the NH3 (aq) is dilute and the solution is hot.

The amino complex can be crystallised from aqueous NH3, but it is described as being distinctly blue.

It doesn't take pure Ni for the reduction product to be attracted to a magnet, a number of nickel alloys are magnetic. Magnetism is a good indicator of having a high content of Fe/Co/Ni and in alloys possibly Mn. The colour of the ppt argues against a high content of anything but Ni, as Fe and Mn would be sensitive to oxidation by air to give red-brown colours.

What coins were the starting point for this?

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[*] posted on 31-3-2010 at 11:04


Quote: Originally posted by not_important  
(cut)
It doesn't take pure Ni for the reduction product to be attracted to a magnet, a number of nickel alloys are magnetic. Magnetism is a good indicator of having a high content of Fe/Co/Ni and in alloys possibly Mn. The colour of the ppt argues against a high content of anything but Ni, as Fe and Mn would be sensitive to oxidation by air to give red-brown colours.(cut)

There is also a composition range of Cu-Mn alloys, containing no Fe or Co or Ni, that is ferromagnetic, although not as strongly as the latter. In fact, I have what appears to be a bronze lamp-standard made of this ferromagnetic Cu-Mn alloy. However, the reason why V, Cr, Mn, Cu, and in general nearly all their alloys, are not ferromagnetic, is because the electron spins of their unpaired "3d" electrons align in parallel but opposite directions so that their magnetic moments cancel out, which is called antiferromagnetism. In Fe, Co, and Ni, these unpaired electrons spins are aligned parallel in the same directions in magnetic "domains" - ferromagnetism.

Of course, the same also phenomenon occurs with the unpaired 4f and 5f electrons in the lanthanide amd actinide metals, also conferring ferromagnetism, especially in those about the middle of the two series where the numbers of unpaired electrons are at a maximum.
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[*] posted on 31-3-2010 at 12:43


Quote: Originally posted by not_important  
You shouldn't be getting NiO as a precipitate with NH3 unless the NH3 (aq) is dilute and the solution is hot.

It doesn't take pure Ni for the reduction product to be attracted to a magnet, a number of nickel alloys are magnetic. Magnetism is a good indicator of having a high content of Fe/Co/Ni and in alloys possibly Mn. The colour of the ppt argues against a high content of anything but Ni, as Fe and Mn would be sensitive to oxidation by air to give red-brown colours.

What coins were the starting point for this?



The coins are common US nickels dissolved sometime ago in HCl+H2O2 since a freind asked how to go about seperation I decided to try. The salt sat there and turned into the oxides over time since I could never figure a means of seperation. Before this I just added HCl to the oxide powders afford a deep greenish almost brown(very dark) salt solution.

There is trace amounts of Fe due to the water I was using for the initial test but this can be seen to precipitate out as a fine brown dust on the Nickel on sitting. There is only trace amounts of Fe in there at best.

If it takes heat to make the oxide from the precipitate why is that not mentioned in the patent (did I over look something)and also why does this precipitate not show the blue color mentioned but appears to have all the properties of NiO?

Im going to setup a spectroscope that I can take pictures with when I get time to see if I can post the spectrum of a flame test.

I got ALOT more dissolving in H2SO4 right now so I will perform more complete test when thats done but it is going boringly slow.





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[*] posted on 31-3-2010 at 18:31


Quote: Originally posted by Sedit  
...(much snippage)...
If it takes heat to make the oxide from the precipitate why is that not mentioned in the patent (did I over look something)and also why does this precipitate not show the blue color mentioned but appears to have all the properties of NiO?

Im going to setup a spectroscope that I can take pictures with when I get time to see if I can post the spectrum of a flame test.

...

Well, possibly because
A) The patent says
Quote:
As used throughout this application, the term "nickel oxide" includes nickel hydroxide, nickel monoxide (NiO), nickel sesquioxide (Ni2O3) and nickelic tetraoxide (Ni3O4) and mixtures of two or more of the foregoing.

B) The patent is about producing "nickel peroxide" and not nickel oxide or hydroxide, which is mearly an incidental step along the way.

And ever reference I can find from Mellor and Friend to a recent one like the attached PDF all say that you get Ni(OH) from adding alkali hydroxides, including aqueous NH3, to solutions of nickel salts. They also state that Ni(OH)2 is soluble in excess ammonia, which is why I asked about the coins; the precipitate just does not sound like pure Ni(OH)2.




Attachment: On the existence of a nickel hydroxide phase which is neither a nor b.pdf (143kB)
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[*] posted on 1-4-2010 at 19:21


Any suggestions on how to test between the hydroxide and the oxide when I attempt to perform this on a larger scale once they finish dissolving?

It's all got me a little curious, more so then I thought I would be over it, because it just don't seem to add up to reports. I guess is possible for the Cu to be pushing it out simular to the common ion effect but I'll wait till I have more material to work with before making any assumptions.

Thanks for your input its been helpful as usual. Feel free to add anything else as you find it.

Various Ni(II) reactions:
http://www.public.asu.edu/~jpbirk/qual/qualanal/nickel.html

These explain a little of what I have been seeing and like you stated before the Ni(OH)2 dissolves in Excess NH3. At first there is a large amount of precipitate with a faint blue color however more NH3 solution formed the deep blue you see leaving less precipitate. I think that correct ratio is key to getting useful seperation of the two compounds.

I also will precipitate some of the solution next time time NaOH and use this as a standard in following test knowing that the NaOH should precipitate both the Ni and Cu alllowing me to know exactly how much I should be recovering in the way of Ni(OH)2





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"I see a lot of patterns in our behavior as a nation that parallel a lot of other historical processes. The fall of Rome, the fall of Germany — the fall of the ruling country, the people who think they can do whatever they want without anybody else's consent. I've seen this story before."~Maynard James Keenan
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