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)?
The sodium salt of acetylsalicylic acid should precipitate the copper selectively. Then just treat with HCl to recover the CuCl2 and acetylsalicylic
acid.
This exact question has also been answered in previous posts, if I remember correctly--you might want to try the search function.martin21 - 15-5-2008 at 19:04
damn I feel stupid... I didn't took the time to search but thanks anyway guys!woelen - 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.12AX7 - 15-5-2008 at 23:13
And produces I2, correct? Could be irritant unless reduced with e.g. H2SO3.
Timwoelen - 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 .12AX7 - 16-5-2008 at 10:03
Oh, skimmed over that :-[Ozone - 16-5-2008 at 17:44
Cation exchange resin?
Cation-exchange separation of copper from zinc, nickel and lead :
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).
Cheers,
O3Jome - 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]ciscosdad - 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. JohnWW - 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).bbartlog - 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. Mr. Wizard - 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.
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.garage chemist - 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. Sedit - 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.ciscosdad - 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. not_important - 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.
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.
ChrisWhewell - 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.ciscosdad - 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?
bbartlog - 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
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.
Sedit - 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.bquirky - 26-3-2010 at 09:07
weird could it be some detergent in the ammonia hydroxide?Sedit - 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.
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.
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?not_important - 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
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]Sedit - 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.not_important - 29-3-2010 at 23:18
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.Sedit - 30-3-2010 at 04:28
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.not_important - 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) This file has been downloaded 688 times
Sedit - 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.not_important - 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.
(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.Sedit - 31-3-2010 at 12:43
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.not_important - 31-3-2010 at 18:31
...(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.
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.
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
.