| Pages:
1
..
5
6
7
8
9
..
11 |
Sedit
International Hazard
Posts: 1939
Registered: 23-11-2008
Member Is Offline
Mood: Manic Expressive
|
|
Ok Pictures coming soon but the kids just got home from school and I have to take care of the important things in life before getting pictures...
I dried the Chloride made from that green precipitate you see in the large jar above leaving me with a green powder.
I dissolved this in a minimum amount of water and added NaOH prills like I stated last night only to yeild a brown precipitate that is the color of
MnO. This I don't understand at all.
Ok, today with my experiment I again dissolved the green powder in a minimum amount of water as before and added a NaOH solution this time. This time
I got a blue precipitate then with a few more drops of NaOH solution the precipitate turned green like that of before. But here's where something odd
took place... The solution above the precipitate took on a blue color exactly like that of the Ammonia complex. I feel this is due to Ammonia being
released from Ammonium Chloride that was still in the mix but there is the possibility that the precipitate itself IS a Nickle amine complex (or
Copper) and that's why things are so weird with this reaction.
I suspected the Cl- ion played a roll in the past hence the reason I asked Blogfast to add NH4Cl in hopes that would produce some sort of results but
it didn't.
I'm starting to wounder if the Nickle amine complex is not very soluble in a solution of the Copper amine complex.
I'm just musing at the moment since I just performed the experiment and got alot of other things to do right this minute so keep the flames at a
minimum. Later tonight I will get the pictures and further this discussion.
Knowledge is useless to useless people...
"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
|
|
|
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
Quote: Originally posted by Sedit  |
I dissolved this in a minimum amount of water and added NaOH prills like I stated last night only to yeild a brown precipitate that is the color of
MnO. This I don't understand at all.
[…]
The solution above the precipitate took on a blue color exactly like that of the Ammonia complex. I feel this is due to Ammonia being released from
Ammonium Chloride that was still in the mix but there is the possibility that the precipitate itself IS a Nickle amine complex (or Copper) and that's
why things are so weird with this reaction.
[…]
I'm starting to wounder if the Nickle amine complex is not very soluble in a solution of the Copper amine complex.
|
Adding prills of NaOH leads to local overheating due to the high solvation heat of NaOH: Cu(OH)2 loses water to form (brown black) CuO easily.
Second point: could be residual ammonia (did you smell any?) but more likely it’s cuprate anions
(Cu(OH)<sub>4</sub><sup>2-</sup> from Cu(OH)2 + 2 NaOH); copper is slightly amphoteric and Cu(OH)2 dissolves somewhat in
strong NaOH to a cobalt blue cuprate, almost indistinguishable from the corresponding ammine complex.
Third point: no. Remember how a 50/50 Cu(OH)2/Ni(OH)2 mix dissolved effortlessly in 1.5 M NH3. The nickel ammine complex’ solubility is unaffected
by the copper ammine complex.
Try using fairly dilute, cool alkaline solutions, like 1 M (40 g NaOH per liter).
|
|
|
Sedit
International Hazard
Posts: 1939
Registered: 23-11-2008
Member Is Offline
Mood: Manic Expressive
|
|
I believe you are correct and that is Copper oxides due to over heating, the blue precipitate in the background is the one where I use dilute NaOH.
The color of the solution cleared up so I will repeat the experiment in an attempt to make sure it stays so I can photograph it.
Its odd as you can see the precipitate is blue but when I first started to add the NaOH solution the precipitate was green and only turned blue after
further addition of NaOH.
Im going to have to make a standard NaOH solution so I can determine the exact concentration at any point in time as this is important. Here are the
photos. I will perform more experiments later. Its dinner time now 
WTF, I cant attach files, is it showing up on yalls computers because I don't see it on mine
[Edited on 23-11-2011 by Sedit]
[Edited on 23-11-2011 by Sedit]
[Edited on 24-11-2011 by Sedit]
Knowledge is useless to useless people...
"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
|
|
|
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
As far as I’m concerned the mystery is now solved, and it really is ‘game over’.
I distinctly remember when I first precipitated the mixed hydroxide that it was this verdigris green/blue colour and that on adding the requisite
small excess of NaOH the stuff turned blue. This is likely to happen anytime there’s sufficient chloride present:
2 CuCl2(aq) + 3 NaOH(aq) === > Cu2(OH)3Cl(s) + 3 NaCl(aq)
On adding some more NaOH the transition from verdigris to blue occurs:
Cu2(OH)3Cl(s) + NaOH(aq) === > 2 Cu(OH)2(s) + NaCl(aq)
As I wrote above, I feel all the green precipitates including the test tube ones we’ve seen are in fact the copper hydroxychloride, probably with
co-precipitated Ni(OH)2.nH2O in it. No real separation was ever achieved, not even at the interface.
The copper hydoxychloride appears to start precipitating a bit earlier (from about pH = 4) than straight copper hydroxide but its precipitation range
(4 to 7 acc. Wiki, in accordance with my observations) overlaps too much with that of nickel hydroxyde (which starts from about 6.5) to be of any
real use as a separation method for Cu2+ and Ni2+. The only conditions in which it might be possible to separate out any nickel is if it is
present in really low concentrations (< 0.001 M) to begin with.
It’s been fun and educational… 
No, I can't see any pics.
|
|
|
m1tanker78
National Hazard
Posts: 685
Registered: 5-1-2011
Member Is Offline
Mood: No Mood
|
|
Yes it has!
Sedit, I don't see any pics either.
Tank
|
|
|
Sedit
International Hazard
Posts: 1939
Registered: 23-11-2008
Member Is Offline
Mood: Manic Expressive
|
|
Trying once again to attach the photo...
Somethings wrong folks, can someone else attempt to attach a file of somekind here to see if its the forum or something odd on my end.
[Edited on 24-11-2011 by Sedit]
Knowledge is useless to useless people...
"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
|
|
|
m1tanker78
National Hazard
Posts: 685
Registered: 5-1-2011
Member Is Offline
Mood: No Mood
|
|
I'm telling ya, create a Photobucket or similar account. Upload pics. Some allow you to edit right in the browser. Copy the url it gives you. Paste it
here (between the img tags). Presto.
If you're hell-bent on uploading the pics directly then search the net for an open source program that will allow you to resize and crop - at minimum.
Has anyone tried precipitating copper from the mixed chlorides with ascorbic acid? I tried yesterday to add AscA solution to dilute-ish Ni/Cu chloride
but nothing precipitated. I've read a few accounts of people using AscA to
precipitate colloidal copper from a Cu2+ solution but no mention of other metals (nickel in this case). What am I doing wrong?
Tank
|
|
|
Sedit
International Hazard
Posts: 1939
Registered: 23-11-2008
Member Is Offline
Mood: Manic Expressive
|
|
Photobucket and the likes are the reason forums look like shit as the years go by because the images vanish, I do not like my photos vanishing because
my memory is very very bad. So bad that I can't even remember why I have been trying to upload this photo I am about to attach without rereading the
last page or so of this thread but here it is and now I got a 30 day trial version of Adobe photoshop (could use a crack if yall got one ) I will post as many pictures as I can.
PS. The acidic precipitate (The green shit in the big jar higher up in this thread)was purely Copper complex of some sorts with no magnetic ability
after reduction with Zinc or Aluminum, If I find that there is no Nickle at all in the mix or at lest a very low percentage then that might be a
viable option for removing the Cu from the mother liquor.
Sorry for the hackish post but I just had to get this one out there as I have been away and trying to fix up my computer at the same time.
Knowledge is useless to useless people...
"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
|
|
|
Formatik
National Hazard
Posts: 927
Registered: 25-3-2008
Member Is Offline
Mood: equilibrium
|
|
I don't know if this method has been mentioned yet, but I was reading some time ago in Die Bestimmungsmethoden des Nickels und Kobalts und ihre Trennung von den anderen Elementen, which mentioned a reference of Brunck who suggests to
separate copper from nickel by precipitating the copper through sodium hydrosulfite (sodium dithionite) in acid solution.
Taking a closer look at the Brunck reference (Ann. Chem. 327 (1903) 244), who says that the separation is quantitative. Brunck points out iron, zinc,
nickel and cobalt are not precipitated through sodium hydrosulfite in acidic solution (but copper is). Zinc is said also to show not even the smallest
tendency to go into the copper precipitate (not the case when precipitation is done with H2S).
|
|
|
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
| Quote: Originally posted by Formatik | | Zinc is said also to show not even the smallest tendency to go into the copper precipitate (not the case when precipitation is done with H2S).
|
With a solubility product K<sub>sp</sub> for ZnS of about 10<sup>-25</sup> as opposed to 10<sup>-36</sup> for CuS,
zinc and copper can almost certainly be separated from each other based on solubility of their sulphides (using H<sub>2</sub>S
saturation) in acidic solutions, in the same way that it's done for Cu/Ni separation by sulphides at pH < 7. The K<sub>sp</sub> for NiS
is similar to that of ZnS.
|
|
|
Formatik
National Hazard
Posts: 927
Registered: 25-3-2008
Member Is Offline
Mood: equilibrium
|
|
This comment on zinc was Brunck's finding on a partial co-precipitation he apparently observed.
Brunck, Ann. Chem. 327 (1903) 244.
|
|
|
m1tanker78
National Hazard
Posts: 685
Registered: 5-1-2011
Member Is Offline
Mood: No Mood
|
|
I scaled up the electrolysis experiment again only this time I used a piece of cupro-nickel I found at a marine scrap yard. I hit a couple of minor
snags (fell asleep as always) so this run wasn't as clean as the previous one. By that I mean that I wasn't particularly careful to remove as much
copper as possible. I stopped the electrolysis and filtered the cell liquor to remove elemental copper and copper(I) chloride (the other snag!).
I believe I devised a simple qualitative test for copper(II) in the cell liquor. In a tiny test tube, I placed some filtered and dilute cell liquor. I
added about the same volume of DMSO and added some granules of NaHSO4. With light heating, a dark precipitate immediately formed which I presume to be
CuS.
I don't know if this method would work (be cost effective, etc) for bulk separation but my focus is on a reliable and simiple qualitative test for
copper in the nickel(II) chloride cell liquor. FWIW, the test solution smells like DMS, not H2S. After sitting for a bit, some small flat colorless
crystals have manifested at the bottom of the test tube.
Could somebody replicate this with the pure chlorides of nickel and copper and perhaps one 50/50?
CuS(?) can be seen in the test tube:
Tank
Chemical CURIOSITY KILLED THE CATalyst.
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
Those following some of the chemistry surrounding Oxalic acid and Oxalates that I have been reporting may find some interesting observations for the
separation of Cu/Ni. In particular :
1. Per "A dictionary of chemistry and the allied branches of other sciences", Volume 4, by Henry Watts, page 257, to quote:
"Oxalate Of Copper. Neutral cupric Oxalate, ....(according to Lowe, Jahresb. 1860, p. 213), is a light greenish-blue precipitate, insoluble in water,
nearly or quite insoluble in oxalic acid, but easily soluble in the neutral oxalates of ammonium, potassium and sodium. It does not give off the whole
of its water even at 120°, but decomposes at a somewhat higher temperature."
Link:
http://books.google.com/books?pg=PA257&id=lYXPAAAAMAAJ#v...
2. Same source, page 262, to quote:
"Oxalate Of Nickel, ... Greenish-white precipitate insoluble in water, soluble in ammonia and in ammoniacal salts. It dissolves also in potash,
forming a crystallisable potassio-nickel-oxalate. Neutral oxalate of ammonium dissolves oxalate of nickel, and the solution yields by evaporation
green prisms of ammonio-nickel-oxalate. On adding to the aqueous solution of this bait a small quantity of ammonia, a pale green precipitate is
'formed, consisting, according to Winckelblech (Ann. Ch. Pharm. xiii. 278), of oxalate of nickel and nickel-ammonium"
3. Source: "Zinc, cadmium, mercury, bismuth, tin, antimony, arsenic, nickel, cobalt ...", by Carl Schnabel, page 578, to quote:
"The powdered ore is heated with concentrated oxalic acid solution, which leaves nickel undissolved. The residue is reduced to obtain nickel." Link:
http://books.google.com/books?id=DWx9AAAAIAAJ&pg=PA578&a...
Suggested Synthesis to Extract Nickel
Dissolve the Cu/Ni in a concentrated Oxalic acid solution. Add a solution of Sodium oxalate to dissolve the Copper oxalate precipitate (I am expecting
this to work as a soluble Nickel-Sodium oxalate double salt apparently does not exist, so Nickel Oxalate will not dissolve). Separate the residue
which should be Nickel, some Nickel oxalate and other impurities for further thermal processing.
I am expecting the creation of some Nickel oxalate as apparently Oxalic acid with mild heating dissolves Nickel Oxide and also can cause the
dissolution of Nickel in Nickel ferrite. References:
"Dissolution of Nickel Oxide in Oxalic Acid Aqueous Solutions"
Per the abstract:
"The dissolution of nickel oxide (bunsenite) in acid solutions containing oxalic acid has been studied at 70.0°C. The dependencies of the rate of
dissolution on total oxalate concentration and on pH have been explained by assuming a mechanism involving the transfer of two different surface
complexes, I and II, that predominate in different pH ranges. The rate law is R=k1{I}+k2{II}, where {} denotes surface concentration. The values
k1Ns=3.04×10−3 mol Ni m−2 s−1, k2Ns=1.84×10−3 mol Ni m−2 s−1, together with the stability constants K1=675
mol−1 dm3 and K2=60 mol−1 dm3 fit all the results very well. The species formed in more acidic media is both more stable and more
reactive."
LINK:
http://www.sciencedirect.com/science/article/pii/S0021979701...
"Dissolution of Nickel Ferrite in Aqueous Solutions Containing Oxalic Acid and Ferrous Salts"
Per the abstract:
"The dissolution of nickel ferrite in oxalic acid and in ferrous oxalate-oxalic acid aqueous solution was studied. Nickel ferrite was synthesized by
thermal decomposition of a mixed tartrate; the particles were shown to be coated with a thin ferric oxide layer. Dissolution takes place in two
stages, the first one corresponding to the dissolution of the ferric oxide outer layer and the second one being the dissolution of
Ni(1.06)Fe(1.96)O(4). The kinetics of dissolution during this first stage is typical of ferric oxides: in oxalic acid, both a ligand-assisted and a
redox mechanism operates, whereas in the presence of ferrous ions, redox catalysis leads to a faster dissolution. The rate dependence on both oxalic
acid and on ferrous ion is described by the Langmuir-Hinshelwood equation; the best fitting corresponds to K(1)(ads)=25.6 mol(-1) dm(-3) and
k(1)(max)=9.17x10(-7) mol m(-2) s(-1) and K(2)(ads)=37.1x10(3) mol(-1) dm(-3) and k(2)(max)=62.3x10(-7) mol m(-2) s(-1), respectively. In the second
stage, Langmuir-Hinshelwood kinetics also describes the dissolution of iron and nickel from nickel ferrite, with K(1)(ads)=20.8 mol(-1) dm(3) and
K(2)(ads)=1.16x10(5) mol(-1) dm(3). For iron, k(1)(max)=1.02x10(-7) mol of Fe m(-2) s(-1) and k(2)(max)=2.38x10(-7) mol of Fe m(-2) s(-1); for nickel,
the rate constants k(1)(max) and k(2)(max) are 2.4 and 1.79 times smaller, respectively. The factor 1.79 agrees nicely with the stoichiometric ratio,
whereas the factor 2.4 implies the accumulation of some nickel in the residual particles. The rate of nickel dissolution in oxalic acid is higher than
that in bunsenite by a factor of 8, whereas hematite is more reactive by a factor of 9 (in the absence of Fe(II)) and 27 (in the presence of Fe (II)).
It may be concluded that oxalic acid operates to dissolve iron, and the ensuing disruption of the solid framework accelerates the release of nickel."
Link:
http://www.ncbi.nlm.nih.gov/pubmed/11254279
Note, as I am currently out of stock in H2C2O4, I am not immediately able to test this suggested synthesis.
[Edited on 16-5-2012 by AJKOER]
|
|
|
Formatik
National Hazard
Posts: 927
Registered: 25-3-2008
Member Is Offline
Mood: equilibrium
|
|
Another way to do separation is through aqueous hydrazine. Starting from an acid mixture of copper and nickel nitrate. Diluting this with some water
then adding portion-wise aqueous hydrazine. Hydrazine first attacks the Cu ion mainly reducing it to form copper. Then, about after that point it will
form a purple lilac explosive complex (nickel nitrate hydrazinate). This residue is hazardous and possibly unstable due to copper contamination. The
hydrazine addition should be done just up until that point at which the lilac salt starts appearing or just before this, otherwise it becomes
difficult to judge and the nickel will have been used up (a colorless solution obviously means the reaction has gone too far). The brown of the copper
makes this difficult to see. But the copper can easily be reduced out, then filter the solution and solutions can be obtained that do not have the
faintest scent of hydrazine, and are pure green color.
|
|
|
m1tanker78
National Hazard
Posts: 685
Registered: 5-1-2011
Member Is Offline
Mood: No Mood
|
|
Formatik, would you happen to have a reference for this?
Ajoker: I'm oxalic acid-poor as well. It'll have to wait.
Tank
Chemical CURIOSITY KILLED THE CATalyst.
|
|
|
Formatik
National Hazard
Posts: 927
Registered: 25-3-2008
Member Is Offline
Mood: equilibrium
|
|
No reference here this time, it's from my own experimentation. Because the amounts weren't noted, I had to kind of feel out the reaction a few times.
It is a quick route for impatient people.
I also doubt the suggested oxalic acid route above will work. The reference: Neueste Erfindungen und Erfahrungen auf den Gebieten der praktischen
Technik, Elektrotechnik, der Gewerbe, Industrie, Chemie, der Land und Hauswirthschaft (1895), Vol.21, p. 353 speaks of a solution of nickel
oxalate in potassium oxalate. So nickel oxalate will build double salts with alkalis.
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
| Quote: Originally posted by Formatik |
I also doubt the suggested oxalic acid route above will work. The reference: Neueste Erfindungen und Erfahrungen auf den Gebieten der praktischen
Technik, Elektrotechnik, der Gewerbe, Industrie, Chemie, der Land und Hauswirthschaft (1895), Vol.21, p. 353 speaks of a solution of nickel
oxalate in potassium oxalate. So nickel oxalate will build double salts with alkalis. |
Per my cited source ("A dictionary of chemistry and the allied branches of other sciences", Volume 4, by Henry Watts, page 262) to quote again:
"Oxalate Of Nickel, ... Greenish-white precipitate insoluble in water, soluble in ammonia and in ammoniacal salts. It dissolves also in potash,
forming a crystallisable potassio-nickel-oxalate. Neutral oxalate of ammonium dissolves oxalate of nickel, and the solution yields by evaporation
green prisms of ammonio-nickel-oxalate. " so Nickel oxalate apparently forms double salts with Potassium & Ammonia, which is in agreement with
your source on the Potassium double salt.
Interestingly, a Sodium-Nickel double salt is not referenced nor does a general internet search find any such salt. However, there is an instance of
Nickel plating in a Sodium oxalate / Oxalate acid medium (see http://www.tandfonline.com/doi/abs/10.1080/03602550500373881... ). If a double salt formation was possible, wouldn't it's formation interfere with
the Ni plating process?
So my doubts are leaning more towards the fact that the Sodium double salt happens not to exist.
[Edited on 29-5-2012 by AJKOER]
|
|
|
Formatik
National Hazard
Posts: 927
Registered: 25-3-2008
Member Is Offline
Mood: equilibrium
|
|
No, there's a reference which confirms it won't work since there is such a complex salt. Analytical Chemistry (1921), Vol. 1, p. 385 by F.P.
Treadwell states nickel oxalate forms a soluble sodium nickel oxalate complex salt with sodium oxalate. The reference from Watts which was
Winckelblech, never attempted to determine any interaction with sodium salts, so that's why it was never noted. It was not comprehensive.
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
Formatik:
You are correct. Found it in an online Googlebook. Source: "Analytical chemistry", Volume 1, by Frederick Pearson Treadwell, page 385 (as you noted).
Link:
http://books.google.com/books?id=ubuaqKkwsQoC&pg=PA385&a...
----------------------------------------------------------
The same source also provided me with another route to investigate. First, here is the authors comment on Copper related reactions on page 277:
"The proper solvent for copper is nitric acid:
3 Cu + 8 HNOs -> 3 Cu++ + 6 N03- + 4 H20 + 2 NO (g)
Bright copper is not dissolved by hydrochloric acid alone, but in the presence of a weak oxidizing agent, e.g., ferric chloride, the solution of the
metal is easily effected. Hot hydrobromic acid dissolves it with evolution of hydrogen, forming cuprous hydrobromic acid: 2 Cu + 6 HBr <=>
H4[Cu2Br6] + H2 (g)
At the beginning of this reaction the solution usually turns dark violet on account of the formation of the cupric salt of cuprous hydrobromic acid,
owing to the copper being somewhat oxidized on the surface. In this case, however, the solution soon becomes colorless, owing to the reduction of the
cupric salt by metallic copper. On adding water to the clear solution cuprous bromide is precipitated: [Cu2Br6]-- ==> Cu2Br2 + 4 Br-.
Copper is not attacked by dilute sulfuric acid, but it dissolves in hot concentrated sulfuric acid, forming cupric sulfate with evolution of sulfur
dioxide:
Cu + 2 H2S04 -> CuS04 + 2 H20 + S02 (g)
The behavior of copper toward acids can be understood by reference to the electromotive series (p. 41). As copper is below hydrogen in the series it
can be oxidized by hydrogen ions only when the concentration of cupric ions is kept very low (cf. p. 43). Hydrobromic acid dissolves copper because a
slightly ionized complex is formed. Sulfuric acid dissolves copper by virtue of the oxidizing power of the hexavalent sulfur."
So, hot HBr dissolves Cu owing to the formation of a slightly ionized complex. Interestingly, per another source with respect to Nickel, to quote:
"Due to the corrosive nature of halogenated gases, special attention should be paid to their handling. High purity WF6 and HBr are supplied typically
in stainless steel, pure nickel or nickel lined cylinders. Passive films formed on the two surfaces are different."
Link: http://www.smlassociates.com/cormetal.shtml
So HBr appears to be unreactive with respect to Nickel owing to the formation of a passive film.
So assuming the existence of the Cu/Ni in the form of an alloy does not presence any issues (?), an idea that might be interesting to investigate is:
> Place the Cu/Ni coin (or first reduce the coin to a fine powder) in hot HBr and see if the Ni remains largely undissolved. I would also collect
the H2 gas, as per the reaction:
2 Cu + 6 HBr <=> H4[Cu2Br6] + H2 (g)
for each mole of Hydrogen, two moles of Copper have been consumed. Knowing the observed volumes of H2 produced and starting weight of the coin and
the residual undissolved mass, one can check on the progress of the reaction. Also, from the filtered solution upon adding water, per the author's
cited reaction:
[Cu2Br6]-- ==> Cu2Br2 (s)+ 4 Br-
the weight of the cuprous bromide precipitated also speaks to the amount of Copper consumed.
[Edited on 30-5-2012 by AJKOER]
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
Here is an extract from an old recipe from "Hand-book of chemistry", Volume 5, page 357 by Leopold Gmelin that mentions some of the steps noted above
to separate out Nickel:
"5. Proust heats roasted copper-nickel with dilute sulphuric acid, and adds carbonate of potash to the filtered solution, to precipitate arseniate of
ferric oxide, till iron can no longer be detected in the liquid by ferrocyanide of potassium. The liquid is then filtered again, and sulphuretted
hydrogen passed through it, to precipitate arsenic, copper, and bismuth, till it is so far saturated as to retain the odour of the gas after being
kept for 24 hours in a closed vessel. The liquid, once more filtered and then evaporated, yields crystals chiefly consisting of sulphate of
nickel-oxide and potash, while the cobalt-salt, for the most part, remains in solution; the former is repeatedly dissolved and recrystallized to free
it from adhering cobalt-salt, and its solution afterwards treated with carbonate of potash, which precipitates the nickel in the form of carbonate."
Also, per page 355:
"The copper-nickel or the cobalt-speiss is generally roasted in a state of powder (at a gentle heat at first, to prevent it from baking together),
whereby the greater part of the arsenic is removed, the nickel oxidated, and a saving of nitric acid thus effected. Since, however, the roasting
process leaves a portion of the arsenic combined in the form of arsenic acid with the oxide of nickel, the roasted ore must be several times
intimately mixed with charcoal dust and again roasted, as long as vapours of arsenic continue to be evolved. Erdmann moistens the roasted cobalt
speiss with water and places it in a cellar till it is converted into hydrate; it is thereby rendered more easily soluble.
1. Laugier dissolves the roasted copper-nickel or the speiss in nitric acid, passes sulphuretted hydrogen through the dilute acid solution till all
the arsenic, copper, bismuth, and antimony are precipitated—then falters
—precipitates all the iron, cobalt, and nickel with carbonate of soda— washes the precipitate thoroughly, and treats it first with oxalic acid and
then with ammonia, as described on page 319, repeating the solution of the nickel-oxalate in aqueous ammonia, till the liquid which stands above the
resulting precipitate no longer exhibits a rose-colour, and is almost wholly free from cobalt."
Link: http://books.google.com/books?pg=PA357&lpg=PA32&dq=h...
[Edited on 12-8-2012 by AJKOER]
|
|
|
cyanureeves
National Hazard
Posts: 737
Registered: 29-8-2010
Location: Mars
Member Is Offline
Mood: No Mood
|
|
this is very similar to what has been done except potassium carbonate is here used and it was also mentioned by someone here that maybe sodium
metabisullfate could be used to drop nickel. sulphuretted hydrogen and sodium metabisulfate both can be used to drop gold from auric chloride i think.
the one thing that is very different here is the roasting in sulfuric acid, dissolving in sulfuric acid has been done but not roasting. its
interesting how in this case hydrogen sulfate is used to drop the copper first and not the nickel.
|
|
|
m1tanker78
National Hazard
Posts: 685
Registered: 5-1-2011
Member Is Offline
Mood: No Mood
|
|
Would somebody please demonstrate separation by bubbling H2S gas through a solution of the combined metals. Seems easy enough for the chemist who has
the reagents, equipment, knowledge and the desire.
Tank
Chemical CURIOSITY KILLED THE CATalyst.
|
|
|
tetrahedron
Hazard to Others
Posts: 210
Registered: 28-9-2012
Member Is Offline
Mood: No Mood
|
|
@tank: this thread's OP is looking to obtain pure nickel salts, but you admit it's hard to achieve:
| Quote: Originally posted by m1tanker78 | | An old plater once told me that the last 1-2% is a PITA. I tend to believe him even though he didn't have any math or graphs to back up his life work.
LOL |
i attempted electrorefining Cu/Ni in a similar way to yours. using a CuSO4 electrolyte, for the anode i soldered one side of the coins to
the tip of some insulated Cu wire individually, then i spray painted the whole side, giving me about 15cm3 total surface area on the front,
for 15g Cu/Ni (3:1). the cathode was a Cu wire. i don't remember the current but it was pretty low (around 0.5A). i stopped the experiment (total
duration <24h) when i noticed the electrolyte seeping under the clear paint. the electrodeposited Cu looked the same as if a pure Cu anode had been
used.
| Quote: Originally posted by m1tanker78 | | Quote: Originally posted by blogfast25 | The metal powders are too intertwined and there may have been some coprecipitation too, making good separation difficult. It’s a recipe for very
impure Ni and very impure Cu, both of which would then need refining anyway.
|
That's precisely what I observed. I was trying to get around having to scrape the Copper(I) Oxide (and/or hydride?) off the cathode every hour or so.
Aside from that, the electrolytic separation of the metals seems to work well. I'm working on starting another batch as I type...
|
have you tried boiling these mixed powders in the electrolyte to reduce the last bits of copper in solution by displacing it with nickel, and and get
pure NiCl2? the Cu metal already precipitated should play no role. i had no mixed powders, so i used some fresh coins for the metathesis
(this has severe limitations due to the small surface area). soon they turned a rusty color. the powder method should be pretty efficient. in
industry, copper impurities in nickel electrorefining are removed with (pure) nickel powder as well.
|
|
|
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
| Quote: Originally posted by m1tanker78 | Would somebody please demonstrate separation by bubbling H2S gas through a solution of the combined metals. Seems easy enough for the chemist who has
the reagents, equipment, knowledge and the desire.
Tank |
As this is a bit of a ‘classic’ of older, wet analytical chemistry, try and consult older texts on this, for full details.
In practice I think you’d have to buffer the solution to a fairly low pH (say 4), then saturate it with H2S gas: CuS precipitates, Ni2+ stays in
solution. Then flush out most of the H2S with air and separate the metals by filtration.
Not difficult in principle but it will require scrubbing out the highly toxic H2S very effectively and decent filtration media. Commercial bleach is a
decent H2S scrubber because it oxidises the H2S to elemental sulphur and water. But use at least two wash bottles.
|
|
|
cyanureeves
National Hazard
Posts: 737
Registered: 29-8-2010
Location: Mars
Member Is Offline
Mood: No Mood
|
|
is this H2S the same gas that gas line workers watch out for with beeper detectors? isnt it as poisonous as cyanide? any how does this H2S work only
on cupronickel in HCl or would it work with cupronickel in sulfuric acid? how is this poison gas made in the lab?
|
|
|
| Pages:
1
..
5
6
7
8
9
..
11 |
![[*]](images/xpblue/default_icon.gif){kind=icon}
