Sciencemadness Discussion Board - Separation of a US nickel

Can you get 500 mg of nickel chloride for $0.50 or a deal similar to that at your ceramics supplier? When one conducts microchemistry on a limited budget, purchases are normally out of the picture. Nickels are not.

Lanth, I think you're counting your chicks before they've hatched. I'm not downplaying the ascorbic acid method to isolate Ni but you haven't provided a yield figure - if you've tried it at all yet.

@Sedit: How does NH3 stack up in sulfate liquor? I have some concentrated Cu/Ni sulfate but ammonium sulfate is much more soluble than NH4Cl so I doubt this would work well. Can you post a link to the topic you're referencing? TFSE sent me on a wild goose chase.

Thanks,
Tank

Random - 27-6-2011 at 14:22

I took some old cupronickel coins and put them into HCl(aq) and then added 1,5mL of ammonium nitrate solution. I will wait until they dissolve and then I will try to separate copper from nickel using potassium metabisulfite solution. Will report the results.

LanthanumK - 27-6-2011 at 14:24

Will the excess nitrate and metabisulfite react, consuming the latter reactant?

LanthanumK - 27-6-2011 at 14:27

The ascorbic acid method is very time consuming and possibly inefficient, as are all of my methods, but does not require any expensive chemicals like nitric acid. Is doing that known as "ghetto chemistry"?

[Edited on 27-6-2011 by LanthanumK]

cyanureeves - 27-6-2011 at 16:04

lanthanumK. when i wrote using nickel as electrode i meant pure nickel metal as anode to plate out the nickel only out of solution. i have tried to plate the copper out of the solution using copper metal as anode also and even though much of the copper plates out to the cathode much will be left in the solution. ghetto chemistry or not ,as long as it works as good as the Hamptons chemistry.

[Edited on 28-6-2011 by cyanureeves]

m1tanker78 - 27-6-2011 at 16:21

...Is doing that known as "ghetto chemistry"?
[Edited on 27-6-2011 by LanthanumK]

It depends on who you ask. I call it resourcefulness. Other here will think it's tacky. Who cares!?

What are you using for a power supply? You might be able to get away with stacking the coins higher than the electrolyte level and devising a simple gravity contact (for the anode). Personally, I'd do that for the cathode as well. The idea is to keep Zn and Fe (from alligator clips or w/e) out of the solution.

Tank

Random - 27-6-2011 at 16:27

Will the excess nitrate and metabisulfite react, consuming the latter reactant?

Maybe, if solution is enough acidic there should be HNO3 that will oxidize SO2 to SO3 (in solution to H2SO4). But I will neutralize the solution with carbonate until almost neutral to not waste too much metabisulfite that could escape into air as SO2. By the way, I will dilute the solution because CuCl is soluble in more concentrated HCl(aq) (also, concetrated here doesn't mean 30%, i seen 19% HCl dissolving CuCl)

cyanureeves - 27-6-2011 at 16:34

i dont know if its too late but the other thread is ( how extrac nickel ) by plante1999

[Edited on 28-6-2011 by cyanureeves]

LanthanumK - 28-6-2011 at 03:40

If I have a pure nickel electrode already, I don't think I would bother doing all this work to extract a few milligrams of nickel from a coin. <s>Tank, I have a video of my typical electrolysis procedure; if I can post it I will</s>I have it here: http://www.youtube.com/watch?v=Q5zEaX5ULOw. You can see that the alligator clips do not touch the electrolyte, although they are heavily corroded from spattering. The container is a medicine cup.

[Edited on 28-6-2011 by LanthanumK]

[Edited on 28-6-2011 by LanthanumK]

problem solved

cyanureeves - 28-6-2011 at 04:26

Quote: Originally posted by The_Davster

Drive to Canada. Go to a bank. Ask for a brick of nickels (100$ worth) Those from 1955 to 1981 are pure nickel metal.

Naturally, only f you are after the metal, and not just looking for some chemistry to do...

davster you must know E.S.P.

Sedit - 8-7-2011 at 08:47

No I can not, if I did I would find a new supplier because that price is outrageous. I can order almost any amount I want since she breaks it down out of 50lb bags for me. Nickle oxides being such powerful coloring agents they are its not unheard of for someone to request only a few grams at a time. I on the other hand by quarter pound bags when I need it.

""
@Sedit: How does NH3 stack up in sulfate liquor? I have some concentrated Cu/Ni sulfate but ammonium sulfate is much more soluble than NH4Cl so I doubt this would work well. Can you post a link to the topic you're referencing? TFSE sent me on a wild goose chase."""

I think its in a thread titled separation of Cu++ and Ni++ ions or some shit like that. I believe there are still pictures of the green precipitate there that I acquired. If you still can't find it ill have a look around. I wish I could remember all the details but all my notes where on a computer thats crashed and my memory is as good as a gold fish.

LanthanumK - 8-7-2011 at 10:27

Of course, you would not buy 100 500mg packets; the price would be outrageous then.

If I use 50mg per experiment, 500g (at a discount price) will be enough to conduct 10,000 experiments; I may do 10. That is a waste of chemicals, storage space, and a toxicity hazard if I ever want to dispose of it. I agree with what woelen says here on his site: http://woelen.homescience.net/science/chem/misc/chemicals.ht... except I take it farther. The 1 oz of NaBr I purchased was way too much; I only didn't buy a smaller quantity because none were available.

m1tanker78 - 21-7-2011 at 12:36

I'm dissolving (electrolytically) a few US nickels as we speak. I'm also simultaneously dummy plating out the copper, with the understanding that it won't completely plate out. The idea is to reduce the copper content AMAP to ease subsequent purification somewhat. I'm using HCl so I can give the ammonia treatment mentioned above a try as well. Keeping with the spirit of reducing waste, I'm collecting the excess spongy copper to try to make this experiment as 'cost-effective' as possible. I don't think it's a good substitute for buying pure job-specific salts but the results will certainly be filed away with my notes, regardless of the outcome.

Tank

Melgar - 21-7-2011 at 13:54

If pre-1981 canadian nickels are too hard to find, pre-1999 canadian quarters and dimes are also pure nickel. If you've got a few bucks handy, go onto ebay and you'll find people selling loads of pre-1981 canadian nickels for the nickel content. I bought around 30 nickels for $10 once. Sure, it'd be cheaper to find them myself, but that's still a great price for pure nickel metal.

cyanureeves - 21-7-2011 at 16:07

tank that is the idea. nickel can be gotten as coins or salt but if i could avoid buying it would be great because if you look around there is stainless and cupronickel everywhere.to me nickel in any form is good because even if you could afford to buy nickel salts or metal for plating it will require alot to plate something bigger than a rim or even a hubcap.i would fill a trash can full of nickel plating solution if i could grab it from pesos or nickels.copper is beautiful but in this case i wish it were'nt ever discovered. patience and determination to you.

m1tanker78 - 22-7-2011 at 12:29

Reeves, you'd have one hell of a task disposing of a trash can full of spent or contaminated nickel solution. Either way, dissolving nickels by electrolysis coupled with Sedit's ammonia separation method makes small scale nickel salt procurement very feasible, I'd say. It's unlawful to use currency for this purpose (as we all probably know). Poor Thomas Jefferson!...

Having said that, I have a ton of Mexican pesos put away somewhere that I never got around to exchanging back to dollars. Care to enlighten me some?

Tank

cyanureeves - 22-7-2011 at 15:30

no tank it was not conclusive yet. sedit separated nickel and copper in their hydroxides by adding ammonium hydroxide as the copper hydroxide is supposed to be insoluble and nickel complexes going into solution with ammonium hydroxide and then precipitating. i was to able to nickel plate a pair of pliers with the solution. i turned it back to the carbonate then eventually to the acetate with vinegar but even then copper plated out as well .sedit went on further to say his blue nickel precipitate did not produce a blue solution any further upon adding ammonium hydroxide again. mine did! although it was much lighter, others said it was not possible as a nickel ion is a nickel ion is a nickel ion and the complex will always turn blue with ammonia. geesh i got to get a life! pesos have very little nickel but they are plentiful and when you dissolve them using current they turn bright red then alternating to fluorescent yellow- green. who needs stinkin nickel when you have pretty colors amigo?

m1tanker78 - 23-7-2011 at 07:59

Speaking of colors: Some of my samples have produced colorful layer effects. A good portion of the copper was reduced out so the liquor was nickel-heavy to begin with. It took A LOT of household ammonia to treat a small volume of liquor so I opted for concentrated ammonia.

I momentarily bumped the voltage up twice toward the end of the run and a nice nickel foil plated out amid the copper sludge (at the cathode). Most of the copper that I periodically removed from the cathode was in the form of Cu2O, Copper(I) Oxide. Because this form of copper is spongy, I'll have to wring it out in a water wash in order to remove much of the nickel solution trapped within it. This conforms to my other goal here which is to collect the copper in a usable form.

I hope to have some time to do a few more experiments when I'm back home later today.

Tank

cyanureeves - 23-7-2011 at 08:44

oh man only once have i produced a nickel foil with the guts of a nickel cadium batttery but i didnt record what all i did so i couldnt reproduced it again. what are you using as solution? hcl acid or sulfuric? and at what voltage do you recall you were using when you got the foil? pesos ( Contenido: 65% de cobre, 10% de níquel y 25% de zinc) 65% copper,10% nickel 25%zinc. but looks as if they add paprika, lemon and cilantro with all the great colors they produce.

Sedit - 23-7-2011 at 08:45

no tank it was not conclusive yet. sedit separated nickel and copper in their hydroxides by adding ammonium hydroxide as the copper hydroxide is supposed to be insoluble and nickel complexes going into solution with ammonium hydroxide and then precipitating.

This is not how it works. Honestly I'm not sure what exactly is taking place. When I started it seemed straight forward and proceeded as planned yet when I brought the experiment to the attention of people here they stated that it wouldn't work

. This is one case where theory does not exactly match experiment.

The idea came from a patent that I posted in another thread that spoke about precipitating Ni oxide/hydroxides onto activated carbon to make a highly active Nickle reducing agent. They used Ammonium hydroxide to precipitate the nickle and that's when I decided that it should work well for separating Cu and Ni since Copper complexes with NH3 and I was unaware that Ni also did.

Experiments showed that Nickle can in no way complex half as much as Copper does and I believe this is why the nickle precipitates out instead of complexing with the Ammonium ion.

The only thing that did not make sense was the fact that after washing the precipitate repeatedly and adding more ammonia there was no sign of any complex yet the precipitate tested positive for Ni. I added HCl to it and it crystallized into the distinct green Nickle chloride hydrate.

The process is in no way stream lined however the principle of it does work. I wish I was 100% sure how it worked but it does indeed precipitate the Nickle while holding the Copper in as its soluble amine complex. I have never obtained yeilds or much other data I wished to have so if someone wishes to further the experiment please be a doll and see what you can gather. Its a huge pain in the ass, the main issue is the fineness of the precipitate. Its next to impossible to filter out.

I also one time much later then those experiments managed to dissolve all of the precipitated into excess NH3. I left this solution sit out for god knows how long and in the bottom large green crystals precipitated out. I still have them crystals but I don't know what they are. I would post a picture but it will be very large and I don't know to re-size it to fit into the window here.

cyanureeves - 23-7-2011 at 10:45

i have it all wrong and i'm sorry for that.

m1tanker78 - 23-7-2011 at 13:43

[...] but looks as if they add paprika, lemon and cilantro with all the great colors they produce.

LMAO! Yeah, I think I'll leave the pesos alone. IIRC, I bumped up to 1.7V (drawing ~ 3A). I used a relatively low concentration of HCl.

Sedit, oddly enough, it appears that the blue liquid layer contains a much higher proportion of nickel than the precipitate. I can't confirm that just yet, though. I took a small sample of both and subjected to heat. Both release a fine white smoke when dry. My first thought was ammonium chloride. I 'condensed' some of the smoke on a cool surface. This left a fine bright yellow powder when heating the blue complex. I neglected to collect the white smoke from the other stuff that precipitates. I performed these experiments outside. I'm not comfortable taking chances with heating unknown nickel salts.

Next up, I would like to reacidify a sample of both and try converting them to carbonate. The color of the precipitate might yield some clues about the content of each. Ammonium chloride should easily remain dissolved and not interfere with the experiment in any way.

Tank

Sorry for double posting...

m1tanker78 - 23-7-2011 at 15:12

I went ahead with the carbonate prep I described above. I used HCl and a strong sodium carbonate solution. The fluff didn't fully redissolve when acidified. It left a murky solution (but no longer brown). I think it was slightly OD green. When the solution returned to basic, the brown fluff reappeared. No sign of copper or nickel but it was hard to tell. BTW, this precipitate was washed meticulously prior...

On the other hand, the blue complex turned a nice green when acidified. I didn't observe any insoluble crap at this point. When basified, the solution turned sky blue and a white precipitate began to fall out. As the precipitate layer grew in thickness, the light lime-green color became more and more obvious. What bothers me is why the solution remains colored baby blue.

[EDIT]: The blue is probably dissolved Ni2+?

Regardless, here's a pic I snapped. I didn't realize it turned out blurry until after I uploaded it...

Tank

[Edited on 7-23-2011 by m1tanker78]

[Edited on 7-23-2011 by m1tanker78]

cyanureeves - 23-7-2011 at 17:40

is the green copper? the green looks like nickel in the carbonate and in the chloride and in the sulfate but copper in the chloride is green also. isnt copper known to look green or black in concentration then turn blue as it gets more diluted?hey maybe copper and nickel swapped around and the more you dilute it the more copper disperses and nickel sinks. i know i'm dreaming but nickel does get pulled down to the bottom in darker concentration in some solutions.

Arthur Dent - 24-7-2011 at 06:07

I prepared some nickel chloride a while ago by dissolving 6 1970's canadian nickels in HCl. The resulting solution was a very deep emerald green, and I stored the concentrated solution in small 125 ml pyrex bottle. I never attempted to reduce the solution by evaporation or boiling off to make crystals because frankly, nickel salts and their vapors scare me!

Is my little bottle of NiCl2 okay for storage in my lab or should i seal it ever further? It's just about 75ml but I treat it as if it were metallic mercury... very carefully!

But definitely, trying to separate nickel from other metals might be an interesting acadmic endeavor, but I'd rather start with pure nickel to start with if I wanted to prepare a nickel salt... you're sure of the end result then. Perhaps we could try to find potential sources for pure nickel... as mentioned above, a lot of canadian coinage was made out of pure nickel, the coating of many rare earth magnets too...

Robert

m1tanker78 - 24-7-2011 at 12:00

I'm open to suggestions for a more conclusive test. Until I can perform such test(s), here's the way I see it. The precipitate is Nickel(II) Carbonate. The solution contains a small amount of Nickel(II) ions aside from Na, NH4, etc. A recap of how I got to this point:

Dissolved some US nickels (cupronickel) and simultaneously dummy plated out some of the copper.

Treated sample of dark emerald green liquor with concentrated ammonia.

Separated blue complex from brown/green precipitate.

Added concentrated HCl to blue stuff (turned emerald green).

Added saturated solution of sodium carbonate until basic. Solution turned sky blue color and white precipitate formed. Precipitate appears light green (accentuated in above pic). Solution remains smurf blue.

++++++++++++++++++

On the other hand, the precipitate from the ammonia treatment appears to be a mixture of copper and iron or some other component. I used distilled water at every stage (as needed) but my HCl is tech grade.

I suppose I'll wash the greenish ppt a few times and go from there?...

Tank

cyanureeves - 24-7-2011 at 12:20

tank on the second part of your answer are you saying that the brown/green precipitated from ammonia was iron and copper? then the blue liquid had all your nickel which turned green when hcl acid was added and precipitated with sodium carbonate as nickel carbonate? why the sad face icon? who gives a rat blankety about iron and copper?boy i hope you got nickel carbonate. if it is then drying it and adding a bit of vinegar and warming the solution will make nickel acetate and should be enough to nickel plate a penny with about 4 or 5 volts. if it plates pink then you know you have copper in your precipitate.

m1tanker78 - 24-7-2011 at 14:17

Well, It's hard to tell what's in the precipitated crap from the ammonia treatment. After several washes, the ppt remains brown. It settles rather neatly and leaves perfectly clear water at the top of the 'test tube'. I was hoping to be able to isolate or at least see some conclusive indications of copper in this stuff.

On the plus side, I'm almost certain that nickel is absent from the ppt (after a couple of washes, anyway).

Now, washing the precipitated carbonate still yields a green-tinted white ppt and a light baby blue supernatant. :cool:

The reason I'm banking on fairly pure nickel(II) carbonate is...

*Apparently, nickel is absent in the ammonia-precipitated crud. By this, I mean after a couple of washes. The as-precipitated crap is too murky to draw any conclusions (but gives an indication of copper).

*Nickel(II) carbonate is verrrry slightly soluble in water so in theory, I'll never have a perfectly colorless solution no matter how many times I wash the carbonate.

*The pure carbonates of nickel and copper can be distinguished somewhat by color. I prepared a small batch of copper carbonate so that I could compare side-by-side.

*Pure copper carbonate will leave a perfectly colorless supernatant, unlike nickel carbonate.

I don't want to jump to conclusions so I'll take this a step or two further. I might even try plating a small part like cyanureeves suggested. I'd be inclined to plate from a sulfate solution, though. Why acetate (for nickel)??

Tank

cyanureeves - 24-7-2011 at 16:59

geesh man i just got too excited, there is no better reason for the acetate over the sulfate as a test. just that i have pure nickel sulfate and as a matter of fact its even easier to plate with nickel sulfate but nickel acetate can readily plate in the cold. you dont even have to warm nickel acetate like i said before but you do nickel sulfate to adhere better. i saw a guy on you tube plate out of a cold acetate solution and i also plated a pair of kleins like that. carry on, carry on youre doing great. i didnt know about copper carbonate leaving a clear solution i thought only the hydroxides were hard to dilute.

Sedit - 25-7-2011 at 07:27

Im a bit confused as to whats going on here mainly due to the talk of the carbonate,

Do I understand correctly that your saying the green precipitate formed from treating the dissolved coin salts in (aq)NH3 tested negative for Nickle?

m1tanker78 - 25-7-2011 at 08:22

Do I understand correctly that your saying the green precipitate formed from treating the dissolved coin salts in (aq)NH3 tested negative for Nickle?

I haven't conducted any conclusive tests yet. The brown/green precipitate from ammonia treatment gives a very strong indication of iron(II). The only possible sources of iron contamination were my tech grade HCl and/or iron leaching from the SS cathode. I took care to remove and replace the cathode live to prevent corrosion but the potential was so low that a little undoubtedly leached and dissolved.

Although I plated out a lot of the copper, there MUST be a measurable amount left in the liquor. The cell liquor started out emerald blue and slowly transformed to emerald green as the nickel concentration rose and copper concentration fell.

The most troublesome part of my series of experiments on this is not being able to find the copper! It may well be hidden among the precipitated nickel carbonate. The iron contamination is also a PITA because it screws the simple tests I've run so far with the ammonia-precipitated stuff.

I may need to repeat this experiment from the very beginning using an insoluble (filter-able) cathode due to the low cell voltage and hence, lack of cathodic protection. Graphite should work just fine for this, I think.

Tank

Sedit - 25-7-2011 at 15:28

Convert it to there respective sulfates and precipitate the Iron by bubbling air in, use Pb next time. Filter this and it should remove most of the Fe contamination.

The precipitate in my test from the Ammonium hydroxide treatment is what was washed repeatedly with water then converted to the chloride yielding green needle crystals presumed to be pretty pure Nickle chloride. HCl and aluminum foil precipitated a black substance that was strongly attracted to a magnet.

cyanureeves - 25-7-2011 at 16:40

both of you gentlemen are at exactly opposite poles of each other unless tank isolated the nickel from the very start with ammonium hydroxide.

m1tanker78 - 26-7-2011 at 07:09

Sedit, the magnet test (which I forgot about) will be useful once I remove iron from the mix. Thanks for the suggestions. Just to clarify, you precipitated nickel powder from an acidic solution of nickel chloride with aluminum, right? Seems easy enough for a qualitative test for nickel...

+++++++++++++++++++++++

I dried a sample of the freshly prepared carbonate (same as the pic I posted before). For clarity, this carbonate is prepared from the ammonia complex, not from the original precipitate.

Although the image is crap, the powder bears a striking resemblance to 'the real deal', nickel(II) carbonate. Then, I'm reminded, "Don't judge a chemical by its color."

So what's better here, acetate for a plating test or chloride then reduce with Al for a magnet test? It's probably too little to practically plate anything except maybe the tip of a copper wire or similar. I have ~ 500mL of cell liquor left from this batch.

Tank

cyanureeves - 26-7-2011 at 15:02

drop the metal with aluminum because whatever it is you can always redissolve it. go for the magnet because really my main interest is all about plating,besides you still have the brown precipitate. is that a bottle cap? its an awful tiny bit but another reason not to plate is because if it is copper then you will introduce nickel into your test and contaminate it with the anode. good work either way. you said you used household ammonia didnt you?was that the regular clear stuff? i am out of ammonium hydroxide and dont know if it is safe to use an aluminum jogger's water bottle to heat ammonium nitrate and sodium bicarbonate to generate ammonia.

[Edited on 26-7-2011 by cyanureeves]

[Edited on 26-7-2011 by cyanureeves]

m1tanker78 - 27-7-2011 at 07:27

Reeves, it's a large pill bottle cap. I poured a little out to show the color. The yield was just over a gram from the sample I took from the cell; I hesitated to do the entire batch since I'm still not sure what's what. If this stuff is relatively pure nickel carbonate then I'll drop the rest of the nickel out of the remaining liquor.

I used concentrated ammonia. 'Clear' ammonia seems to work but you'll need a s*it load of it. The extra water isn't a problem if you're collecting the precipitate. If you're going for the ammonia complex, you'll probably lose some nickel to the extra water. That's if my assessments so far are correct.

If you have or can get ahold of some ammonium sulfate and calcium hydroxide (hydrated lime), that's another convenient route to ammonia. Even a crude distillation setup will yield concentrated ammonia of high purity from that mix. The byproduct is calcium sulfate (plaster).

Quote:

i am out of ammonium hydroxide and dont know if it is safe to use an aluminum jogger's water bottle to heat ammonium nitrate and sodium bicarbonate to generate ammonia.

Aqueous ammonia will corrode aluminum so, no.

m1tanker78 - 28-7-2011 at 17:01

I plated a piece of copper pipe today. The idea was to determine whether nickel ions are present in the ammonia complex that forms when a mixture of copper(II) chloride and nickel(II) chloride is treated with aqueous ammonia. For this, I had to prepare a new batch since I plated out much of the copper from my previous batch in an attempt to pre-refine the liquor. The problem with iron contamination was eliminated in this new batch. This experiment can't quantify nor rule out the presence of copper ions in the complex.

Contrary to what has been proposed and reported before, the experiment clearly demonstrates that nickel ions are present, possibly in abundance, in the ammonia complex. The proportion of copper ions (if any) of the same is unclear as of yet.

Again, forgive the poor quality photo. It's the cell camera, not me (I swear).

The overall nickel concentration per volume of electrolyte is very low so it's difficult to produce an even deposit, let alone a lustrous one.

I'm still open to suggestions for a relatively garage-friendly way to quantify the copper proportion in the complex.

Tank

cyanureeves - 28-7-2011 at 18:41

wonderful. nickel plating is not like silver plating in which silver is atractted to the cathode at all areas. nickel is not as simple and an even deposition on a round object is harder than laying it on a flat surface. even if you plate a flat surface it will depend on which side is facing the anode thats why multiple anodes are used. dont worry about the brightness because usually it takes a bit of nickel chloride,nickel sulfate and even epsom salt and boric acid to make the famous watts solution.that is just great.

m1tanker78 - 29-7-2011 at 18:08

Updates: The ammonia complex contains copper ions so this method isn't useful for separating copper and nickel. It does appear to be useful for removing other contaminants such as iron, though (see below). Lowering the current density during my plating tests produced a distinct copper deposit on top of the nickel deposit pictured above. Bummer...

On the flip side, I collected a carefully washed sample of the precipitate that's produced from adding ammonia to the chlorides. I non-destructively dried it by passing a gentle stream of air over it for 2 days or so. I'd previously pyro'd a sample of both products which was what aroused my suspicions of the ammonia technique in the first place.

It's hard to fathom how a tiny pellet (see bottom of shot glass) can make such a voluminous precipitate. The shot glass was literally filled to the brim with the decanted slurry...

What impressed me even more was that the pellet is rigid and retained its shape as it shrunk down. Weird...

I always thought the inside bottom of the shot glasses were perfectly rounded. Obviously, they ain't!...

The pellet is only barely attracted to a strong magnet. The attraction is almost negligible but then, I don't expect there to be much nascent metal or magnetic oxide(s) within.

It's possible that some of the nickel and copper precipitates as *ides along with [God knows what else]. These two metals are a modern-day Romeo and Juliet (so far). **Sings "Love Hurts"**

Now I have that damned song stuck in my head.

Tank

cyanureeves - 29-7-2011 at 19:44

i tried the salicylic method also and i probably have acetysalicylic because i could not produce copper nor nickel salicylic.salisylic acid supposedly does not combine with nickel carbonate but will with copper carbonate to make copper salicylic. i turned all the nickel/copper chloride solution to the carbonate and added salicylic acid but it all turned baby blue colored and i never got the bright blue insoluble copper salicylic.

m1tanker78 - 31-7-2011 at 07:47

Reeves,you went straight from electrolyte liquor to the carbonates without an intermediate ammonia step? I'm trying to picture your procedure.

I'm very tempted to try a few things in the molten state (outside!). That's a little far-reaching for academic purposes, though.

Tank

[Edited on 7-31-2011 by m1tanker78]

cyanureeves - 31-7-2011 at 10:46

yes tank i just dumped a whole lot of sodium carbonate in the chloric mixture but i think baking soda would've worked better. i dried the carbonates and went on to do the procedure that i saw on youtube to make coppersalicylate by aunonomus. he calls it copper aspirinate. first i made salicylic acid using myfanwy's method with dollar general aspirins which by the way dissolved cleanly as non of it was caught in the filter.myfanwy probably learned how to make salicylic acid here because there is a thread about it here and about copper salicylic,i just used his method because he laid it out on video. i need to make pure copper sulfate before i attempt to make copper salicylic again to see if my salicylic acid is good because my copper sulfate was gotten from epsom salt and copper and probably has alot of magnesium sulfate.

[Edited on 31-7-2011 by cyanureeves]

Mixell - 31-7-2011 at 11:08

You can use dimethylglyoxime to separate the nickel from the copper (Ni(dmgH)2 precipitates), although its not very affordable, may be only on small/demonstration scale.

UPDATE- Just read on the German wiki that it forms a complex with copper too, but my question below still stands.

I'm planning to form this complex myself, as I understand, it is used as a solution in ethanol, can acetone replace it?
Also, do I need any special conditions to dissociate the dimethylglyoxime into dmgH- and H+ like a basic environment? Or a dilute solution at neutral pH will also give good yield?

[Edited on 31-7-2011 by Mixell]

m1tanker78 - 31-7-2011 at 19:06

Reeves, your description sounds analogous to the ascorbic acid method that was proposed (but never attempted AFAIK). I'll look for solubility data and see if it's viable for this purpose.

Mixell, I never used DMG. Like you said, it reacts with both metals so it would be back to square one.

It appears that both metals form a carbonyl complex as well. Yet another dead end on a dangerous path. Ni(CO)4 is very toxic, flammable and volatile even at room temp. Everything I'd ever want in a metal complex (NOT). I won't lie, it IS tempting.

So far, the most success I've had is in dummy plating the copper (sponge) while electrochemically dissolving the metals. This procedure removed a significant amount of copper from the cell liquor and dissolved the coins in relatively short order. Still, I'd have to find a way to remove the remainder of the copper before I can call it 'doable'. Switching to a graphite anode toward the end should allow me to continue with removing copper without introducing more metal ions. If the anode happens to shred a little, filtration will remove the carbon fluff.

Tank

Mixell - 1-8-2011 at 01:15

What about iodide ions to precipitate the copper as CuI (Cuprous iodide, the decomposition of the extremely unstable CuI2, which serves as an intermediate in this reaction) Ni2+ should stay in solution, because NiI2 is stable.
All according to Wikipedia, but it seems right.

m1tanker78 - 3-8-2011 at 12:20

I'm iodine-poor so someone else will have to experiment with that. It'd be interesting to see it work experimentally. H2S is another one that would be reserved for a more dire situation.

Illustrated follow-up

m1tanker78 - 17-10-2011 at 20:48

My method of simultaneously dissolving US nickels (coins) and reducing out Copper(I) by electrolysis appears to have produced satisfactory results. HCl was the acid of choice for that experiment - cheap and available. Like most of my 'successful' experiments, I used my senses to make adjustments.

After it was all said and done, I had some concentrated nickel chloride solution with which I could perform some simple tests.
Below is a small sample:

After adding aluminum, the nickel 'crashed' out of solution and left a sludgy mess.

Sludge after drying:

Easily crushed to a fine powder:

Nickel powder influenced by a magnet (underneath):

Copper residue recovered from the process:

I'll try to reproduce this soon and note down voltage, current, temperatures, procedures, so on. Not bad for having started with 75%/25% Cu/Ni.

Tank

cyanureeves - 18-10-2011 at 15:49

congratulations! can't wait til you write up the whole process. google this very same thing and you'll see a whole bunch of theory but hardly anyone shows proof.

Neil - 18-10-2011 at 15:54

I second cyanureeves, that is nifty.

m1tanker78 - 23-10-2011 at 08:33

I dissolved a single nickel several months ago in concentrated HCl + 18% hydrogen peroxide. I don't know how long it took because I flat out forgot about this one. I had to add distilled water and a bit of HCl to redissolve the solid. It should have been roughly 3/4 copper and 1/4 nickel. I used aluminum metal to drop out the metal powders.

The experiment revolved around using a magnet to separate the dry, powdered metals. All attempts of doing so were unsuccessful. I tried dragging the nickel powder up with the magnet and tapping off the copper. I also placed the magnet on the cap of the container and gently shook. This method alone would constitute an unacceptable loss of nickel in each 'purification' step and will still have too much copper mixed in.

This would have made the electrolytic process much more hands-free but it didn't work out. I should have some free time today so I'll make another electrolytic batch and take some notes and pics.

Tank

blogfast25 - 23-10-2011 at 09:26

The acid of choice for both Cu and Ni is nitric acid. If you haven’t got any, try HCl + NaNO3. HCl alone is a slow boat to China for both.

To separate the precipitated metals purely on magnetism is a pipe dream, IMHO. 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.

Separation by sulphides is highly efficient but smelly.

Perhaps you could try taking advantage of Cu’s amphoterism? Cu(OH)2 does dissolve appreciably in strong NaOH/KOH, forming cobalt blue cuprate [Cu(OH)42-] anions. Ni(OH)2 does not. At 75/25 Cu/Ni that would take a lot of NaOH though…

m1tanker78 - 23-10-2011 at 10:35

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...

Tank

Sedit - 23-10-2011 at 14:04

Quote: Originally posted by m1tanker78

Sedit, the magnet test (which I forgot about) will be useful once I remove iron from the mix. Thanks for the suggestions. Just to clarify, you precipitated nickel powder from an acidic solution of nickel chloride with aluminum, right? Seems easy enough for a qualitative test for nickel...

+++++++++++++++++++++++

I dried a sample of the freshly prepared carbonate (same as the pic I posted before). For clarity, this carbonate is prepared from the ammonia complex, not from the original precipitate.

I haven't abandon you in your quest Tanker just been very busy. Its glad to see your making progress.

In reply to your quote may I offer another suggestion, Covert the carbonate powder you have to the respective chlorides and heat it till its anhydrous, Copper contamination will show a brown color where as anhydrous Nickle will be yellow. It should give you some idea if there is Cu in the mix to some extent.

I need to get on the ball repeating this experiment and documenting it completely just so back up my claims but like I stated before time has not really been on my side. That magnet test should be the most productive as far as determining the amount of Ni present but will not exclude excess Cu from carrying over into the magnetic particles. Me and Vesp tried in the past to use a magnetic stirrer to separate the two after precipitation but it proved relatively useless due to the small particle size. I may right more later when I get a chance but I just wanted to let you know I am still here watching along so when I get a chance to repeat I will offer more information as to the process I toyed with sometime ago on precipitating just the Ni out of the Cu/Ni complex.

Good to see someone more motivated then myself in this experiment, carry on the good work.
~Sedit

The Anode:

m1tanker78 - 23-10-2011 at 15:11

A late start; I tack welded 17 nickels together in the form of a fly swatter. The total starting mass is 84.2g:

It's important to make sure the welds penetrate all the way though so that there's good conductivity even as the anode erodes. Back side:

I used a sharpened graphite gouging rod - DCEN - Poor man's TIG

More to come...

p.s.: Sedit, glad you're following along.

Tank

Sedit - 23-10-2011 at 18:35

I just noticed something that may be of value to someone more apt then me in chemistry and I feel it has to do with the common ion effect of some sorts.

I mixed H2O2 and H2SO4, added 15 USA nickles and after a few hours noticed a shit load of CuSO4 precipitating with a green solution on top.... No evidence of precipitated Nickle sulfate has shown itself only the Copper sulfate.

Anyone care to elaborate on what I can't fully explain?

PS: I have seen this before with the sulfate however the green solution still contains SOME Cu salts yet much has been precipitated already so it may aid in separation in one form or another.

[Edited on 24-10-2011 by Sedit]

m1tanker78 - 23-10-2011 at 18:56

Sedit, I noticed the same thing:

Look familiar?? hehe

Sedit - 23-10-2011 at 21:28

VERY!.... Make that Exactly!

Reduction with Al proved there is still Copper ions in there but it is obvious much has precipitated. Im running it with Aluminum in a magnetic stirrer as we speak but there is still a significant amount of Copper present in the green solution.

Arthur Dent - 24-10-2011 at 04:46

Quote: Originally posted by m1tanker78

A late start; I tack welded 17 nickels together in the form of a fly swatter (...)
Tank

What type of electrode are you using? Tungsten? Titanium? Iron?Wouldn't this welding introduce another metal impurity to your solution? Even carbon arc electrodes have a copper shielding that would alter the metals present in your nickel anode...

Robert

blogfast25 - 24-10-2011 at 04:48

Both:

You've obviously exceeded the solubility limit of CuSO4 in your solvent, the excess of which crystalised out, but not the solubility limit of NiSO4 which remains in solution in its entirety: partial separation! Naturally the supernatant liquid is saturated with CuSO4 but it won't be that much, going by colour. There must be some common ion effect going on here.

Precipitate the supernatant liquid as mixed hydroxides, then treat with hot, stronng NaOH to leach out the copper as cuprate and Bob should be your uncle!

For the CuSO4, wash with cold water, then recrystallise once to get relatively pure copper suphate pentahydrate...

The FBI should now be onto both of you for destroying legal tender!

[Edited on 24-10-2011 by blogfast25]

m1tanker78 - 24-10-2011 at 06:23

Quote: Originally posted by Arthur Dent

Hi Robert, the copper cladding on gouging rods can easily be peeled or etched off. Our nickels contain 75% copper - of which >97% will be electrolytically removed so copper is hardly a contaminant. I didn't use any filler metal to tack the coins.

Tank

fledarmus - 24-10-2011 at 06:42

I think you're safe as far as the legal tender issue goes - Title 18, United States Code, 331 says only that it is illegal to "fraudulently" mutilate coins.

On the other hand, deliberately rendering a bill unfit for use does appear to be illegal, under Title 18, United States Code, 333.

Not legal advice, just one of those bits of trivia that accumulates after a while, thought it might be worth passing on.

m1tanker78 - 24-10-2011 at 06:51

The experiment is coming along nicely. I was afraid I'd develop the 'scaling up blues' but that hasn't been the case. Unfortunately I'm going to have to interrupt the electrolysis experiment until I get back home. I'll leave a few pics showing the progress and fill in details this evening...

The cell:

"First fuzz" and a hint of tint:

Just before the first cathode scrape:

What it looks like now. Nice emerald green:

Anode has thinned considerably (hard to see in the pic):

Stay tuned..

Tank

blogfast25 - 24-10-2011 at 08:52

Now that's what I call copper!

So you're selectively plating out the copper on a steel cathode, while the nickel enters into solution, as NiSO4 right?

What's the electrolyte... dilute H2SO4? And what voltage and amps are you reading?

Very interesting, Tank, much more so than the 'magnetic effort'...

cyanureeves - 24-10-2011 at 16:08

blogfast25 you wrote: precipitate the supernatant as mixed hydroxides? with what? this solution would have all of the nickel since copper solubility maxed out reducing the percentage of copper left in solution with nickel,right? the precipitate should then be heated with NaOH and leached out. leached out? how? filtered?poured out? will the nickel remain as insoluble hydroxide?

m1tanker78 - 24-10-2011 at 19:49

Quote: Originally posted by blogfast25

Now that's what I call copper!

So you're selectively plating out the copper on a steel cathode, while the nickel enters into solution, as NiSO4 right?

What's the electrolyte... dilute H2SO4? And what voltage and amps are you reading?

Very interesting, Tank, much more so than the 'magnetic effort'...

The initial electrolyte is HCl in distilled water. I set the power supply to provide 777mV. I add just enough HCl to pull 500mA initially. As time goes on, the current draw increases to as much as 2.5A. Copper(II) is selectively reduced to Copper(I) at the cathode. Nickel cannot exist as Nickel(I) in this setting which is precisely what's exploited. The copper can be removed by gently lifting the cathode out of the cell and scraping it into a separate container. This form of copper is very spongy so you'll undoubtedly drag out some acidic nickel solution with it.

p.s.: Blogfast, the 'magnetic effort' was something I had to get off my chest. Done.

Tank

m1tanker78 - 25-10-2011 at 17:21

I don't know when I'll be able to pick up where I left off on this experiment so I'm going to report what I have. I had to interrupt the electrolysis; I was aiming for 20 to 24 hours (or until the anode eroded away).

*The total run time was about 14 hours.
*Electrolyte was dilute hydrochloric acid.
*Voltage was 0.777V throughout.
*Current varied from 0.5A to 2.5A.
*Actual average power consumption (measured at the wall) was 51W x 14 hrs = ~$0.06.
*17 coins x $0.05 = 85 cents.
*The anode weighed 84.2g before and weighed 46g after (38.2g dissolved).
*I scraped the copper off the cathode (into distilled water) about 8 times throughout the experiment.

Here's a look at the solution:

And the copper (don't know the precise composition) removed periodically:

The copper is brick red when fresh and turns a lighter color as the dragged-in acid begins to slowly attack it. The water slowly takes on a blue color as the copper dissolves. The copper is extremely porous and now has a glittery appearance - probably micro crystals of metallic copper mixed in with Cu+??

The total cost of this experiment was less than a buck - if I don't count my time and labor.

Tank

cyanureeves - 25-10-2011 at 17:35

to not have to be at the mercy of online nickel compound sellers: priceless!

m1tanker78 - 25-10-2011 at 17:53

to not have to be at the mercy of online nickel compound sellers: priceless!

I have to disagree with you, Reeves. The fun factor and learning experience is priceless. I should probably take this opportunity to learn about Moles, Coulombs and knock some rust off of standard red/ox potentials. THAT would be priceless!

Tank

Sedit - 25-10-2011 at 18:33

Have you checked the precipitated Copper with a magnet to ensure that not to much Nickel is coming along for the ride?

blogfast25 - 26-10-2011 at 04:30

blogfast25 you wrote: precipitate the supernatant as mixed hydroxides? with what? this solution would have all of the nickel since copper solubility maxed out reducing the percentage of copper left in solution with nickel,right? the precipitate should then be heated with NaOH and leached out. leached out? how? filtered?poured out? will the nickel remain as insoluble hydroxide?

The solution at that point still contains small amounts of copper sulphate.

Precipitate the lot as Ni(OH)2/Cu(OH)2 with NaOH or KOH and filter. Treat the filter cake with strong NaOH or KOH (hot, preferably): as Cu(OH)2 is mildly amphoteric (it forms cobalt blue cuprate anions = Cu(OH)42-

it will enter solution, leaving behind the pure Ni(OH)2 (which isn't amphoteric). Wash profusely with clean hot water and the 'pure' Ni(OH)2 is yours. This method is especially workable if the copper is the minority constituent.

Tank:

Thanks for the data. Surprisingly low currents there but all in all a nice method for separating Cu and Ni. Nice work and well done!

In future and to keep run time lower, try larger electrodes, put closer together. Larger electrodes allow higher currents with acceptable current densities. Shorter distances between electrodes allows higher currents w/o overheating (ohmic power) because of lower overall cell resistance. In industrial electrolysis, electrodes are often frighteningly close together: time is money!

[Edited on 26-10-2011 by blogfast25]

[Edited on 26-10-2011 by blogfast25]

blogfast25 - 26-10-2011 at 04:37

Have you checked the precipitated Copper with a magnet to ensure that not to much Nickel is coming along for the ride?

Nickel is less ferromagnetic than Fe: small amounts would go undetected by magnetism. Also, if the applied voltage is correct, the separation should be near 100 %.

Neil - 26-10-2011 at 04:40

This is great... now for a condensed write up for the pre-pub folder...

m1tanker78 - 26-10-2011 at 06:59

Have you checked the precipitated Copper with a magnet to ensure that not to much Nickel is coming along for the ride?

Yes, I've poked around with a strong magnet in a ziplock back. Nothing was picked up but just to be clear, some Nickel(II) Chloride solution is dragged in from the cell at each cathode scraping. Most of it should dissipate after rinsing the copper 3 or 4 times with distilled water. It also helps to break-up the clumps before rinsing.

Quote:

Shorter distances between electrodes allows higher currents w/o overheating (ohmic power) because of lower overall cell resistance. In industrial electrolysis, electrodes are often frighteningly close together: time is money!

Blogfast: I agree with you about the electrode spacing but I'll point out that you'd run into a big problem by closing the gap. It won't take long for the copper deposit to bridge over to the anode and create a short circuit.

On that note, one of my next improvements would be to find a suitable 'bag' to hold the cathode. This would greatly facilitate the removal of the copper. As it is, it's hard to avoid dropping a little of it back in the cell each time the cathode is removed unless you scrape every 20 to 40 minutes.

Tank

blogfast25 - 26-10-2011 at 08:10

Tank:

Or some membane of sorts, to prevent shorting...

Chemistry Alchemist - 28-10-2011 at 23:19

Does the Copper/Nickel form an alloy that is soluble in HCl? i dissolved some coins in HCl, the colour started off as green (Nickel II Chloride) but after a while of dissolving still it turned a brown colour... i pored off about 50 mls and added Aluminum until it stopped fizzing, this produced a red precipitate that looks alot like Copper, would Nickel also be precipitated as well? could i add the precipitate to more HCl to dissolve the Nickel and leave the Copper as it is?

cyanureeves - 29-10-2011 at 06:42

wont aluminum drop all metals from solution?adding hcl would alloy the sponge back into solution unless your solution remained nickel green but i bet it went clear.why dont you add hydroxide to the chloride solution like blogfast25 suggests then strong heat the dropped hydroxides with more hydroxide and leach?plante1999 and i have successfully separated chromium from iron this way.maybe nickel or copper will remain as insoluble hydroxide.

Chemistry Alchemist - 29-10-2011 at 06:53

Yeah Aluminum will precipitate all metals but Copper is not reactive to HCl so why would i have copper chloride in the solution?

cyanureeves - 29-10-2011 at 07:04

yikes! i understand your initial question. i will go back to the bleachers now!

blogfast25 - 29-10-2011 at 09:19

maybe nickel or copper will remain as insoluble hydroxide.

If there's enough alkali, Cu(OH)2 will enter solution as Cu(OH)42-, cuprate, a cobalt blue solution. Copper is slightly amphoteric. Nickel isn't at all, it stays as Ni(OH)2.

Chemistry Alchemist - 29-10-2011 at 09:35

could i just add more HCl to the Copper/nickel powder to dissolve the nickel and leave teh copper powder?

m1tanker78 - 29-10-2011 at 10:07

Quote: Originally posted by Chemistry Alchemist

could i just add more HCl to the Copper/nickel powder to dissolve the nickel and leave teh copper powder?

That won't work because both metals will enter solution as +2. You can dissolve the metal powders and try a mild reducing agent. I used electrolytic Hydrogen to selectively reduce Cu(II) to the insoluble Cu(I). I'm still working on the half cell reactions and so forth. Other reducing agents have been proposed and suggested upthread.

Tank

m1tanker78 - 1-11-2011 at 11:30

Here are some progress images on the copper side of the electrolytic process. Bear in mind that I didn't take much care in rinsing or otherwise preventing oxidation. I rinsed once with tap water (not recommended) and once with distilled water (better).

Final settling before decanting off as much of the water as possible. I wasn't able to figure out what the crystals were. Crystals of alkali carbonates (from the tap water) seem out of place in this setting:

Copper 'mud' after decanting off water:

Lack of care on the rinsing step (IMO) quickly manifests itself upon exposure to air:

I suppose it isn't too late to give it a proper rinse but I'll leave it as is for now. Most of the drab green areas seen in the above pics now appear (edit) lime green after several days.

Tank

[Edited on 11-1-2011 by m1tanker78]

blogfast25 - 1-11-2011 at 12:56

Very nice! Now wash, dry and melt down ;-) !

Sedit - 1-11-2011 at 19:09

I am finding the Sulfate salt of the dissolved coins WILL precipitate with the addition of Ammonium hydroxide ONLY the Nickle compounds.

A wash of Ni and Cu sulfates took up ONLY the Cu and left the Ni as a sludge.

I want to personally ask Woelen to perform a simple inorganic experiment for us all if he may since I have great faith in his skills, materials and attention to detail. I wish to know the solubility of CuSO4 in concentrated H2SO4... also the solubility of NiSO4 in concentrated sulfuric acid.....

In the end I would like to see these two saturated solutions mixed together. I have a huge suspicion that the Copper sulfate will drop out due to A common ion effect leaving the Nickle behind much more pure then before still in solution.

I can not seem to keep CuSO4 in the solution of sulfates very well.... it keeps precipitating out and this could be a great find IMO.

PS: I cant edit pictures very well at all and many are quite large, I miss posting pictures... If I posted some very large ones is there any Mods willing to resize and fix them for me so I can actively post my progress in various areas until my computer is fixed?

[Edited on 2-11-2011 by Sedit]

blogfast25 - 2-11-2011 at 03:41