Nemo_Tenetur
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Best way to get green nickel oxide into solution ?
Hi!
I got some very cheap green nickel oxide (fine powder, like dust, great surface) and want to prepare different nickel salts.
Unfortunately, this batch of nickel oxide refuses to go into solution. According to different web sources, the solubility of nickel oxide can differ
greatly between easy soluble and almost insoluble. Generally, light green nickel oxide should be better soluble as dark brown or even black nickel
oxide.
I´ve refluxed a sample with an excess 80% acetic acid overnight and the acid is still almost colorless. Same happens even with perchloric acid (about
40-50%, I do not dare to reflux azeotropic 70 % for a prolonged time).
Ideas/suggestions?
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Precipitates
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Can you try a stronger acid e.g., sulphuric, nitric or hydrochloric, and then perform a double displacement to get the desired salt?
Probably what I'd try if the nickel oxide is just refusing to dissolve in weak acids.
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Nemo_Tenetur
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I thought perchloric is one of the strongest acids and nickel perchlorate is also very well soluble? I also found a statement that nickel oxide is
soluble im ammonia water, maybe i should try this even if i´m not convinced.
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Nemo_Tenetur
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This is the source of information about nickel oxide green and black:
https://www.todini.com/chemicals/nickel/nickel-oxide-green
https://www.todini.com/chemicals/nickel/nickel-oxide-black
The German Wikipedia states that nickel oxide is soluble in ammonia water and refers to "Georg Bauer, handbook of preparative inorganic chemistry"
whereas the English Wikipedia states that it is soluble in KCN (what I really don´t want to try).
https://de.wikipedia.org/wiki/Nickel(II)-oxid
https://en.wikipedia.org/wiki/Nickel(II)_oxide
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Precipitates
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Quote: Originally posted by Nemo_Tenetur  | | I thought perchloric is one of the strongest acids and nickel perchlorate is also very well soluble? I also found a statement that nickel oxide is
soluble im ammonia water, maybe i should try this even if i´m not convinced. |
It is*, but it appears nickel oxide doesn't dissolve in perchloric acid as well as other strong acids such as sulphuric (1).
I believe senior forum members have previously produced nickel perchlorate from nickel oxide (2), so they may chip in on how they did this/how this
went etc.,
1. Effect of Solid State Impurities on the Dissolution of Nickel Oxide
2. Uses for nickel (II) perchlorate hexahydrate
*Sorry previous post was misleading, perchloric acid is a very strong acid, stronger than those I mentioned.
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Nemo_Tenetur
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Thank you for this information. The first link doesn´t work, but the second work well. Maybe @woelen can give more details about the preparation from
nickel oxide and perchloric acid.
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Precipitates
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Attachment: Effect of Solid State Impurities on the Dissolution of Nickel Oxide.pdf (625kB)
This file has been downloaded 42 times
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clearly_not_atara
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I would definitely prefer hydrochloric acid, since chloride is a weak ligand for nickel. Perchlorate on the other hand is a highly non-coordinating
anion.
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bnull
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| Quote: Originally posted by Nemo_Tenetur | | The German Wikipedia states that nickel oxide is soluble in ammonia water and refers to "Georg Bauer, handbook of preparative inorganic chemistry"
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We have an old English edition of Brauer's in the Library. Nickel (ii) oxide is in page 1548.
[Edited on 29-1-2025 by bnull]
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Boffis
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Have you tried adding a little hydrogen peroxide? This certainly assists in the dissolution of Co and Mn oxides though it doesn't seem to make much
difference to Fe3+ oxides.
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DraconicAcid
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| Quote: Originally posted by clearly_not_atara | | I would definitely prefer hydrochloric acid, since chloride is a weak ligand for nickel. Perchlorate on the other hand is a highly non-coordinating
anion. |
Agreed. You probably need a ligand to help get the nickel into solution- hydrochloric acid, or concentrated ammonia/ammonium chloride.
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
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bnull
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You don't need the nickel salts right away, do you? Put a little oxide into a vial, add some acid and leave the vial alone for a couple of days or a
week. Do that with the acids you have (and ammonia too) and see later what dissolved the oxide the most. Give them time and go do something else.
I remember some iron oxide sold as pigment that refused to dissolve in cold, hot, and boiling hydrochloric acid. I gave up and forgot the matter. Some
time later, while digging through my stuff, I saw that a good part of the oxide had dissolved.
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Nemo_Tenetur
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| Quote: Originally posted by bnull | You don't need the nickel salts right away, do you? Put a little oxide into a vial, add some acid and leave the vial alone for a couple of days or a
week. Do that with the acids you have (and ammonia too) and see later what dissolved the oxide the most. Give them time and go do something else.
I remember some iron oxide sold as pigment that refused to dissolve in cold, hot, and boiling hydrochloric acid. I gave up and forgot the matter. Some
time later, while digging through my stuff, I saw that a good part of the oxide had dissolved. |
That´s a good idea! I´ll give the ammonia solution a try, put it in a closed bottle and make a review in a month or two. In contrast to acetic and
perchloric acid, I can´t reflux it as I have no pressure-resistant equipment (otherwise most of the ammonia goes off).
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woelen
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I have different excperiences with different oxides, even of the same metal.
I have black nickel oxide (I think it is a higher oxidation state or mixed oxidation state oxide), and it dissolves with great difficulty in acids. I
also have NiO, and my oxide dissolves easily in strong acids (I made solutions of nickel perchlorate from this, used for doing experiments with nickel
complexes of amines).
The solubility of a metal oxide strongly depends on how it is prepared. If it is strongly heated during preparation, then it becomes very very inert.
I have two samples of Nd2O3, both looking very similarly (grey powder with a purplish hue). One of these samples dissolves in a strong acid quickly,
and even hisses a little bit on addition to the acid, and leads to noticeable heating up of the liquid. The other sample does not dissolve at all. Not
even an hour of refluxing with any of the common strong acids gives more than trace amounts of Nd in solution. I have a similar experience with my
dark red Fe2O3, green Cr2O3, and white SnO2.
You indeed could put the nickel oxide aside in concentrated HCl (30% by weight) and leave it that way for days or weeks. I think that HCl is the best
option, because it forms complexes, which may help dissolving the oxide. In my experience, HNO3, H2SO4 and HClO4 all are less suitable for dissolving
hard to dissolve oxides, because their anions are not coordinating to the metal ions.
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Sulaiman
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as aqua regia dissolves both nickel and nickel oxide, I'd give it a go.
CAUTION : Hobby Chemist, not Professional or even Amateur
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Nemo_Tenetur
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Thank you very much for this detailed writeup. The oxide preparation temperature is critical, without question, but impurities are also important if
I understand the pdf-document from the user "Precipitates" correct.
The source of my cheap nickel oxide is pottery, therefore contaminants are very likely.
Chromium impurities act as inhibitor, whereas lithium promotes dissolution. Maybe I should add a trace amount lithium salt?
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Axt
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If you hadn't seen my post from before.
"1.5g (0.02mol) nickel oxide and 11.5g (0.08mol) 70% perchloric acid in 20ml water.
Heat with stirring on 150 degree hotplate (solution temp 75 rising to 90) until a slurry with no black oxide left. Dilute and redissolve with 50ml
water for a clear emerald green."
Unfortunately you do need to heat it but there's no problems with aqueous perchloric acid at these temperatures. I used "pottery" black NiO. I had to
use the 2x excess HClO4 otherwise it just wouldn't fully react. I've used 32% HCl and 68% HNO3 too and let it heat to dry without issue.
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RU_KLO
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maybe if you want to try another route (not tried by myself) but why not...:
use oxalic acid, at 70-90°C for at least 5 hs. (refluxing)
Read somewhere that "complexation power" of oxalic acid helps with the solution of some difficult metal oxides.
"Dissolution of Nickel Oxide in Oxalic Acid Aqueous Solutions"
https://www.sciencedirect.com/science/article/abs/pii/S00219...
Then you will have insoluble nickel oxalate (or maybe some "Oxalatonickelate" https://en.wikipedia.org/wiki/Oxalatonickelate).
Treat this with a calcium chloride strong solution to get insoluble calcium oxalate + soluble nickel chloride
Go SAFE, because stupidity and bad Luck exist.
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Nemo_Tenetur
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Update
Just for your information, a prolonged refluxing with about 35 % perchloric acid for 48 - 72 hours dissolved most of the nickel oxide.
After cooling to room temperature, a vacuum filtration with a buchner funnel yielded a heavy somewhat viscous clear dark green solution.
I´ve done two batches. First batch with about 48 hours refluxing yielded a dark green solution with a measured density of 1.49 gram per milliliter at
seven degree centigrade. It didnt´t crystallize after three days at 7 degree centigrade. So the whole batch was evaporated to dryness at about 40 -
50 degree centigrade in a very large open pan. Yield about 60% of the weight from the heavy solution.
The second batch was refluxed for about 72 hours, the nickel oxide went almost completely into solution, only a few percent residue. I´ve used an
almost stoichometric quantity of perchloric acid (no excess, it´s much more expensive and valuable than nickel oxide) and I guess it was almost
completely used.
The measured density of the dark green solution of this batch was 1.53 gram per milliliter. It´s still evaporating and I expect a slightly better
yield, maybe somewhat above 60% (calculated by the weight of the solution by evaporation to dryness). The total conversion rate (nickel oxide to
nickel perchlorate) is probably about 95% .
Thank you @Axt for the perchloric acid refluxing suggestion!
@woelen I think there´re some typos in the corresponding sciencemadness wiki article, see attached screenshot.
"Nickel(II) nitrate is the chemical compound with the formula Ni(ClO4)2."
I doubt that the solubility data is correct. From my experience, a nickel perchlorate solution with a concentration of more than 60% by weight is
possible even at 7 degree centigrade. So even if the solubility is calculated as the anhydrous salt , it´s too low.
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woelen
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Nickel perchlorate is extremely soluble in water. I have a wet slurry of nickel perchlorate 6-hydrate. It is a wet mass of green crystals under a
layer of dark green solution. Never tried to dry it further, I only use this stuff in aqueous solution, so no need to first dry it.
It is good that you used excess nickel oxide. You don't want excess acid in the solution. But even with the excess nickel oxide, I can imagine that
the solution still is quite acidic. You could try testing it with pH-paper, or with a little solid NaHCO3.
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Nemo_Tenetur
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Yes, I don´t want acid residue. I´ve dried it in a large open pan and finally (required almost two days) got, after grinding with a stainless steel
spon, apple-green crystals, free flowing. Therefore I think the acid residue is negligible, but I should confirm this with pH-paper.
Does perchloric acid have a detectable smell at moderate elevated temperatures (about 50 degree centigrade)? I´ve always avoided to test it because
it´s said to be very corrosive to organic tissue.
Anhydrous perchloric acid is said to be a volatile fuming liquid, very dangerous, whereas somewhat diluted solutions (the commercial 60 % and
azeotropic 70 %) aren´t fuming and more viscous.
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woelen
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Perchloric acid at 70% or less is completely odorless, even if somewhat heated. It does not give off any fumes (other than water vapor) at all, not
even at somewhat elevated temperatures. The hydrated acid has a high boiling point (well over 200 C), that's why it is not fuming at all.
I aslso noticed that perchloric acid is not extremely corrosive to skin, not more so than e.g. HCl. of course, you should not allow it to remain on
your skin for minutes, but if you spill something on you, it has no immediate effect like sulfuric acid. It is more like getting some HCl on your
skin. Just rinse it off with some water and no further effect is noticed.
Anhydrous acid is another thing. It is very dangerous and fumes a lot in contact with air.
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Nemo_Tenetur
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Thank you for explanation. That means, even a considerable residue of perchloric acid is not detectable by smell.
By the way, do you know if the hexammine complex of nickel perchlorate is labile? I´ve dissolved the residue of my nickel perchlorate in the bowl
and glassware with several milliliter distilled water and added excess 25 % ammonia water. Immediately a deep blue complex precipitated out and was
filtered yesterday in the evening, followed by washing with distilled water (no more ammonia smell was detectable) and I let it stand outside in the
cold.
Today the precipitate wasn´t dry (that´s no surprise as the temperature outside was somewhat below zero degree centigrade), but it was partially
discolored back to greenish.
See attached jpeg.
I also noticed some weak ammonia smell.
Is it possible that it loses some ammonia? I have no other explanation for this oberservation. I guess (pure speculation!) that the ammonia loss was
only partially, not completely, because otherwise the extremely soluble nickel perchlorate probably wouldn´t precipitate in such a wet environment.
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