Nemo_Tenetur
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Mixed cobalt oxide
Hi!
From a private pottery supply I got about 350 gram cobalt oxide. I want to dissolve it in perchloric acid and during the calculations it turned out
that it isn´t so easy.
The supplier claimed it to be "CoO", but I doubt that this is true.
According to wikipedia, CoO is dark olive to gray, wheras my sample is black. So I guess it´s at least contaminated with a considerable amount Co +3,
forming mixed Oxides commonly called Co3O4 or CoO*Co2O3.
https://en.wikipedia.org/wiki/Cobalt_oxide
I want to avoid perchloric acid residues and I decided to add a little bit less than stoichometric amount, calculated as "CoO" as I don´t know how
much Co+3 it contains.
Which part of the mixed cobalt oxide will dissolve easier, the Co+2 or Co+3?
I guess it´s finally a mixed solution of Co(ClO4)2 and Co(ClO4)3 or what will happen with the Co+3 fraction?
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Boffis
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Again, I suggest you try using a mixture of hydroen peroxide and acid. The hydrogen peroxide acts as a reducing agent in this case. Other reducing
agents often work eg ascorbic acid, sodium metabisulphite, HI etc but they are less "clean" and you usually need to precipitate the Co2+ as a
carbonate, filter it off and then redissolve in the appropriate acid. Alteratively dissolve it first in boiling 18% hydrochloric acid (tends to be
slow) in a fume cupboard or out side until chlorine evolution ceases, then go to the perchlorate via the carbonate. Avoid precipitating as the
hydroxide as, like manganese, there is a tendency for it to oxidised spotaneously to hydrated Co3+ oxides.
Excess H2O2 is simply removed by evaporation.
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Nemo_Tenetur
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Thank you for feedback and explanation. Is there any visible change between cobalt (II) perchlorate and cobalt (III) perchlorate?
BTW, I was unable to find any information about pure cobalt (III) perchlorate. It seems to be not commercially available and I found only one brief
description about it´s preparation from cobalt chloride and AgClO4 (I guess metathesis, double displacement?).
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DraconicAcid
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Cobalt(III) isn't stable without the right ligands. Aqueous cobalt(III) will readily decompose to give cobalt(II), and you're not going to be able to
keep cobalt(III) perchlorate as aqueous solution or isolate it as a solid.
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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woelen
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Your black material most likely is Co3O4. This stuff is amazingly hard to dissolve. I once tried in conc. HCl, 30% or so. After keeping the material
in boiling hot acid, I only managed to get a pale blue solution, meaning that there were only traces of cobalt in the acid (e.g. CoCO3 immediately
gives an intensely colored sky blue solution in conc. HCl and even a thin film of this solution on glass has a bright blue color).
I think that in HClO4 you won't be able to dissolve this black oxide. I tried in 30% H2SO4 and this did not work at all. Same with HNO3. Only worked
with HCl.
Cobalt(III) is not stable in water, unless some specific ligands are present. Hydrated Co(III) decomposes, giving oxygen and Co(II). With cyanide ion,
you can get a stable complex, [Co(CN)6](3-), which is pale yellow.
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Nemo_Tenetur
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It´s really surprising how much the statements about the solubility of cobalt oxide differ. After the successful dissolution of my green nickel oxide
(48 - 72 hour refluxing with 40-50 % perchloric acid) I decided to reproduce this with cobalt oxide. Hopefully looking forward to get it dissolved,
otherwise I have a not-so-pleasant workup to do (filter dust-like black cobalt oxide out of semi-concentrated perchloric acid and perform a vacuum
distillation because I don´t want to discard the expensive acid).
In this case, I don´t know how to proceed with the hard-necked cobalt oxide. Maybe It´s easier and cheaper to discard the cobalt oxide and buy some
cobalt carbonate as precursor.
But cobalt carbonate is expensive, much more than nickel. So I´ll let it boil and stirr, if necessary for a week or two, no problem. The heating is
with an electrical mantle, the reflux condenser is conncected with a big water reservoir and the reflux rate is very low, less than a drop per
second. It´s not necessary to heat it stronger - just a waste of energy. Next week I´ll present the results ...
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woelen
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I have been disappointed with metal oxides quite a few times (not soluble in aqueous strong acid, nor in aqueous strong base) and now I have the
following oxides lying around, without good use:
- Cr2O3
- Fe2O3
- Co3O4
- SnO2
- TiO2
- Nd2O3
- Ta2O5
- Al2O3
- BeO
- TeO3
With other oxides I had success dissolving them:
- V2O4
- V2O5
- GeO2
- Ag2O
- CuO
- ZnO
- MnO2
- MoO3
- WO3
- Bi2O3
- TeO2
- CaO
- BaO
- BaO2
- diverse lanthanide oxides, including another sample of Nd2O3
Non-metal oxides dissolve easily, sometimes even in water and borderline oxides (e.g. SeO2, B2O3) also can be dissolved easily.
Right now, if I want metal compounds, then either I buy the carbonate if possible, or some water soluble compound (sulfate, chloride or nitrate).
Buying oxides of rarer/expensive metals takes too much risk for me, in that I add another useless oxide powder to my collection.
[Edited on 13-3-25 by woelen]
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DraconicAcid
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For the last couple of years, we've used V2O5 in one of our undergraduate labs- dissolve it in base, then acidify to give a nice yellow solution. We
got a new batch last month, and it just refused to dissolve in base without being near-boiling for half an hour (which really mucked up our three-hour
lab period).
I wonder if some of your useless oxides could be dissolved by fusion with sodium hydroxide, or reduced with aluminum or heating under hydrogen gas.
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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Nemo_Tenetur
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I´ll stop my dissolution efforts today after 48 hour refluxing. The reason isn´t the slow dissolution rate but a more and more detectable smell of
chlorine compounds, which escape out of the (very effective, total length 45 cm with water mantle and helix inside) reflux condenser.
It´s definitely not perchloric acid what escapes, it´s a degradation product. Maybe hot perchloric acid is affected by the Co +3 fraction of the
oxide?
I don´t know what kind of chlorine compounds are evolved. But I know that many of such chlorine oxides are explosive (at higher concentrations at
least), poisonous and - most important for me - I don´t want to destroy my expensive perchloric acid.
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bnull
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Quote: Originally posted by Nemo_Tenetur  | | It´s definitely not perchloric acid what escapes, it´s a degradation product. Maybe hot perchloric acid is affected by the Co +3 fraction of the
oxide? |
I'd say it is the other one. You're oxidizing Co2+ to Co3+ and turning perchloric acid into whatever lower acids and oxides of
chlorine are possible. Then you have a hot, acidic solution that gives off chlorine and its oxides.
Read this: R. C. Hubli, J. Mittra, A. K. Suri, Reduction-dissolution of cobalt oxide in acid media: a kinetic study (Hydrometallurgy, 44, 1 (1997),
125-134, https://doi.org/10.1016/S0304-386X(96)00036-9). They used 4 M sulfuric acid but I suppose you can get away with HCl; the important part is the
addition of sulfite. Bisulfite (with an i) will do if you happen to have any, or any stronger reducing agent that works in acidic solution.
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teodor
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For reacting with an acid a compound should be soluble in that acid.
NaOH pellets don't react with concentrated H2SO4 because they are soluble in water but not soluble in concentrated H2SO4.
The same way most of carbonates don't react with organic acids or reacts slowly without completion.
That means that it's not enough to be an acid and oxide, that relation can only work if water is present as a solvent.
Co2O3 is not soluble in any solvent which is liquid at room temperature.
But it is soluble in some "solvents" which you never can think it could be a solvent.
For example Co2O3 is dissolved by molten SiO2 exotermically starting from 650C with decomposition to O2 and CoO.
At lower temperature (400-500C) it reacts with Na2O forming Na4CoO4.
So, if you have some "insoluble" oxides check their solubility in molten salts. After getting the oxide in solution you can make some reactions
considering the property of the molten salt which could be different from water but still working by similar rules.
@woelen I have some interest to do experiments with some of your "insoluble" oxides. Those with not so common elements are most interesting.
[Edited on 14-3-2025 by teodor]
Some reactions with acids are actucally oxidation/reduction reactions and they go by different rules. For example, dissolution of Zn in H2SO4. I am
not sure, from the point of crystallography, are the rules of oxidation of metals are the same as ionic compounds. As far as I know oxidisers H2O2,
HNO3, NOCl could be used with many metals but I can't remember any examples with oxides of lower valency (which are not soluble in not-oxidising acids
but soluble in oxidised ones). As I guess, to be oxidised by acid compound should be an electrical conductor. Or is it more complex mechanism of solid
oxidation?
[Edited on 14-3-2025 by teodor]
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Nemo_Tenetur
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Update
After about two days refluxing and cooling down to room temperature (which is at the moment between 8 and 10 degree centigrade) the undissolved black
cobalt oxide settled down followed by crystallization of assumed cobalt perchlorate as long needles.
The supernatant solution is really dark red, almost black, with a tingle salmon to brown colored. Even a thin layer is intensely colored.
After suction filtration I´ve recovered about 260 gram Cobalt oxide, that means about 90 gram went into solution. But the cobalt oxide is still not
completely dry, so maybe a few grams more went into solution.
I´ve placed the filtrate in the freezer for several hours which lead to additional crystallization of long needles. Tomorrow I´ll filtrate the
cystals and try to dry it . It seems that the dissolved cobalt perchlorate has a much more intense coloration compared to the needle-shaped
crystals.
Then I intend to reduce the volume of the filtrate under reduced pressure (vacuum distillation) and put it in the freezer again to get a second crop
of crystals.
I do not dare to evaporate to complete dryness. Cobalt perchlorate hexahydrate is heat-sensitive and I doubt that I could get it dry without at least
partial decomposition. There must be a large quantity of unreacted perchloric acid in the solution. The complete separation without partial
decomposition is the next challenge for me.
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woelen
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Good to see that you get some cobalt(II) perchlorate from your oxide. It does dissolve, but you need a lot of patience. I did not have that patience,
when I experimented with Co3O4 ;-)
I would not worry too much about decomposition of the cobalt perchlorate on drying. Perchlorate ion really is stable, certainly up to 100 degrees.
Only covalent perchlorates (anhydrous acid, organic esters) are prone to decomposition, or even explosion. But if you dry your material at 50 C or so,
you'll never get anhydrous acid (the acid forms an azeotrope with water at 72%, which boils a little over 200 C).
Another mode of decomposition is possible at elevated temperature, and that is internal hydrolysis. If you heat it too much, then water of
crystallization will combine with perchlorate ion, giving free azeotropic acid, leaving a basic salt behind (mixed perchlorate/hydroxide/oxide). I
myself once had this, when making copper nitrate. On heating, the deep blue solid turned lighter blue/green and the dry powder was insoluble in water,
it was something like Cu(OH)NO3. Water and HNO3 vapor boiled off. In your case, with perchloric acid, this is less likely to happen, because the
azeotrope boils well over 200 C, so if you keep the temperature lower than 100 C, no acid will boil off. At reduced pressure, some acid may boil off
and then the combination of heating and reduced pressure may lead to internal hydrolysis.
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Nemo_Tenetur
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The first crystallization crop was about 140 gram (wet crystals as I didn´t know how to remove better than with suction filtration). Then the still
dark red (but not black) filtrate was vacuum distilled to concentrate it. About 500 ml ( almost pure water ) was distilled off in the temperature
range 20 - 40 degree centigrade. At the end of the distillation suddenly crystals appeared in the still hot RBF.
Cobalt perchlorate seems to possess a much lower solubility in azeotropic perchloric acid than in dilute (about 40 %) perchloric acid.
Then the RBF was placed in a freezer over-night and next day I got a whooping scond crop (more than 300 gram after suction filtration, also still wet
with perchloric acid). Now the filtrate was only pale red, indicates that most of the cobalt perchlorate crystallized out and only a minor quantitiy
remained in solution.
Again, the filtrate was vacuum distilled for further concentration and a possible thirth crop of crystals. The vacuum distillation temperature for
the azeotropic acid was between 105 and 106 degree centigrade.
The distilled perchloric acid wasn´t colorless but pale yellow with a distinct smell resembling a mixture of chlorine and bleach.
.
Some crystals appeared at the end of the distillation in the still hot RBF, but the quantity is much much lower. I doubt that it is more than a few
gram.
The RBF is now in the freezer and I´ll perform suction filtration next week.
Two problems are still waiting for a solution:
1.) How can I purify the perchloric acid to remove the pale yellow color ? Vacuum distillation doesn´t help.
2.) How can I remove the perchloric acid from the wet crystals? I doubt that the perchloric acid is volatile enough at room temperature. It´s even
more likely that the perchloric acid is hygroscopic.
Normally, I would flush the crystals, but I can´t use water (highly soluble) nor can I use alcohol (explosive mix with perchloric acid) or any other
solvent I can think about.
Ideas/suggestions?
.
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bnull
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| Quote: | | 1.) How can I purify the perchloric acid to remove the pale yellow color ? Vacuum distillation doesn´t help. |
Try bubbling air into the solution. It looks like chlorine dioxide. I think Brauer has something on it.
| Quote: | | 2.) How can I remove the perchloric acid from the wet crystals? I doubt that the perchloric acid is volatile enough at room temperature. It´s even
more likely that the perchloric acid is hygroscopic. |
What about removing it as potassium perchlorate? As far as I know, it doesn't form a double salt with cobalt(ii). Dissolve, neutralize, chill, filter
and let it crystallize. What with the common ion and so on, and you can recover the acid from KClO4 later anyway.
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woelen
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Getting rid of perchloric acid from your cobalt perchlorate indeed can be done by bnull's method. I would use KHCO3 for neutralization, not KOH. This
assures that there will not be strongly alkaline hotspots in which Co(OH)2 will be formed. In strong alkalies, cobalt can easily be oxidized by oxygen
from air, producing basic mixed oxidation state impurities, which are hard to get rid of. So, I would use a solution of KHCO3 and neutralize with
that. Allow a slight excess (just over exact neutralization) of KHCO3 to neutralize all acid. You get a precipitate of KClO4 and a tiny amount of
CoCO3. Cool down the liquid to precipitate all KClO4. After filtering your solution you will have a solution of Co(ClO4)2 with only a very little
amount of potassium ions in it.
Recovering HClO4 from KClO4 is extremely dangerous and I would not recommend doing that. You need to distill HClO4 from a mix of concentrated H2SO4
and KClO4, but the acid you distill is anhydrous and can explode without any apparent reason.
Purifying your precipitated wet KClO4 from any cobalt in it can be done easily, however. Add a litte acid to the solution (e.g. a few drops of H2SO4
or HNO3) and then add boiling water, so that the solid just dissolves. Then allow cooling down and if at room temperature allow cooling down further
in a fridge (but do not let it freeze). Filter the solid and suck away as much of the water. Then recrystallize a second time from boiling water and
do the same cooling step again. After this, you will have nice snow-white KClO4 and you'll recover more than 90%.
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bnull
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| Quote: Originally posted by woelen | | You need to distill HClO4 from a mix of concentrated H2SO4 and KClO4, but the acid you distill is anhydrous and can explode without any apparent
reason. |
I was thinking of an alternative, aqueous route, something I saw in Schilt's Perchloric Acid and Perchlorates. It ended up being
H2SiF6, so never mind that.
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Nemo_Tenetur
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Thirth batch different
Today I worked the very small thirth batch up. It´s obviously different to the first two batches, much lighter in color.
See attached jpeg.
It´s just speculation, but could it be some cobalt chloride , chlorate, hypochlorite or the like?
If you prepare, for example, silver perchlorate by heating of a silver compound with concentrated perchloric acid, some silver chloride is formed as
a by-product.
So my guess is that some cobalt chloride is formed as a by-product.
And, by the way, the combined first two batches (right part of the jpeg) attracted moisture from the air.
The adherent perchloric acid is no longer azeotropic, considerably diluted. So I think about a different way to remove the perchloric acid, dissolve
in distilled water and crash it out of solution with plenty amounts of IPA. Wikipedia states that cobalt perchlorate is insoluble in ethanol.
Therefore a higher alcohol with a longer unpolar carbon chain should further reduce the solubility, even if water is present in the mixture.
I´ll perform a small scale test with appropriate safety precautions (kevlar gloves, face shield, safety goggles etc.). I doubt that dilute
perchloric acid (below azeotrope) is as dangerous as concentrated or even anhydrous acid, but I´m not sure.
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woelen
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Diluted acid will not lead to explosion with organics. I have mistreated dilute perchloric acid with lots of organics and even made solid salts of
them (e.g. all kinds of amine-HClO4 salts).
As long as the perchlorate ion is present as ion and temperatures remain below 60 C or so, you're pretty safe.
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Osmiridium
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Oh, I know these issues with refractory metal oxides very well from my own experience.
I found that metal oxides that are heated to very high temperatures, possibly even to the sintering or melting point, somewhere during production,
tend to be quite difficult or just slow to dissolve. It may have something to do with porosity reduction and the effective removal of more reactive OH
ions as many "oxides" obtained by wet chemistry are never really well defined pure oxides, even after drying, but still contain varying fractions of
hydroxide.
If the metal ion can be further oxidized an oxidizing melt can help (works e.g. with Cr2O3 or MnO2 in KOH/KNO3). If the oxide is more basic an acidic
or if acidic a basic melt can help. (e.g. NaHSO4 or Na2CO3/Hydroxide respectively). It can be quite frustrating sometimes.
[Edited on 25-3-2025 by Osmiridium]
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