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AJKOER
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Here is a new approach. Use the apparent fact that Citric acid can effectively leach copper oxide. The following extract (see http://link.springer.com/article/10.1007%2Fs12613-012-0628-9 ) provides some detail:
"Leaching of an oxidized copper ore containing malachite, as a new approach, was investigated by an organic reagent, citric acid. Sulfuric acid is the
most common reagent in the leaching of oxide copper ores, but it has several side effects such as severe adverse impact on the environment. In this
investigation, the effects of particle size, acid concentration, leaching time, solid/liquid ratio, temperature, and stirring speed were optimized.
According to the experimental results, malachite leaching by citric acid was technically feasible. Optimum leaching conditions were found as follows:
the range of particle size, 105–150 µm; acid concentration, 0.2 M; leaching time, 30 min; solid/liquid ratio, 1:20 g/mL; temperature, 40°C; and
stirring speed, 200 r/min. Under the optimum conditions, 91.61% of copper was extracted."
So dissolve an excess of powdered Cu in Citric acid/dilute H2O2 solution at 40 C with stirring. Replenish the H2O2 as needed. Add Na2CO3 to form a
precipitate of Copper carbonate. Treat with acetic acid to create the Copper acetate.
Alternative Synthesis: As oxidation of Cu maybe a key and difficult step, alternately, one could heat the powdered Copper (or, a large number of
copper pennies as we will be using only the surface oxide) in air and/or treat with NaOCl. Rinse and treat with Citric acid (or Citric acid/H2O2 as
before) and contnue with the prior synthesis. Namely, add Na2CO3 to form a precipitate of Copper carbonate. Finally, add acetic acid to the CuCO3 to
create the Copper acetate.
[EDIT] As Wikipedia cites the following reaction:
2 CuSO4 + 2 Na2CO3 + H2O → Cu2(OH)2CO3 + 2 Na2SO4 + CO2
most likely we will be treating Basic copper(II) carbonate with Acetic acid and not CuCO3.
[EDIT EDIT] However, the reaction with Baking Soda, NaHCO3, could be more favorable forming CuCO3(?) in conc solutions.
CuSO4 (aq) + NaHCO3 --> CuCO3 (s) + NaHSO4 (aq)
so I would substitute NaHCO3 for Na2CO3 in the synthesis.
[Edited on 23-2-2013 by AJKOER]
[Edited on 23-2-2013 by AJKOER]
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KonkreteRocketry
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So i just keep to put CuO in Acetic acid and heat it a bit ?
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Nathaniel
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Quote: Originally posted by blogfast25  | | 7.2 g/100 ml in cold water isn't that poorly soluble (Wiki). And 20 g/100 ml in hot water (Wiki) is someting that could be exploited: from a hot,
saturated Cu(OAc)2 with excess acetic acid to suppress hydrolysis, crystals of the salt should be formed on cooling. |
Yeah, I already looked that up, but the problem is that I don't have that much acetic acid. I'm making 1mol copper acetate, which dissolves in almost
1l of hot water. So if I would want the solution to be let's say 10% acetic acid AFTER reaction is complete, that's almost 100ml pure acetic acid
excess (210ml total) and I don't have that much...
I guess I'll try to dissolve the basic carbonate in multiple portions of 9% vinegar...with a large excess it should work 
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blogfast25
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Nope. Commercial CuO is likely to have been calcined hard. That stuff doesn't usually dissolve into anything but the strongest and most concentrated
acids.
[Edited on 23-2-2013 by blogfast25]
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blogfast25
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| Quote: Originally posted by Nathaniel | I guess I'll try to dissolve the basic carbonate in multiple portions of 9% vinegar...with a large excess it should work
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Or you could try and concentrate the acetic acid a bit. React the vinegar with stoichiometric amounts of slaked lime (Ca(OH)2) to obtain calcium
acetate, by boiling in. Treat the solid with conc. H2SO4 and distill over the acetic acid (ethanoic acid).
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KonkreteRocketry
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| Quote: Originally posted by blogfast25 |
Nope. Commercial CuO is likely to have been calcined hard. That stuff doesn't usually dissolve into anything but the strongest and most concentrated
acids.
[Edited on 23-2-2013 by blogfast25] |
I got my CuO by decomp from Copper IInitrate, so that shall work ?
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blogfast25
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If you didn't calcine it to death then probably, yes.
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KonkreteRocketry
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what do u mean by calcine it to death ? would heating CuO do something to it ? I heat it in a 200 degree alcohol flame.
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woelen
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Calcining to death of an oxide is heating very strongly (1000 C or so) for an extended period of time. Many oxides become very inert when this is
done. E.g. they cannot be dissolved in acids anymore when they are calcined.
Some noteworthy examples are:
Cr2O3
Al2O3
Fe2O3
Co3O4
TiO2
These oxides do not dissolve in strong acids, not even boiling hot 35% hydrochloric acid, 65% nitric acid or pure H2SO4 are capable of dissolving
these oxides after calcining. This change of property is due to formation of a more compact crystalline form, which slowly occurs at very high
temperatures, even well below the real melting point of the oxides. An inert calcined oxide usually requires workup with molten NaOH or molten NaHSO4
in order to dissolve it. This requires working with melts at several hundreds of degrees centigrade.
I think that CuO, on the other hand, remains fairly reactive, even after strong calcining.
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KonkreteRocketry
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| Quote: Originally posted by woelen | Calcining to death of an oxide is heating very strongly (1000 C or so) for an extended period of time. Many oxides become very inert when this is
done. E.g. they cannot be dissolved in acids anymore when they are calcined.
Some noteworthy examples are:
Cr2O3
Al2O3
Fe2O3
Co3O4
TiO2
These oxides do not dissolve in strong acids, not even boiling hot 35% hydrochloric acid, 65% nitric acid or pure H2SO4 are capable of dissolving
these oxides after calcining. This change of property is due to formation of a more compact crystalline form, which slowly occurs at very high
temperatures, even well below the real melting point of the oxides. An inert calcined oxide usually requires workup with molten NaOH or molten NaHSO4
in order to dissolve it. This requires working with melts at several hundreds of degrees centigrade.
I think that CuO, on the other hand, remains fairly reactive, even after strong calcining. |
So yeah mine was not even over 200 degree, so my CuO shall work with household vinegar right ? My vinegar is 6% Acetic acid. Shall i heat it while i
put the CuO insdie or its ok if i just put CuO.
And in the end i shall get Copper acetate right ?
Cool.
I have some K2S and idk what to do with it,
Do u know what K2S thermal decmpose into ?
and when i put K2S with water, this will happen right ?
K2S + H2O = KOH + KSH
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blogfast25
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Maybe. Vinegar is a weak solution (5 - 6%) of a weak acid, i.e. acetic acid (ethanoic acid). Only a small part of the ethanoic acid is actually
dissociated via:
HOAc(aq) + H2O(l) < === > H3O+(aq) + OAC-(aq)... (1)
And it's the oxonium ions (H3O+) that are supposed to do the work:
CuO(s) +2 H3O+(aq) === > Cu2+(aq) + 3 H2O(l)...(2)
Since as the concentration of oxonium ions in a commercial vinegar is really small, the reaction rate for the second reaction is small, compared to
when you use a strong acid like HCl or H2SO4 where the dissociation step (1) is almost 100 %.
What's more, even if your CuO dissolved completely in your vinegar you still only have a very weak solution of copper (II) acetate, from (2).
The proof is in the trying, of course...
[Edited on 24-2-2013 by blogfast25]
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kingkey24
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why not just react CuCO3 with acetic acid, then boil down the solution and place in a desiccator for crystallization?
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AJKOER
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| Quote: Originally posted by panziandi | | If you keep copper wire in acetic acid you will get a blue-green solution of copper (II) acetate. I have done this with vinegar and 33% acid, but it
is SLOW. If you heat copper wire in a blue flame you can oxidise it to copper (II) oxide which will dissolve more easily in the acetic acid. Copper
(II) oxide and copper (II) carbonate (basic) can be had cheaply and can be dissolved in the acetic acid quickly to yield the acetate. Copper (II)
acetate are beautiful crystals, enjoy |
I tried the following reaction recently and was surprised. I started by adding dilute aqueous household ammonia to some copper coated pennies (US
currency). I was not expecting much, and 'SLOW' is too rapid an adverb for what does occur.
However, per this not too dated 1962 paper "Kinetics and Mechanism of Copper Dissolution In Aqueous Ammonia" fully available after signing on to ones
Facebook account at http://www.academia.edu/292096/Kinetics_and_Mechanism_of_Cop... the author cites a rate for Cu dissolution as a function of available O2 and NH3.
Some of the underlying reactions include:
2 Cu + 4 NH3 + 1/2 O2 + H2O --> 2 [Cu(NH3)2]OH
2 [Cu(NH3)2]OH + 4 NH3 (aq) + 1/2 O2 + H2O --> 2 [Cu(NH3)4](OH)2
Cu + [Cu(NH3)4](OH)2 <---> 2 [Cu(NH3)2]OH
Side reaction:
2 NH3 (aq) + 3 O2 + [Cu(NH3)4](OH)2 --> [Cu(NH3)4](NO2)2 + 4 H2O
and I suspect further oxidation to Copper ammonium nitrate as well.
As a result, I added some dilute H2O2 to add oxygen to speed up the reaction.
To my surprise, the pennies became readily covered with O2 in agreement with a cathodic reduction reaction of oxygen at the copper's surface per the
electrochemical dissolution model. The reaction is also apparently exothermic as the solutions became warmer. Within an hour, a dark blue was
apparent. In 8 hours, a different lighter shade of blue was apparent that is characteristic of the usual cupric salts. Expected products could include
tetraamminediaquacopper(II) dihydroxide, [Cu(NH3)4(H2O)2](OH)2, as well a monohydroxide, tetraamminecopper(II) nitrite and also the nitrate. Note, the
reaction produced more gas than I suspected (do not used sealed vessels) perhaps due to the formation of both O2 and N2 (via some nitrite formation
and decomposition reaction).
Conclusion, if you wish to form a cupric salt, this appears to be a fairly quick and inexpensive procedure using dilute ammonia, 3% H2O2 and copper.
Caution: The presence of Copper Ammonium nitrite and/or Ammonium nitrite may present a potential spontaneous nitrogen gas decomposition issue, which
are more likely in slightly acidic or concentrated solutions. I would also be concerned on heating an acidified form of the solution just prepared
due to known stability issues with hot aqueous NH4NO3 in the presence of metallic impurities (including Copper, Tin and Nickel see http://www.google.com/url?sa=t&rct=j&q=ammonium%20ni... ).
--------------------------------------------------
Now, having forming one or more cupric salts, to prepare Cupric acetate, just add NaHCO3 and filter out the CuCO3. Add Acetic acid to the washed
precipitate to form Cupric acetate.
--------------------------------------------------
[EDIT] Here is a less authoritative 2011 study ("Copper-Mediated Non-Enzymatic Formation of Nitrite from Ammonia and Hydrogen peroxide at Alkaline pH"
) that is pertinent relating to nitrite formation noted above (please see
http://www.google.com/url?sa=t&rct=j&q=reaction%20of%20nh3%2Ch2o2%20and%20cu&source=web&cd=4&ved=0CDwQFjAD&url=http%3A%2F%2Fsp
hinxsai.com%2Fvol3.no2%2Fchem%2Fchempdf%2FCT%3D23(646-656)AJ11.pdf&ei=iS-mUfCNN4nr0gGYw4D4BA&usg=AFQjCNFaObAi5_3NNOdt8e1DiRoiHzg9bg&bvm=bv
.47008514,d.dmQ ). To quote:
"Hydrogen peroxide with lowest recorded redox
potential of - 0.68 V compared to that of Cu++ / Cu+, +
0.15 V15 acts as a strong reducing agent particularly in
presence of hydroxide ions [13], [18] to donate electrons to
copper (II) forming copper (I) oxide,
H2O2 + 2 OH- → 2 H2O + O2 + 2 e- (1)
2 Cu++ + 2 e- + H2O2 → Cu2 O + H2O (2)
Reddish-yellow cuprous oxide is rendered colorless in
presence of sufficient ammonia to form
diamminecopper (I) [15],
Cu2 O + 2 NH4OH → 2 [Cu (NH3)2] OH + H2O (3)
[ not balanced, corrected per ajkoer:
Cu2O + 4 NH3 + H2O → 2 [Cu(NH3)2]OH (3)]
Diamminecopper (I), generated from reduction of
copper (II) or added exogenously facilitates oxidation
of ammonia, a reducing agent [14], by hydrogen
peroxide,
...[Catalyst].....Cu (NH3)2]OH.........................
NH3 + 3 H2O2 -----------------> HNO2 + 5 H2O (4)
[ not balanced, corrected by ajkoer:
NH3 + 3 H2O2 -----------------> HNO2 + 4 H2O (4)]
Further studies are required to elucidate the actual role
of diamminecopper (I) in the reaction; whether it is
converted to tetramminecopper (II), or undergoes a
reversible changes during the process."
With additional ammonia, the reaction with nitrous acid proceeds as follows:
HNO2 + NH3•H2O --> NH4NO2 + H2O
Interesting observations by the author includes "The reaction is mediated by copper (II) as it fails to occur in absence of copper", and that the best
order of addition of reactants is Cu then aqueous NH3 and finally H2O2. The author also notes the need for excess ammonia, to quote: "as it is needed
to maintain: (i) solubility of copper; (ii) optimal alkalinity for expression of reducing potential of hydrogen
peroxide; (iii) adequate concentration of free ammonia; and (iv) conversion of nitrous acid to ammonium nitrite."
[Edited on 30-5-2013 by AJKOER]
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subsecret
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You can damage the fauna in your septic tank if you throw large quantities of copper salts down the drain. Same goes for sodium hypochlorite.
Fear is what you get when caution wasn't enough.
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thebean
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Copper Acetate Synthesis
I've never seen this synthesis before and I attempted the other day. I believe it worked and thought I'd share my findings. I'm going to transcribe
the notes.
31.2 grams of CuSO4·5H2O are weighed and dissolved in minimal water with ice. 9.99 grams of NaOH are added over the course of a few minutes. The
solution was then filtered. Cu(OH)2 is an is a blue insoluble compound. After filtration the Cu(OH)2 was added to a beaker and acetic acid (6%
vinegar) was added until no reaction occurred. The solution goes a dark green and then after sitting for a while turns blue which indicates Cu2+ in
solution. This is then boiled down to a solid.
You could use higher concentration acetic but I don't have a lot of glacial acetic acid and didn't want to use it on this synthesis so I used a lot of
vinegar instead of a lot less glacial. The acetic acid strips the hydroxide off of the Cu(OH)2 and gives you Cu2+ in solution and water. Once you boil
the solution down the Cu2+ ion attaches itself to the C2H3O2- ion and forms copper (II) acetate. I've seen a lot on vinegar, pennies and hydrogen
peroxide being boiled but this gives you impure copper acetate and a mediocre yield.
"You need a little bit of insanity to do great things."
-Henry Rollins
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violet sin
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I just made a nice pile of copper acetate by taking all the sludge from copper hydroxide electrolysis experiments, and leaving it in vinegar. After
reducing water content a few times, then adding more vinegar to get all the sludge dissolved, I filtered and crystallized it. Apparently I still had
some sulfate in there from the Epsom salt electrolyte. This was evident from crust on top of the sol. When viewed with a back light. ~+80% teal
crystals with a few much more blue rhombus patterns every so often. But I ended up with a lill more than two pill bottles worth of nice sized
crystals. Rock salt-ish size, maybe a bit smaller. Fun and at least useful in the future as opposed to mixed oxides and hydroxide. I used a coffee
carafe on low heat with a computer fan blowing on the surface took a while, but I didn't end up with more hydroxide/carbonate by decomp. The last
lill bit had some of the magnesium salts in it, but the first 4 or so crystallizations never went to dry, so they were much cleaner.
I was hoping to leave some to crystallize to much larger size like in the 3rd page of pretty pictures 2 thread.
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cyanureeves
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nice!i will give this a try for sure thebean because it took me weeks to get just a few tiny crystals of copper acetate using peroxide and vinegar and
natural summer heat.
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thebean
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Violet Sin, you could make larger crystals by putting the solution of copper acetate in a more moist environment because this slows down the
crystallization process and gives you larger crystals. I would also make sure that the solution is as saturated as possible and then add a seed.
"You need a little bit of insanity to do great things."
-Henry Rollins
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PHILOU Zrealone
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Simply leave metallic Cu wire or pieces into a closed jam glass jar with white vinegar up to 1/4 height of the recipient.
The metallic Cu wire have to come from the liquid phase to the aerian phase!
The air is important above the liquid!
Swirl the closed jar from time to time and open from time to time to allow fresh air to come inside.
The oxygen is consumed in the process so the jar will be in slight depression.
A nice blue colour will appear in a day, then green colour will darken from day to day until you get concentrated Cu acetate.
Cu is oxydisable by air and forms a layer of CuO and Cu(OH)2 what is dissolved by the vapours of water and acetic acid. The naked metal is then
further oxydised and the cycle can continue...theorically until all Cu has dissolved (if enough reactants are present).
Same happens with metallic Cu and concentrated NH4OH. You can form concentrated Cu(NH3)4(OH)2 solution and crystallization.
The Eurocents (1, 2 and 5) Cu layer (Cu plating) can be completely dissolved that way leaving the silvery steel core of the coins.
PH Z (PHILOU Zrealone)
"Physic is all what never works; Chemistry is all what stinks and explodes!"-"Life that deadly disease, sexually transmitted."(W.Allen)
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Metacelsus
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I like to use basic copper carbonate (from copper sulfate and sodium carbonate) to react with acetic acid. Sodium carbonate is easier to obtain for me
than sodium hydroxide.
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bfesser
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Threads Merged
14-11-2013 at 06:29 |
bfesser
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Thread Pruned
14-11-2013 at 19:38 |
cjevancich
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I just wanted to add, since I didn't notice it already mentioned, that Copper acetate has the (annoying) property of decomposing under heat.
Therefore, if you synthesize it, in a large, aqueous solution, you will have to slowly evaporate it. You cannot get away with boiling off the water.
At best, you can keep the water warm with something like a crock pot on a low setting. If you try to boil it, the Copper acetate will decompose into
Copper hydroxide and you'll lose acetic acid.
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Arthur Dent
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I like the Copper Carbonate method, even if it's a bit labour intensive.
After dissolving pure copper wire in HCl (which at first, looks like a brownish, turbid solution until it oxidizes into a bright emerald green
solution), I neutralize it by adding sodium bicarbonate, which produces an insoluble pale greenish blue Copper Carbonate precipitate and a Sodium
Chloride solution. It's a bitch to filter so my suggestion is to let the precipitate settle, decant as much liquid as you can, add more distilled
water and repeat the procedure a few more times, finally filter off the precipitate with ample quantities of water. The resulting copper carbonate mud
can then be poured on a plate of glass to dry, then scraped off. The other annoyance is that if you don't do several water/decant cycles, then a
thorough filtering, there will still be a lot of Sodium Chloride/Bicarbonate impurities in the Copper Carbonate.
The resulting light greenish blue dust readily dissolves in Acetic Acid and turns into a clear blue solution. Gently heat the solution to evaporate
the water and voilà! The resulting crystals will look like dark blue sand. I tried this experiment with glacial acetic acid, never tried with plain
vinegar (but it should work equally, although there would be a lot of water to evaporate).
Robert
Robert
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DraconicAcid
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Copper(II) carbonate (actually a basic carbonate) is a goopy mess, but it becomes much more tractable if you wash it with acetone. Same with
copper(II) oxide.
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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BobD1001
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I made some relatively pure Copper Acetate recently. Simply dissolved some pure copper wire in dilute glacial acetic acid, with added H2O2 to force
the reaction. Let it dissolve for several days until no more reaction was noted. Then I simply vacuum dessicated it,and got some nice dark copper
acetate crystals.
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Ascaridole
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Many years ago I remember making copper (II) acetate from copper (II) oxide formed via thermal decomposition of copper (II) carbonate in an oven. The
still hot copper oxide was then added to concentrated acetic acid and crystals began to form quite rapidly. The acetic acid was not glacial as I
distilled it from sodium acetate and sodium bisulfate that was not perfectly dry. Most likely I formed the monohydrate.
The stupid part was I attempted to purify the already beautiful crystals in a dilute acetic acid solution and over heated the solution and began
forming basic copper salts. Never did recover my beautiful crystals. Lesson learned, copper likes to form basic salts.
Perhaps a repeat of the experiment is in order...
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