TheMrbunGee
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Copper (II) acetate in boiling water
Hey.
So I was trying to purify my Copper acetate, I grinded it up and started to add in near boiling water. (98° C)
It bubbled and became insoluble and in color of copper carbonate. Does it decompose in higher temperatures?
I tried to put it in warm water, everything was fine. When I heated the solution at one point - bubbles and copper carbonate color appeared again.
I was not using distilled water, but there just can’t be that much CO3 ions in my water. :?
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DraconicAcid
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If you get it too hot, you'll have basic copper(II) acetate precipitating, as the acetic acid evaporates.
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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TheMrbunGee
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Thank you! 
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AJKOER
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My take on the chemistry is that the presence of oxygen is also a factor.
Apparently, Copper (also Iron and Cobalt) salts in their lower valent states can react directly with dioxygen with a source of H+. The acid and
sunlight in the presence organic reducers (like citrate, ascorbate, acetate,..) can lead to transition metals in their lower valent state. Also, NaCl
can increase the solubility of the cuprous (by forming a complex in the case of copper), to move the reaction along towards the formation of the basic
salt.
For more details including references, see my comments in a recent thread at http://www.sciencemadness.org/talk/viewthread.php?tid=66247#...
Also, see related discussion on the preparation of basic copper chloride which involves hot brine and O2 at https://en.m.wikipedia.org/wiki/Dicopper_chloride_trihydroxi... .
[Edited on 22-9-2016 by AJKOER]
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DraconicAcid
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Quote: Originally posted by AJKOER  | My take on the chemistry is that the presence of oxygen is also a factor.
Apparently, Copper (also Iron and Cobalt) salts in their lower valent states can react directly with dioxygen with a source of H+. The acid and
sunlight in the presence organic reducers (like citrate, ascorbate, acetate,..) can lead to transition metals in their lower valent state. Also, NaCl
can increase the solubility of the cuprous (by forming a complex in the case of copper), to move the reaction along towards the formation of the basic
salt.
For more details including references, see my comments in a recent thread at http://www.sciencemadness.org/talk/viewthread.php?tid=66247#...
Also, see related discussion on the preparation of basic copper chloride which involves hot brine and O2 at https://en.m.wikipedia.org/wiki/Dicopper_chloride_trihydroxi... .
[Edited on 22-9-2016 by AJKOER] |
This isn't a chloride, and acetate is not a significant reducing agent. Redox chemistry is not a factor in this reaction, Joker.
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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AJKOER
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Per my Wikipedia reference on what I contend is a similar reaction scheme occurring with acetate, cited equations to quote:
" CuCl2 + Cu + 2 NaCl → 2 NaCuCl2 (eq.6)
6 NaCuCl2 + 3/2 O2 + H2O → 2 Cu2(OH)3Cl + 2 CuCl2 + 6 NaCl (eq.7) "
where Equation (7) indicates redox chemistry.
So, assuming we can move the cupric into cuprous, a redox reaction could proceed. In the above cited system, the presence of copper metal assisted in
forming cuprous from cupric. Alternately, to quote a source ( https://www.researchgate.net/publication/11374766_Generation...):
"The process is enhanced by contaminating Fe3+ and Cu2+;"
"The addition of Fe2+ and Cu+ (0-20 microM) to KH resulted in a concentration-dependent increase in *OH formation, as measured by the salicylate
method."
where an iron contamination could arise from using tap water (containing some ferrous bicarbonate, for example, and noting in the opening thread, to
quote MrbunGee, "I was not using distilled water, but there just can’t be that much CO3 ions in my water. :?" ).
[Edit] Yet another reference:
Fe2+ + Cu2+ ↔ Fe3+ + Cu+ (coupled redox reaction)
See: https://www.google.com/url?sa=t&source=web&rct=j&...
[Edited on 23-9-2016 by AJKOER]
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nezza
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If MrBunGee is talking about copper(II) acetate there is no redox reaction occurring. It is simply partial hydrolysis to a basic salt and the
precipitation of that salt. At it's simplest the hydrolysis of copper acetate goes as :-
Cu(CH3COO)2 +2H2O <> Cu(OH)2 +2CH3COOH
That is an oversimplifictation and intermediate species no doubt occur.
Excess water with heat will drive the reaction to the right and excess acid will drive it to the left. Note that because acetic acid is a weak acid
there will always be some hydrolysis of copper acetate in solution.
If you're not part of the solution, you're part of the precipitate.
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TheMrbunGee
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| Quote: Originally posted by nezza | If MrBunGee is talking about copper(II) acetate there is no redox reaction occurring. It is simply partial hydrolysis to a basic salt and the
precipitation of that salt. At it's simplest the hydrolysis of copper acetate goes as :-
Cu(CH3COO)2 +2H2O <> Cu(OH)2 +2CH3COOH
That is an oversimplifictation and intermediate species no doubt occur.
Excess water with heat will drive the reaction to the right and excess acid will drive it to the left. Note that because acetic acid is a weak acid
there will always be some hydrolysis of copper acetate in solution. |
That seems right, if 2CH3COOH looks like copper carbonate.  could not find a
picture of it. The bubbles might just be air trapped in the powder..
Bad thing is that I partially wasted 100g of nice acetate crystals! never
mind, I’ll make some more!
Thanks, guys! You are helpful!
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DraconicAcid
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Just add some more acetic acid to it, and you'll get it back.
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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TheMrbunGee
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Yeah, but that’s not the problem, problem is getting from solution to crystals. I produced those in about 7 runs and few weeks of crystalizing.
I over read the post. CH3COOH really does not look like copper carbonate.
I will attach the video of my failure!
https://youtu.be/UrA3zmI2pZM
In the video you can see precipitate, that looks like Cu(OH)2, but there was a lot of precipitate, that looks like copper carbonate. Teal, solid. Not
the blue clot stuff Cu(OH)2 is like.
[Edited on 23-9-2016 by TheMrbunGee]
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AJKOER
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To quote from Wikipedia on one of the commercial processes used to manufacture basic copper chloride (https://en.m.wikipedia.org/wiki/Dicopper_chloride_trihydroxi...):
"Cu2(OH)3Cl can be prepared by air oxidation of Cu(I)Cl in brine solution. The Cu(I)Cl solution is usually made by reduction of CuCl2 solutions over
copper metal. A CuCl2 solution with concentrated brine is contacted with copper metal until the Cu(II) is completely reduced. The resulting Cu(I)Cl is
then heated to 60 ~ 90 °C and aerated to effect the oxidation and hydrolysis. The oxidation reaction can be performed with or without the copper
metal."
With respect to Copper(ll) acetate, only dilute solutions on boiling are said to undergo hydrolysis (see, https://books.google.com/books?pg=RA1-PA65&lpg=RA1-PA65&... ).
As mentioned previously, tap water is rich in transition metals that can fuel a "coupled redox reaction" along with dissolved oxygen. The key phrase
per above is "aerated to effect the oxidation and hydrolysis". As aerating a solution costs money, it is unlikely a commercial preparation of a basic
copper salt would waste $$$ if the role of oxygen was not material.
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DraconicAcid
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| Quote: Originally posted by TheMrbunGee |
In the video you can see precipitate, that looks like Cu(OH)2, but there was a lot of precipitate, that looks like copper carbonate. Teal, solid. Not
the blue clot stuff Cu(OH)2 is like.
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Probably basic copper(II) acetate.
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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AJKOER
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No, it could be basic copper carbonate!
Assuming Cu(2+) behaves similarly to Pb(2+) in the presence of HCO3- noting a comment on the sensitivity of Lead(II) acetate to CO2 (from your ever so
accomodating tap water) by Wikipedia ( https://en.m.wikipedia.org/wiki/Lead_carbonate ) to quote:
"Lead carbonate is manufactured by passing carbon dioxide into a cold dilute solution of lead(II) acetate, or by shaking a suspension of a lead salt
less soluble than the carbonate with ammonium carbonate at a low temperature to avoid formation of basic lead carbonate."
OK, I have referenced possible issues per your use of tap water relating to the presence of transition metals salts (including those of Fe, Cu, Co,
Mn,..), dissolved O2 and now CO2. I should add that the presence of minerals makes the solution more ionic and a better electrolyte for any favorable
electrochemical reactions. The obvious lesson here is to avoid tap water.
[Edited on 23-9-2016 by AJKOER]
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Fidelmios
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I made the same mistake, I assumed it was copper hydroxide. As said already in this thread, add acetic acid and you're good to go.
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Texium
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Thread Moved
26-9-2016 at 08:03 |
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