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ficolas
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[*] posted on 14-5-2016 at 11:27
Green "copper acetate" solution


I created some copper acetate to then crystalyce into beautiful crystals (hopefully)
I added some copper wire to a acetic acid and hydrogen peroxide solution, and after some time, I got a beautiful blue solution, that I then evaporated off, to get some little greenish-blue crystals.

After some time, I went to prepare a solution with those crystals, I heated some water, and started adding the copper acetate, but to my surprise, its seems to be very unsoluble, and the solution its green. Why can this be? Did I get some weird impurities in my copper acetate? Did my copper acetate decompose into something? (I had it in a transparent bottle, since I didnt read anywhere that copper acetate can decompose from the action of light or anything like that) Or did my copper acetate react with some chlorine compound in the water to form copper chloride? Because I made the solution with tap-water, since I was just testing how easily it disolves in a very small scale, I will use distilled water for the recrystalization and the solution for the big crystal.

[Edited on 14-5-2016 by ficolas]
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Boffis
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[*] posted on 14-5-2016 at 13:23


It is a curious feature that some transition metal acetate noteably copper and nickel dissolve very slowly and if you heat the solution the salts turn into insoluble basic salts. So be patient and use cold water with a little acetic acid in!
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[*] posted on 14-5-2016 at 13:30


Boffis is right. Hydrolysis of acetate anion is present as it is a base. Crystallize it out of acetic acid solution. It will work. :cool:



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AJKOER
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[*] posted on 16-5-2016 at 06:45


Cu(ll) can be reduced to Cu(l) with light or the presence of an organic reducer, like ascorbate, citrate,..

Add O2 which will directly react with Cu(l) and any available H+ source including from the dissociation of water. This leaves OH- which is the source of hydroxyl ion for the formation of basic salt. As an example of this oxygen path to a basic salt see Wikipedia on dicopper chloride trihydroxide at https://en.m.wikipedia.org/wiki/Dicopper_chloride_trihydroxi... .

Note, the direct reaction between Cu(l) or Fe(ll) with molecular oxygen is rare among transition metals.

The copper based reaction is given by

4 Cu+ + 4H+ + O2→ 4 Cu2+ + 2H2O

Source: https://www.researchgate.net/publication/262451840_Review_of...

Similarly, for the iron based reaction, to quote:

"This reaction occurs spontaneously in 2 steps. The first reaction is:

2 Fe(s) + O2(g) + 2 H2O(l) = 2 Fe(OH)2(s)

The second reaction is:

4 Fe(OH)2(s) + O2(g) = 2 H2O(l) + 2 Fe2O3 · H2O(s) "

Link: https://www.google.com/url?sa=t&source=web&rct=j&...

Or, per another source, expressed as:

4 Fe2+(aq) + O2(g) + 4 H2O(l) + 2 xH2O(l) → 2Fe2O3•xH2O(s) + 8H+(aq)

which is more interesting as ferrous/oxygen is here a provider of H+ and thus could complement the action of Cu(l) with O2 which consumes it. This could account, in part, for the syngeristic benefits of mixed iron/copper salts in Fenton-like advanced oxidation processes.

[Edited on 16-5-2016 by AJKOER]
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Boffis
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[*] posted on 16-5-2016 at 11:05


@Ajoker; what has all that got to do with hydrolysis of transition metal acetates salts in solution?
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AJKOER
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[*] posted on 16-5-2016 at 12:56


My supposition is that the so called hydrolysis of Cu(ll) could really be the result of a chain reaction. It starts with the formation of some Cu(l) (perhaps from the action of light on Cu(ll), or the action of Cu on Cu(ll) ), its subsequence reaction with O2 and H+ (leaving OH- if the hydronium ion is sourced from water), which further can result in the formation of a basic copper salt.

A variation of the path sourcing Cu(l) from the action of Cu and Cu(ll) in the presence of aqueous NaCl and dissolve oxygen is discussed in the Wikipedia reference and is the basis of a commercial patent for dicopper chloride trihydroxide.

If you can offer an alternate path to the product of the hydrolysis resulting in a basic copper salt, I would be interested.

[Edited on 17-5-2016 by AJKOER]
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Boffis
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[*] posted on 16-5-2016 at 17:21


@Ajoker, I'd hate to have you design a nutcracker. You always manage to come up with the most obscenely complex route to anything.

This reaction is a simple hydrolysis reaction; water produces an equilibrium concentration of OH- ions and H+ ions (actually H3O+). Acetic acid is a weak acid so when copper acetate ionises in water it produce Cu2+ and C2H3O2- ions the latter compete for the H+ ions leaving behind the OH- ions which quickly render the Cu2+ less soluble by generating hydroxide bearing compound ions and free acetic acid. This is why the addition of acetic acid reduces hydroylsis (check out Chatelier's principle). The accumalation of acetic acid eventually reaches a level that prevents further uptake of H+ ions by acetate ions and so OH- generation; equilibrim is reached.

This reaction appears to be slow at room temperature. This may be because copper acetate appears to act as a complex compound in which the copper is coordinated with acetic acid rather than with water (as in say copper sulphate hydrate) and warming may simply accelerate ion exchange.
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[*] posted on 17-5-2016 at 04:19


Here is a good description of the copper salts in question, see https://books.google.com/books?id=Zk0z22smWUoC&pg=PA80&a... and there usual modes of preparation including my esoteric O2 path for the basic salt.

I am not entirely convinced by your discussion presenting half of the bufferring reaction. This combined with it not being cited as a path to the basic salt makes it a weaker argument in my opinion.

Complex path yes, but I was kind enough not to present the competing Fenton (iron rich tap water) and Fenton-like (via cuprous) reactions, which are relevant here providing OH- (and also hydroxyl radicals).

Note, acetate is an electron donor and has been cited in the reduction of the transition metal manganese, for example, see http://www.jove.com/visualize/abstract/24266405/hydrogen-ace... . The presence of a transition metal in one of its lower valence states is a starting condition for the reaction I cited.

Memo: A better reference on the action of O2 on iron rich water is given by:

O2(aq) + 4Fe2+ + 6H2O ↔ 4FeOOH(s) + 8H+

Source: "Air Oxidation of Ferrous Iron in Water" by Ahmet Alıcılar, Göksel Meriç, Fatih Akkurt and Olcay Şendil, link: https://www.google.com/url?sa=t&source=web&rct=j&...

[Edited on 17-5-2016 by AJKOER]
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Boffis
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[*] posted on 18-5-2016 at 06:16


@Ajoker; hydrolysis of metal salts is NOT a redox reaction. You are mixing up ideas. Yes atmospheric oxygen oxidizes Fe2+ to Fe3+ but this is not the cause of hydroylysis but if the acid concentration is sufficiently high no basic salt is preceipitated.

Consider the case of hydrated aluminium chloride, there is no possibility what so ever of a redox reaction and yet on dilution of a solution of aluminium chloride (don't get confused with anhydrous aluminium chloride here we are talking about the hydrated form) it deposits either a hydroxide or a basis chloride salt (from much personal experience). No redox reaction in site!!!
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ficolas
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[*] posted on 18-5-2016 at 08:19


Well firstly, thanks to you all for your answers.
I cant seem to fully understand the reason why adding acetic acid to the solution increases the solubility, if somebody could explain in a way so that somebody without that much chemistry knowledge (The last things I have studied in high school equilibriums, redox reactions, acid/bases, and some other stuff that I learned online), hopefully next year ill start studing a chemistry degree, or chem. E, and i'll start learning more about it :)
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[*] posted on 18-5-2016 at 09:53


Ficolas: The high school chem version is this.

Copper(II) salts hydrolyze in aqueous solution. Cu(2+) + 2 H2O = Cu(OH)2 + 2 H(+)

Adding more acetic acid (a source of hydrogen ion) drives the equilibrium to the left.




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ficolas
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[*] posted on 18-5-2016 at 11:28


But why the reaction driving to the left makes the salt more soluble? acording to what I have learned, having more Cu(2+) ions in solution, would make the salt less soluble, because
Cu(CH3COO)2 -> Cu(2+) + 2(CH3COO-)
Qs = [Cu(2+)][CH3COO-]^2, so Qs would be higher with more Cu(2+) ions, and it would shift drive the reaction towards the left to reach the equilibrium, so I probably got something wrong there.
Also adding more acetic acid, would increase the amount of CH3COO-, and acording to what I have learned, making the salt less soluble

[Edited on 18-5-2016 by ficolas]
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[*] posted on 18-5-2016 at 11:36


Copper(II) acetate is perfectly soluble. It's the copper(II) hydroxide or basic copper(II) acetate (basically a double salt of copper(II) acetate and copper(II) hydroxide) that is your precipitate.



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[*] posted on 21-5-2016 at 08:47


Quote: Originally posted by AJKOER  
...

Note, the direct reaction between Cu(l) or Fe(ll) with molecular oxygen is rare among transition metals.

The copper based reaction is given by

4 Cu+ + 4H+ + O2→ 4 Cu2+ + 2H2O

Source: https://www.researchgate.net/publication/262451840_Review_of...
....


First point, I should have added Cobalt to the transition metal list along with Copper and Iron. See discussion at https://books.google.com/books?id=KhvqCAAAQBAJ&pg=PA302&...

Second, I did more research on the Oxygen/Copper(l) reaction. Here is a reference (see https://www.google.com/url?sa=t&source=web&rct=j&... ) citing the galvanic half cell reactions supporting the contention that Cu(I) is readily oxidized to Cu(II) in the presence of oxygen:

4(Cu+ = Cu2+ + e–) –Eo = –0.153 V

O2 + 4H+ + 4 e– = 2H2O Eo = +1.229 V

Net cell reaction: 4 Cu+ + O2 + 4H+ = 4 Cu2+ + 2 H2O Eo cell = 1.076 V

Here is also a 2013 radical reaction supplement, "Impacts of aerosols on the chemistry of atmospheric trace gases: a case study of peroxides radicals"', by H. Liang1, Z. M. Chen1, D. Huang1, Y. Zhao1 and Z. Y. Li, link: https://www.google.com/url?sa=t&source=web&rct=j&... that will be used to detail a radical pathway, in particular citing reaction numbers:

R24 O2(aq) + Cu+ → Cu2+ + O2− ( k = 4.6xE05 )
R27 O2− + Cu+ + 2 H+ → Cu2+ + H2O2 ( k = 9.4xE09 )
R25 H2O2 + Cu+ → Cu2+ + .OH + OH− ( k= 7.0 xE03 )
R23 .OH + Cu+ → Cu2+ + OH− ( k = 3.0×E09 )

Net radical reaction: O2 + 4 Cu+ + 2 H+ → 4 Cu2+ + 2 OH-

Note: The above implies that the reaction should proceed with less acid with the formation of a basic cupric salt. It is also seemingly (possible pH effects) equivalent to the prior net reaction upon adding 2 H+ to each side of the equation.

Not cited, however, is the following reverse radical reaction, R29, which even has a higher reaction rate k than for R24:

R29 O2− + Cu2+ → Cu+ + O2(g) ( k = 8.0×E09 )

This reversibility of the Oxygen/Copper(l) reaction (which is apparently also the case for another transition metal, Cobalt, for Oxygen/Cobalt(ll) reaction) may account for the use of a hot concentrated brine, in the prior referenced commercial path for basic cupric chloride, to favorably alter reaction conditions.

Related reference, see "Generation of OH initiated by interaction of Fe and Cu with dioxygen; comparison with the Fenton chemistry" available at https://www.google.com/url?sa=t&source=web&rct=j&...

[Edit] An interesting conjecture to explain the apparent reversible of the O2/Cu(l) interaction is that there may be an electrochemical path (the descriptive term metal auto-oxidation sounds galvanic) for the action of oxygen in the presence of Copper ions and an electrolyte for the forward reaction, and a radical reaction based path for the reverse reaction.

[Edited on 21-5-2016 by AJKOER]
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[*] posted on 24-5-2016 at 14:17


Weak base and weak acid. And people somehow get surprised when their salt undergoes hydrolysis...



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[*] posted on 25-5-2016 at 04:19


Quote: Originally posted by ave369  
Weak base and weak acid. And people somehow get surprised when their salt undergoes hydrolysis...


Add that copper is a fairly active transition metal (note, cobalt salts, for those who have some, is even more so), plus chlorinated (?) tap water (free Cl2, HOCl, CO2, dissolved O2, along with iron salts, perhaps a touch of Mn,..), plus possibly some light exposure,..., isn't a recipe for nothing to happen!
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