# How may copper acetate ligands be manipulated to change colors?

I have a solution of copper acetate and I would like to play around with the ligands to get different colors.

Background: The copper acetate was made through mixing vinegar (5% acetic acid), NaCl, and C$_{\textrm{(s)}}$. The deep blue-colored copper acetate spontaneously formed during a month in my dark storage room.

 I neglected to mention that the experiment simultaneously produced a 0.5 inch deposit of what appears to be Copper Carbonate or Verdigris on the bottom of the 1 Liter beaker. Also, I have let the copper acetate solution evaporate for several years now. The former 1L is now 1/2 liter and has begun precipitating crystals (like the ones on the wiki page). Fun fact: During the Renaissance, glacial acetic acid was made by dry distilling metal acetates and primarily copper(II) acetate.[/edit]

Question(s):

What easily obtained household chemicals may be mixed with samples of the copper acetate to change the ligands attached to the copper and thereby alter the color?

Will heating or cooling the solution change the color and/or ligands?

• That sounds like a cool project! Do you have enough solution for some trial and error? All manner of organic oxyanions bind to copper. You might try EDTA, which can be found in a number of beverages, although I don't know what effect that would have on color. – Pat Apr 27 '12 at 4:28
• Common off-the-shelf eye drops are one of the best sources of EDTA if I recall correctly. Citric acid might be worth a try, instead, though, since it's cheaper and will probably behave pretty similarly. – Aesin Apr 27 '12 at 19:05
• Copper arsenate is a wonderful green… but I'm not sure you could consider arnite sources as “home chemicals” :) – F'x Apr 28 '12 at 3:40
• @Pat I have around 1/2 Liter of deep blue copper acetate. – Dale Apr 29 '12 at 23:02
• I have found an interesting article on making Cu2O (Cuprite crystals) from Copper acetate titled, "Seed-mediated synthesis of polyhedral 50-facet Cu2O architectures." available here: rsc.org/suppdata/ce/c1/c1ce05243h/c1ce05243h.pdf I wonder if that polymorph of Cuprite is also a superconductor. – Dale May 10 '12 at 20:54

I may have accidentally created copper cloride in addition to the copper acetate. I decided to just let the solution (made form Copper, vinegar, and salt) evaporate with a small fan speeding the evaporation for the last few weeks. The tallest copper acetate crystals started breaking the surface and then the surface of the liquid started to form a skin (pic 1) so I poured off the now green liquid into another jar. Copper acetate is blue, which is why the green solution seems to suggest that there is copper cloride.

Evaporation! Note the blue-green layer of copper carbonate below the dark copper acetate crystals.

Top View

The green liquid on the left is what was left from evaporation (and I think is copper chloride) - next to another jar of copper acetate.

The copper acetate crystals were covered in a thin layer of copper carbonate because last week I disturbed the solution and stirred up a cloud of copper carbonate accidantally. I rinsed them with vinegar and they cleaned nicely.

The largest crystal was just over 1 cm wide. These crystals were grown from a solution that started evaporating 5 years ago.

The copper carbonate goo in the bottom needed a use so I added more vinegar to it and it formed the copper acetate solution three pictures up (next to the green copper chloride).

Copper (II) is almost always blue, so I'm not sure how much luck you'll have trying to change the color. CuCl$_2$ is one of the few solutions I can think of off the top of my head that isn't blue - it's more greenish-blue.

CuCl$_2$ can be made via electrolysis of NaCl and copper plates. You could also make CuCl$_2$ if you add hydrochloric acid and some NaCl.

Here are two more blue shades:

• the bright blue of the hexa-hydrated $\ce{[Cu^{II}(H2O)6]^2+}$ (dissolve $\ce{CuSO4}$, maybe not household chemical, but used e.g. as fungicide in farming incuding organic farming),

• if you drop ammonia into this solution, you'll first see a white (blueish with the remaining hexahydrate) precipitate or clouds of the hydroxide

• which dissolves with more ammonia into the deep blue tetra ammine complex $\ce{[Cu^{II}(NH3)4(H2O)2]^2+}$

Old question but interesting topic. I don't intend to reply, but rather to add some relevant info. There is no "one" copper acetate. There are many, some neutral, some basic. They all follow the rule $$\ce{[Cu(CH3COO)2]_x·[Cu(OH)2]_y·nH2O}$$ in which x ≥ 1, y ≥ 0 and n ≥ 0 (i.e., with ratios x:y:n, being basic if y ≠ 0).

Neutral copper acetates exist in different hydration levels: $\ce{{[1:0:0] }Cu(CH3COO)2}$ or $\ce{{[1:0:1] }Cu(CH3COO)2·H2O}$ or $\ce{{[1:0:2] }Cu(CH3COO)2·2H2O}$, being the monohydrated one the most stable and commonly found/produced.

The known types of basic copper acetates (Kühn 1993) are 2:1:5 (blue), 1:1:5 (blue), 1:2:0 (blue, unstable) and 1:3:2 (green), but eventually other formulas and/or structures are mentioned (e.g., $\ce{Cu4(OH)(CH3COO)7·2H2O}$, allegdly isostructural with $\ce{Co4(OH)(CH3COO)7·2H2O}$, according to López-Delgado et al. 1998 & 2001).

Scott (2002) has some neat "concentration maps" (Pourbaix diagrams) which show the products obtained depending on (reagents and respective) concentrations.

References:

• Kühn, Hermann (1993): Verdigris and Copper Resinate. In Ashok Roy (Ed.): Artists' Pigments. A Handbook of Their History and Characteristics, Volume 2. Washington, London: National Gallery of Art; Archetype, pp. 131–158.

• López-Delgado, Aurora; Cano, Esmeralda; Bastidas Rull, José Maria; López, Félix A. (1998): A Laboratory Study of the Effect of Acetic Acid Vapor on Atmospheric Copper Corrosion. In J. Electrochem. Soc. 145 (12), p. 4140-4147. DOI: 10.1149/1.1838928.

• López-Delgado, Aurora; Cano, Esmeralda; Bastidas Rull, José Maria; López, Félix A. (2001): A comparative study on copper corrosion originated by formic and acetic acid vapours. In Journal of Materials Science 36, pp. 5203–5211. DOI: 10.1023/A:1012497912875.

• Scott, David A. (2002): Copper and Bronze in Art. Corrosion, Colorants, Conservation. Los Angeles: Getty Conservation Institute. Partially available online at https://books.google.com/books?id=yQKuSOzkLvcC.