Sciencemadness Discussion Board - Copper acetate problem. Muddy brown solution

Copper acetate problem. Muddy brown solution

Stearic - 9-7-2013 at 16:46

I was synthesizing copper acetate from copper metal in boiling 50/50 3%h2o2 and 5%AcOH. The reaction proceeded as planned forming coppers characteristic blue tinge, until about a minute in. It then began evolving a brown foam rapidly. Once I calmed the foam down I was left with a muddy brown solution. Could the h2o2 have formed coppers oxide?

[Edited on 10-7-2013 by Stearic]

CuO catalyzed hydrogen peroxide decomp.

bfesser - 9-7-2013 at 16:50

Please type formulae properly—it matters.

H2O2

The answer is yes. You formed <a href="http://en.wikipedia.org/wiki/Copper(II)_oxide" target="_blank">copper(II) oxide</a> <img src="../scipics/_wiki.png" />, which catalytically decomposes hydrogen peroxide—esp. ca. 100 °C.

2 H2O2(aq) —cat. CuO→ 2 H2O + O2↑

References:<ul type="circle"><li>Fred Senese. Will hydrogen peroxide blacken copper?. Frostburg State University General Chemistry Online!. http://antoine.frostburg.edu/chem/senese/101/redox/faq/h2o2-... (accessed July 9, 2013).</li><li>V. Múčka. Decomposition of hydrogen peroxide on copper(II) oxide. Collect. Czech. Chem. Commun. 1976, 41, 1717-1726. DOI: <a href="http://dx.doi.org/10.1135/cccc19761717">10.1135/cccc19761717</a> <img src="../scipics/_ext.png" /></li><li>M.H. Robbins, R.S. Drago. Activation of Hydrogen Peroxide for Oxidation by Copper(II) Complexes. Journal of Catalysis. 1997 170, 295–303.</li>DOI: <a href="http://dx.doi.org/10.1006/jcat.1997.1754" target="_blank">10.1006/jcat.1997.1754</a> <img src="../scipics/_ext.png" /></li></ul>

[Edited on 7/11/13 by bfesser]

Stearic - 9-7-2013 at 16:55

I apologize I'm usually more careful about capitalization. I appreciate the help.

Shows a graph on decomposition rates dependent on catalyst. Copper (II) Oxide would most certainly cause decomposition at 100C.
http://www.gcsescience.com/rc13-catalyst-hydrogen-peroxide.h...

[<a href="u2u.php?action=send&username=bfesser">bfesser</a>: merged sequential posts]

[Edited on 7/10/13 by bfesser]

bfesser - 9-7-2013 at 17:08

Stearic, I appreciate your willingness to search for yourself, and your effort in posting (punctuation, full sentences, etc.). It's quite refreshing. One suggestion, however; when you've just made a post, there is an <img src="./images/xpblue/edit.gif" /> button on the upper right, if you need to add or change something. It's preferred over making a sequential post. The "Preview Post" button is also handy. Other than that, you're doing everything right so far. Keep up the good work!

[edit]
Oh, and pictures of colorful copper chemistry are always appreciated!

[Edited on 7/10/13 by bfesser]

AndersHoveland - 10-7-2013 at 00:51

Both copper and manganese(III) can form unusual structured complexes with acetate. These might not be the color you would expect.

Metacelsus - 10-7-2013 at 05:00

A better way to make copper acetate would be to react basic copper carbonate (precipitated from copper sulfate and sodium carbonate) with acetic acid. Also, mixing hydrogen peroxide and acetic acid can form the dangerously unstable peroxyacetic acid; however, this probably won't happen with vinegar and 3% hydrogen peroxide.

AJKOER - 10-7-2013 at 06:27

OK, I have performed this experiment by adding metal copper (actually US pennies with the intent of cleaning) to a mixture of vinegar and 3% H2O2. It never occurred to me to apply heat as I felt the H2O2 would just decompose too quickly. The reaction, as I understand it, is that exposed Cu reacts with the H2O2 forming CuO, which is more readily attacked/dissolved by the weak acetic acid exposing new copper, and the reaction cycle repeats. As such, it may be wise in this slow preparation of Copper(II) acetate to periodically add more H2O2 in a sealed vessel to avoid CO2 exposure. To quote from Wikipedia (http://en.wikipedia.org/wiki/Basic_copper_carbonate ):

"Copper in moist air slowly acquires a dull green coating because its top layer has oxidised with the air. Some architects use this material on rooftops for this interesting colour. The green material is a 1:1 mole mixture of Cu(OH)2 and CuCO3:[1]

2 Cu (s) + H2O (g) + CO2 + O2 → Cu(OH)2 + CuCO3 (s) "

In a day, without heating, and sealed from air, the solution will form the characteristic color of a copper salt. More precisely, Cu2(OAc)4(H2O)2 is bluish-green, whereas anhydrous Cu(OAc)2 is a dark green crystalline solid. If using a copper plated source (a post 1982 US Penny, for example, is at least 95% Zn, but pure Cu plate, see http://en.wikipedia.org/wiki/US_penny ), do not dissolve too much of the penny as exposed Zn could enter the reaction mix and displace Cu. As Zn acetate is colorless, you will clearly notice the color change in the solution. To quote Wikipedia (link: http://en.wikipedia.org/wiki/Zinc_acetate ):

"Zinc acetate is the chemical compound with the formula Zn(O2CCH3)2, which commonly occurs as a dihydrate Zn(O2CCH3)2(H2O)2. Both the hydrate and the anhydrous forms are colorless solids that are commonly used in chemical synthesis and as dietary supplements. Zinc acetates are prepared by the action of acetic acid on zinc carbonate or zinc metal."

[EDIT] Having said some nice things about Zinc acetate, please note that these comments do not generally apply to Copper acetate. In fact, some history per Wiki, to quote:

"Copper(II) acetate was historically prepared in vineyards, since acetic acid is a byproduct of fermentation. Copper sheets were alternately layered with fermented grape skins and dregs left over from wine production and exposed to air. This would leave a blue substance on the outside of the sheet. This was then scraped off and dissolved in water. The resulting solid was used as a pigment, or combined with arsenic trioxide to form copper acetoarsenite, a powerful insecticide and fungicide called Paris Green or Schweinfurt Green."

Apparently many copper salts are highly toxic to lower organism, including fish, small animals and even fungus, and find limited use as copper pesticides. To quote Wikipedia on Copper pesticides (http://en.wikipedia.org/wiki/Copper_pesticide ):

"A copper pesticide is a copper compound used as a pesticide or fungicide. In the UK the Soil Association (one of the organic certification authorities) permits farmers to use some copper fungicides on organic land used for the production of certified organic crops only if there is a major threat to crops. [1] During the 2008 growing season the compounds were applied to much of the organic potato crop to control potato blight (Phytophthora infestans).[2] The compounds permitted are copper sulfate, copper hydroxide, cuprous oxide, copper oxychloride, copper ammonium carbonate (at a maximum concentration of 25 g/l), and copper oxtanoate. According to the Soil Association the total copper that can be applied to organic land is 6kg/ha/year.[3] This limit is designed so that the amount of copper in the soil does not exceed the limits specified in the Soil Association standards for heavy metals."

Copper ammonium carbonate is particularly nasty, and can be formed quite accidentally. For example, the exposure of a newly installed brass fitting in a fish tank to waste products leading to NH3 dissolving a small amount of Cu, which further combines with CO2 from aeration, or via:

H2N-CO-NH2 (aq) + H2O <--Bacteria--> CO2 (g) + 2 NH3 (aq)

can completely poison a fish tank killing fish and plants (see http://www.thereeftank.com/forums/f267/help-please-holy-crap... and http://koi2day.com/forum/index.php?topic=4193.0;wap2 ). Bottom line, be mindful when employing Copper (or Brass) and when disposing of aqueous Copper salts.
---------------------------------------------------------------------
[EDIT][EDIT] An interesting fact is that by adding a chloride to the acetic acid/H2O2 mix, is that the dissolving/etching of the Copper metal appears to be visibly accelerated. My take on some of the chemistry in our case (see http://www.hkedcity.net/article/project/hkcho/1P.pdf ):

2 Cu(s) + 2 H2O2 + 4 H+ --> 2 [Cu 2+] (aq) + 4 H2O
[Cu 2+] (aq) + Cu (s) <--> 2 [Cu +] (aq)
[Cu +] (ag) + 2 Cl- (aq) --> [CuCl2 -] (aq)
2 [CuCl2 -] (aq) + H2O2 (aq) + [H +] (aq) + 4 Cl- (aq) < --> 2 [CuCl4 2-] (aq) + OH- (aq) + H2O
[CuCl4 2-] (aq) + Cu(s) <--> 2 [CuCl2-] (aq)

where the 3rd equation requires a high chloride concentration and the last equation causes a circling back to the 4th equation. Upon evaporation, the addition of say NaCl to accelerate the reaction, may cause at least a more costly salt separation issue to isolate the pure Copper acetate (as Cu(OAc)2 is soluble in alcohol whereas NaCl is not).

[Edited on 11-7-2013 by AJKOER]

[Edited on 11-7-2013 by AJKOER]

bfesser - 10-7-2013 at 17:03

http://www.sciencemadness.org/talk/viewthread.php?tid=11144#...

AJKOER - 11-7-2013 at 10:24

Quote: Originally posted by AJKOER

.... An interesting fact is that by adding a chloride to the acetic acid/H2O2 mix, is that the dissolving/etching of the Copper metal appears to be visibly accelerated. My take on some of the chemistry in our case (see http://www.hkedcity.net/article/project/hkcho/1P.pdf ):

2 Cu(s) + 2 H2O2 + 4 H+ --> 2 [Cu 2+] (aq) + 4 H2O
[Cu 2+] (aq) + Cu (s) <--> 2 [Cu +] (aq)
[Cu +] (ag) + 2 Cl- (aq) --> [CuCl2 -] (aq)
2 [CuCl2 -] (aq) + H2O2 (aq) + [H +] (aq) + 4 Cl- (aq) < --> 2 [CuCl4 2-] (aq) + OH- (aq) + H2O
[CuCl4 2-] (aq) + Cu(s) <--> 2 [CuCl2-] (aq)

where the 3rd equation requires a high chloride concentration and the last equation causes a circling back to the 4th equation. Upon evaporation, the addition of say NaCl to accelerate the reaction, may cause at least a more costly salt separation issue to isolate the pure Copper acetate (as Cu(OAc)2 is soluble in alcohol whereas NaCl is not).

In my depiction on the dissolution of copper metal by addition of a chloride to the acetic acid/H2O2 mix as means of accelerating the reaction, I may be a little remiss in not properly describing the process (please correct me, if you disagree with my comment below).

Similar to the discussion at "Kinetics and Mechanism of Copper Dissolution In Aqueous Ammonia", available at http://academia.edu/292096/Kinetics_and_Mechanism_of_Copper_... the dissolution mechanism is, in my opinion, best described as also electrochemical in nature. To quote:

"The generally accepted theory on the corrosion of a metal (Evans[18]), is that when a metal comes into contact with an aqueous salt solution to which oxygen is accessible, oxygen takes up electrons at one part of the surface (the cathodic zone) while the metal gives it up at another (the anodic zone). In this way the attack of the metal proceeds at an appreciable rate at room temperature. These principles are well established and they were successfully demonstrated in many cases, e.g. the dissolution of zinc in sodium chloride solution in contact with air, or gold in a cyanide solution saturated with air, Thompson [19]"

Here are some of the actual equations cited by the author occurring in the overall electrochemical reaction:

1. 1/2 O2 + H2O + 2 e- ---> 2 OH- (Cathodic reduction of O2 at surface of the Copper)

2. Cu + 4 NH3 ---> [Cu(NH3)4]2+ + 2 e- (Anodic dissolution of Cu by a complexing agent)

Overall:

Cu + 4 NH3 + 1/2 O2 + H2O ---> [Cu(NH3)4]2+ + 2 OH-
-------------------------------------------

This net electrochemical reaction is superimposed onto a cupric-cuprous equilibrium (also consuming copper) which can be described as occurring as follows:

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

The similarity to the copper/acetic acid/H2O2 reaction sequence I cited above is apparent. Note also, that the cathodic reaction is a diffusion controlled process (wherein the O2 is consumed upon reaching the Cu surface) while the anodic reaction is chemically controlled.

[Edited on 11-7-2013 by AJKOER]

bfesser - 11-7-2013 at 18:39

Quote: Originally posted by AJKOER

2 Cu(s) + 2 H2O2(aq) + 4 H+(aq) → 2 Cu2+(aq) + 4 H2O(l)
[corrected equation formatting]

I'd like to point out that the interaction between hydrogen peroxide and copper is much more complex than in this portrayal, as described in the paper(s) I posted earlier in the thread.

Please work on cleaning up the way you write equations on the board, they are very difficult to read. You don't use subscript/superscript, and your application of brackets and physical state designations (s/l/g/aq) seems entirely arbitrary. As an example, there was no need to put Cu2+ in square brackets (as I understand it, this is customarily used to indicate unstable intermediates or inferred species), and the ionic charges must be superscripted to avoid confusion—I do understand that the cupric ion exists as a complex in aqueous solution, but it's standard practice to write Cu2+(aq) and leave it to the reader to have a sufficient level knowledge and understanding. The equations you're posting are essentially unreadable. I, for one, would really appreciate some effort shown on your part in properly formatting your reaction equations.

Finally, I don't see any evidence being provided for what appears to be merely conjecture (the [CuCl42-] in particular). I'm not saying that I doubt it can exist, but where did you get this idea from? I'm working on reviewing the PDF posted, but from what I've seen so far, I'd label it as only loosely scientific and of poor quality and reputability. I also note a disturbing lack of citations in the paper. Unless I somehow—and I truly hope this is the case—overlooked them. There were none (ZERO)! . . .

<center><img src="../scipics/_warn.png" /> Warning: What follows is off-topic and may inadvertently offend some readers. <img src="../scipics/_warn.png" /></center><hr width="800" /> . . . Not to discriminate, but this is typical of what I've witnessed from 'academia' and 'science' in China as of late. There seems to be no emphasis on proper scientific procedures, poor quality of literature research, and no respect at all for original work—plagiarism is rampant and widely accepted! I don't have them in hand at the moment, but I recall seeing papers out of China in which the original authors names have merely been replaced and the papers submitted to journals and re-published as original work. There have even been entire doctoral theses plagiarized and accepted! This is quite upsetting and saddening to me.

Yes, I saw that the paper is only from a secondary school 'Science Olympiad' program… but come on! I guess it's easiest to impression pupils while they're young that plagiarism is acceptable, as long as they come out ahead in life. Anyway, I'll see if I can find those papers and related news stories to post in another thread, if anyone cares to see them.

[edit]
I just realized that you may have been trying to show, with brackets, overall ionic charge of the complex ions. Is this the case?

[Edited on 7/12/13 by bfesser]