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Author: Subject: Cu + HCl + H2O2; Solution color
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[*] posted on 25-2-2006 at 17:46
Cu + HCl + H2O2; Solution color


I put some stripped Cu speaker wire into a flask and added 31% HCL then 3% H2O2 to dissove it. It produced a nice emerald green which would turn so dark it appeared black, yet greenish-yellow on the edges of the top of the solution. I periodically added more H2O2 to speed it up, but what I didn't understand is what was going on with each H2O2 addition.

When the solution turned very dark, and I poured a slight amount of H2O2 into it, it would turn light green where the H2O2 hit then dark again. When I poured just the right amount of H2O2 into the solution, the whole thing would quickly turn the nice emerald green, then eventually very dark again. When the Cu was completely dissolved, I added H2O2 until it made that color transition, and it did not revert back to dark again.

It did not seem at all like the 3% H2O2 was diluting it to the light green, but that in the right amount, it would turn the whole solution emerald green. What is happening here?

Another question: I poured a small amount into a beaker and placed it on a hot plate at 54c to get a nice green solid, but when it was dry it was brown. How do I get the nice green? I don't have to evaporate it all at room temp do I?
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[*] posted on 25-2-2006 at 18:08


Copper has three oxidation states, 0, 1 and 2. 0 is of course metal. +2 is emerald green with chloride (like concentrated HCl, but also saturated salt solutions), and blue without chloride (i.e., plain aqueous solution). A quirk of a chloride solution is, the +1 state is soluble, when ordinarily it isn't (CuCl, copper (I) chloride, is insoluble in water for instance). For some reason, it makes a dark green to brown colored solution. Woelen has made posts on this curious subject.

At any rate, the copper metal dissolving in the solution reduces some of the +2 copper (in solution) to +1, in exchange, oxidizing the copper metal from 0 to +1. Since +1 is dark-colored, it uh, makes a dark colored solution...

When you add the H2O2, even weak as it is, it oxidizes the Cu(I) to Cu(II), removing (bleaching, you could say!) the dark colored compounds, leaving an "emerald" copper solution.

I don't know why you got a smutty product on drying. The first batch of CuCl2.2H2O I evaporated partially dehydrated, making a mottled-color product. (Ya, it's very dry in this house when winter comes around!) The second (smaller) batch I did much more carefully and obtained an even, granular, light blue product. Both of them dissolve in water with little turbidity (hydrolysis), go figure!

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[*] posted on 25-2-2006 at 19:04


I'll have to hunt down Woelen's posts. I have seen the cupric/cuprous chloride samples on his site though.

Anyway, I put the rest (~50mL) in a beaker and sub-boiled it down fast. When almost dry there were three colors to it. There was the green Cu(I), the turquoise Cu(II), and that pesky brown that was a very thin layer on some of the top, and a little on the bottom. I carefully scraped off the brown, trashed it, and let the blue/green air dry until a sniff didn't beat up my nose (no more HCl). I poured it onto a paper towel and powdered it. It looks, colorwise, like it's inbetween the (I,II) states, mostly cupric, but not quite as blue as woelen's sample.

Interestingly it seems the brown from the first dried sample might be reverting back to blue or green. I scraped the bottom of the beaker to free the brown powder, and the super thin layer still covering the bottom of the beaker is now very light in color. It's hard to tell the actual color because it almost looks white it is so thin. I'm waiting for the brown powder in the corner of the beaker to make a change...
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[*] posted on 25-2-2006 at 22:58


Anhydrous CuCl2 is brown. Try adding the brown stuff to water and wait and see it turns blue again.



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[*] posted on 26-2-2006 at 01:58


A solution of CuCl (thats Cu(I)) in HCl is colourless. I think the brown/black colour is due to mixed Cu(I) Cu(II) complexes- compounds with the same metal in 2 oxidation states are often very dark colours.

The Cu(I) compounds are not stable in air- they oxidise so, if you want to see the colourless CuCl in HCl solution you need to prepare it without air present and with an excess of some reducing agent that is powerfull enough to get Cu(II) to Cu(I) but not to get Cu(I) to Cu(0),
The easy way is to use copper wire. If you take that black/ brown soln and put it in a stoppered bottle with copper wire+ leave it you get the brown colour to fade to clear.

BTW copper has Cu(III) and Cu(IV) oxidation states too- but they are a bit obscure.
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[*] posted on 26-2-2006 at 02:15


The H<sub>2</sub>O<sub>2</sub> speeds up the reaction by sucking electrons off the copper. It's very good at that in acidic environments: H<sub>2</sub>O<sub>2</sub> + 2H<sup>+</sup> +2e<sup>-</sup> -> 2 H<sub>2</sub>O

the green colour you see is the colour of the tetrachlorocuprate complex, CuCl<sub>4</sub><sup>2-</sup> which is formed in solutions with high Cu<sup>2+</sup> and high Cl<sup>-</sup> concentations.

Copper can be +1 +2 +3 and +4. +2 is the most common form.
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[*] posted on 26-2-2006 at 06:16


I indeed have done a lot of experimenting with copper (I) and copper (II) ions. The combination of copper (I) and copper (II) in the presence of chloride forms deep brown complexes, with mixed oxidation states with average value between +1 and +2. Here follow the links on this subject:

http://woelen.scheikunde.net/science/chem/riddles/copperI+copperII/index.html
http://woelen.scheikunde.net/science/chem/solutions/cu.html

For me, the special phenomena have not completely settled, but I have a reasonable explanation for all phenomena. The precise explanation, with formula's of all involved compounds, I unfortunately cannot give. Some people believe that the brown complex is Cu(CuCl2)Cl or ClCu(μ-Cl)-CuCl. Both proposed complexes have copper in oxidation state +1.5, but their structures are different. With the first one, CuCl2 is proposed as ligand for Cu(+). These, however, are not established facts, these are theories from some people, who investigated these compounds in more detail.

[Edited on 26-2-06 by woelen]




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[*] posted on 26-2-2006 at 12:35


I have found the reaction between the various Cu oxidation states to be of practical use in the etching of circuit boards. [For those not familiar those boards are an insulating material covered with Cu -- a circuit pattern is exposed and developed on the surface and then an etchant solution is used to selectively remove Cu to give a circuit.] I have used a solution of CuCl2 in HCl as an etchant. As it disolves the Cu from the board some Cu(II) is reduced to Cu(I) and the solution darkens and weakens. By bubbling air thru the solution I can maintain a stronger solution with a green color. After the etching is done I just continue to bubble air thru the solution until it is returned to its original bright green color and can be re-used. Much nicer than etching with FeCl3 which can't be so easily rejuevenated.



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[*] posted on 26-2-2006 at 22:32


The brown powder did grab some H2O and is now blue/turquoise and stuck to the side of the beaker. The brown that I scraped off the top of the 50mL batch after drying and placed outside on the dirt, is the same blue/turquoise. Both are more blue than the (I and II combo) I put into a sandwitch bag left open.

Woelen, I never saw that Cu riddle page so thanks for posting the link.

I Put some of the mixture from the baggie into a test tube and added water. It dissolved fast, but the green crystals settled to the bottom and took a few seconds before they were also dissolved.

I then added a NaOH solution in dropwise. The Cu solution turned blue, then reverted back to turquoise. I then added some NH4Cl, but I detected no smell of NH3. I tested the pH, and sure enough, neutralish. I added more NaOH, and it turned blue then reverted again. I did this a few times until the blue color stayed. At that time there was an unmistakable smell of ammonia. I added more NaOH, but no change (except more ammonia smell.)

This Cu is getting more interesting. I can see why woelen spent so much time working with it.
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[*] posted on 26-2-2006 at 22:38


NeutralIon, how does adding CuCl2 to the HCl affect the etching? I always wanted to try something else for etching beside FeCl3, but I have a gallon of it, so I never got aroung to trying anything else out. I ditched electronics though for the most part, so now I etch a pc board every few months or so.
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[*] posted on 26-2-2006 at 23:16


CuCl2 in HCl can etch copper due to this reaction:

CuCl2 + 2HCl <--> [CuCl4]2- + H+

[CuCl4]2- + Cu <--> 2[CuCl2]- This complex may react with CuCl2 to form the mysterious dark colored product
2[CuCl2]- + O2 + 4HCl <--> [CuCl4]2- +H2O

More info on http://members.optusnet.com.au/~eseychell/PCB/etching_CuCl/

[Edited on 2/27/2006 by guy]




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[*] posted on 27-2-2006 at 17:00


innervision:

Basically the Cu2+ ion acts as an oxidizer to oxidize the Cu metal to Cu+ [and the Cu2+ gets reduced to Cu+ at the same time]. HCl by itself is rather slow to etch Cu, but the addition of an oxidizer helps.

guy:

Great link you posted :) Thanks. It explains the chemistry and the practical issues very well. I had found this process a couple of years ago on another site that I don't remember right now -- I'll post it if I can find it again.

Incidently I made my initial CuCl2 by reacting CuSO4.5H2O [root killer, Zep brand] with CaCl2 [ice melter]. If you do this make separate solutions of each chemical and then mix. The CaSO4 is insoluble and will settle out. Do not try to filter -- its too thick for that. Instead let it stand for a day or so and collect the green CaCl2 solution off the top [syringe, or bail with a scoop]. Filter that to get rid of any remaining CaCl2 and evaporate for CuCl2 crystals.




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[*] posted on 27-2-2006 at 17:08


The etching makes sense now. That was a great link. I saved that page so I can try it out sometime.
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