Again a strange copper observation

[woelen]

The element copper remains a very special element to me. It keeps on giving strange results and I want to post one over here again. It is a really simple experiment, just add two chems to each other. In jsatan's words, I saw a "little thing", which I cannot neglect. The "little thing" here is that the copper (II) chloride turns very dark brown, almost black in the concentrated acid. When water is added, then it dissolves again and the results become predictable again.

 

http://woelen.scheikunde.net/science/chem/riddles/copper+h2so4/index.html

 

If anyone over here has an explanation for the observation, then I would be pleased to know.

[xeluc]

Is the dark compound stable when all H2SO4 is boiled away or do you think this complex is only stable under acidic conditions... I understand you said it couldn't be dehydrated CuCl2, but I looked at your pictures and the brown particals resembled my dehydrated CuCl2. Of course, your probably right about the pictured brown stuff being too dark, I don't have any more CuCl2 at the moment, althoguh I am brewing some up . Here's something, do you get the same results with Copper Sulphate? If so, you can rule out Chlorine as having anything to do with the reaction. Also, I'm jsut guessing that CuCl2 is insoluble to H2SO4, so if you mix CuCl2 and some other acid other than H2SO4 that isn't a solvent for CuCl2, if the brown stuff appears, you can say that teh sulphate Ion plays no part. Your a really smart guy though and I'm sure you've thoguht this out much more thoroughly than me, I'm jsut trying to bounce some ideas. It seems reasonable that you could figure out what's in this compound by testing for when it is present, as for another acid that is insoluble with CuCl2, no idea.

 

EDIT: assuming that any water in the H2SO4 didn't hydrate the CuCl2, you could try adding Anhydrous Copper Chloride and seeing how much the compound darkens, if it does.

[woelen]

When you add hydrated copper sulfate to conc. H2SO4, then it remains blue. I tried that before, so the chloride is part of the reaction.

 

I can try adding some anhydrous CuCl2 to sulphuric acid. I'll see if I can make some by heating CuCl2.2H2O. I must be absolutely sure that no remains of HCl are left in the solid and I must also be absolutely sure that no basic oxychloride is formed, so that's why I did not try it yet.

 

Many hydrated salts decompose when heated, not by simply loosing water, but also loosing HCl. E.g. CuCl2.2H2O is prone to loosing HCl when heated:

 

CuCl2.2H2O ---> Cu(OH)Cl + H2O + HCl

 

2Cu(OH)Cl ---> CuO.CuCl2 + H2O

 

Here we have a basic copper chloride, also known as copper oxychloride. This is not a mix of the two, but a crystalline solid, with copper, chlorine and oxygen in a single crystal lattice.

 

With this kind of compounds, the experimental outcome will be quite different when it is added to H2SO4. However, I certainly will try your suggestion and once I'm sure that I have pure anhydrous CuCl2 I'll do.

 

One big problem with investigating this compound is that it is so hard to isolate. Remember, it is in concentrated H2SO4. So, simply boiling off the liquid is out of the question (H2SO4 boils at appr. 300 degrees C and the fumes are insanely corrosive and lethal). Hot H2SO4 at 300 C eats everything in seconds. Filtering also is not an option with the viscous and very corrosive liquid. Conc. H2SO4 really is nasty stuff.

 

If you have some of your CuCl2, could you please compare the colors/darkness and let me know? The pictures I have made are fairly close to the real color. The compound in the acid really is almost black as you can see on my site.

[woelen]

Xeluc,

 

I must admit that I made some big mistake with this copper riddle . You were totally right.

 

You are right with your conclusion that it could be copper chloride, anhydrous. I was fooled by the observation of the light yellow/brown stuff I made by boiling a solution of CuCl2 in HCl to dryness. The stuff I had made before is not pure anhydrous copper chloride, but it most likely was something with HCl incorporated in the solid.

 

Now, I've heated some of my reagent grade CuCl2.2H2O very slowly and carefully, taking care not to overheat it in order to avoid the making of basic copper chloride. The product I obtained is exactly the same as on the picture, but now as a dry powder. A nice dark brown and dry powder, which dissolves in water, giving a green solution, without any residue.

 

With this observation, I regard this riddle as solved .

 

The webpage is updated (not removed).

[xeluc]

well I'm glad you figured it all out. The fact that H2SO4 is used in dehydration led me to believe that. Anyhow, as your usually the one who comes to the rescue when I don't understand something, I figured I'd just post this here.

 

(Note: Nothing is balanced just to simplify things (Actually, I did a little..))

I believe that a solution of HCl and CuCl2 with Cu will Oxidize Cu into Cu+, then O2 Combines with Cu+ into Cu2O and the HCl Oxidizes that into CuCl2 and H2O. I worked out the charges, and they ar ethe same on both sides of the reaction. The Cu+ changes to Cu2+ and the O2- changes to OH-. So does that mean that the O is reduced? It's harder for me to understand since it ends up with an H there also. It'd make more sense for a Peroxide anion to be formed, but that would create H2O2, not H20. Could you clarify this?

EDIT: I now understand that if a Peroxide Ion was formed, there would not be a conservation of charge..

 

Also, What if there was no HCl, just CuCl2 and H2O? I believe one of 2 things could happen.

1: Nothing.. Maybe you need an acidic PH for the Oxidation of Cu2+ to Cu+ to occur.. Dunno.

2: Cu2+ Oxidizes Cu to 2Cu+, 4Cu+ will bind with O2 to create 2Cu2O. This will then precipitate out of solution.

 

Are any of these right? Thanks..

[jdurg]

Oxygen does tend to oxidize a lot of things, and if oxygen is doing the oxidizing, then it's only natural that it would be reduced.

[xeluc]

Oxygen does tend to oxidize a lot of things, and if oxygen is doing the oxidizing, then it's only natural that it would be reduced.

 

Ouch, that was a stupid mistake on my part... Well what about part 2 of my post?

[woelen]

OK, let's go carefully through this post...

 

well I'm glad you figured it all out. The fact that H2SO4 is used in dehydration led me to believe that. Anyhow, as your usually the one who comes to the rescue when I don't understand something, I figured I'd just post this here.

Now you came to rescue me .

 

(Note: Nothing is balanced just to simplify things (Actually, I did a little..))

I believe that a solution of HCl and CuCl2 with Cu will Oxidize Cu into Cu+

Yes, this is right, with the added notion, that both the Cu(2+) and Cu(+) are complexed, so:

 

[math]Cu(s) + Cu^{2+}(compl) -> 2Cu^+(compl)[/math]

 

With (s) I mean solid and with (compl) I mean complexed in some way. Under these conditions there are no free copper (II) and copper (I) ions.

 

, then O2 Combines with Cu+ into Cu2O and the HCl Oxidizes that into CuCl2 and H2O.

No, this is not correct. O2 does act as oxidizer, but there definitely is no intermediate Cu2O. HCl is not oxidizing at all in this place, although it plays an important role.

 

What really happens is that the complexed Cu(+) ions are oxidized by O2 with the help of H(+) ions from the acid:

 

[math]4Cu^+(compl) + O_2 + 4H^+ -> 4Cu^{2+}(compl) + 2H_2O[/math]

 

The H(+) ions are not oxidizing, but they help the oxidation to be carried out. Many oxidizers need H(+) for their oxidizing properties, e.g. permanganate, dichromate.

 

What I did was introduce the notation (compl). I did this to tell you that the reaction between copper (II) and copper metal indeed occurs, but not as simple ions. In reality, it is the complex ion [math]CuCl_4^{2-}[/math], which acts as oxidizer. The resulting ion is [math]CuCl_2^-[/math]. This latter ion contains copper in the +1 oxidation state. When both ions are present, then even more complex things are formed, the deep brown mix-valency complexes, probably ClCu(μ-Cl)CuCl, which contains both copper in the +1 and copper in the +2 oxidation state (or two copper ions in the +1.5 oxidation state).

 

There is another thing I want to point out. You state that Cu2O is oxidized by HCl to CuCl2 and H2O. Just as an exercise, try to balance the following equation:

 

Cu2O + HCl ---> CuCl2 + H2O

 

Please do this exercise and give feedback here. You'll learn a LOT of this exercise! After your feedback I'll come back on this topic.

 

 

 

I worked out the charges, and they ar ethe same on both sides of the reaction. The Cu+ changes to Cu2+ and the O2- changes to OH-. So does that mean that the O is reduced? It's harder for me to understand since it ends up with an H there also. It'd make more sense for a Peroxide anion to be formed, but that would create H2O2, not H20. Could you clarify this?

EDIT: I now understand that if a Peroxide Ion was formed, there would not be a conservation of charge..

Good that you found out yourself. No peroxide can be formed here.

 

Also, What if there was no HCl, just CuCl2 and H2O? I believe one of 2 things could happen.

1: Nothing.. Maybe you need an acidic PH for the Oxidation of Cu2+ to Cu+ to occur.. Dunno.

2: Cu2+ Oxidizes Cu to 2Cu+, 4Cu+ will bind with O2 to create 2Cu2O. This will then precipitate out of solution.

 

Are any of these right? Thanks..

You'll get something between (1) and (2). A little copper will be oxidized and you get a complicated mix of copper (I) and copper (II) complexes, together with some basic stuff, but the reaction will come to an end really soon. The acid is really needed.

[xeluc]

Ok, I'm fairly certain this is impossible to balance. I also now don't know how the oxidation state could even be changed. It seems more likely that instead of that happening, Cu2O just disolves with the help of HCl to make the Cu+ Complex and water.

 

Cu2O + 2HCl --> 2CuCl2(-) + H2O. Interestingly (And logically now that I think about it) Everything works out. Like the Cl's I mean. At first I put Cu+ in the reaction and I'm like.. so where do the Cl's go... And then I remembered the complex that it is in.. It all makes sense now. I really needed to figure that out on my own.

 

So then your Cu+ complex just turns into the Cu2+ complex With the help of O2. If I had some Cu2O I could test this hypothesis, but I have none..

 

Anyway... Thanks for helping to understand all of this, unless I'm still wrong. In which case....hah...

[woelen]

Ok, I'm fairly certain this is impossible to balance. I also now don't know how the oxidation state could even be changed. It seems more likely that instead of that happening, Cu2O just disolves with the help of HCl to make the Cu+ Complex and water.

Exactly, this is what I wanted you to find out yourself. You cannot balance the equation and the reaction is not possible.

 

 

Cu2O + 2HCl --> 2CuCl2(-) + H2O. Interestingly (And logically now that I think about it) Everything works out. Like the Cl's I mean. At first I put Cu+ in the reaction and I'm like.. so where do the Cl's go... And then I remembered the complex that it is in.. It all makes sense now. I really needed to figure that out on my own.

Indeed, Cu2O will dissolve in HCl, forming the complex CuCl2(-). If the HCl is absolutely free of oxygen and there is no contact with air, then you would get a colorless solution. In practice, however, dissolving Cu2O results in formation of a dark brown liquid, because O2 oxidizes part of the CuCl2(-) and then the mix of copper (I) and copper (II) is formed, in which the deep brown multi-valence complex is formed.

 

So then your Cu+ complex just turns into the Cu2+ complex With the help of O2. If I had some Cu2O I could test this hypothesis, but I have none..

 

Anyway... Thanks for helping to understand all of this, unless I'm still wrong. In which case....hah...

I have the impression that you understand this quite well now. As you see, sometimes doing the math yourself may be very enlightening.

[xeluc]

Sorry, on elast thing and I'm out of your hair. if you could keep everything out of reach of Oxygen, if you added NaOH to the CuCl2(-) complex would you get CuOH? I cant find any ino on it, the closest thing is like CuCl2OH or something. I'm trying to make Cu2O, using the limited chems I have....

 

I was asking about the OH because I know most Cu+ compounds are unstable and I was hoping the CuOH would drop out water and become Cu2O, but I don't even think CuOH even exists...

[akcapr]

So Woelen, do you have liek a Phd in like ionic chemistry or something?

[woelen]

Sorry' date=' on elast thing and I'm out of your hair. if you could keep everything out of reach of Oxygen, if you added NaOH to the CuCl2(-) complex would you get CuOH? I cant find any ino on it, the closest thing is like CuCl2OH or something. I'm trying to make Cu2O, using the limited chems I have....

 

I was asking about the OH because I know most Cu+ compounds are unstable and I was hoping the CuOH would drop out water and become Cu2O, but I don't even think CuOH even exists...[/quote']

Indeed, ideally you would get CuOH, which however has another structure. Real copper (I) hydroxide does not exist as a pure compound. What you'll get is hydrous Cu2O, which best can be written as Cu2O.xH2O.

 

In realily, situations are even more complex. Also quite some chloride will precipitate and what you really get is a dirty brown hydrous copper (I) oxide/chloride precipitate. Making Cu2O from a concentrated CuCl2(-) solution, hence is not the best thing to do, because of the coprecipitation of chloride.

 

 

So Woelen, do you have liek a Phd in like ionic chemistry or something?

No, I'm a completely home-made chemist, but I've read a lot of books, Internet pages and did a lot of experiments. That is how I gained my knowledge. In my daily working-life I'm an ICT-consultant and software engineer.

[xeluc]

I can only hope to accomplish what you have, heh. anyway thanks a lot.