Sciencemadness Discussion Board - Unknown copper compound

Unknown copper compound

CyrusGrey - 5-2-2008 at 15:14

Several months ago I used HCl with a little CuCl to dissolve the copper cladding off of some graphite welding rods. Well that had been sitting with the now stripped rods in it until now. I took the rods out and thought that I would try and do something useful with the liquid left over, it was a dark green at this time. I put a copper pipe in it that had been corroded with HCl fumes. After several days I ended up with a very dark brown liquid that I cant identify (It looks black, almost so even in a test tube).

I did the following tests on it:
Dilution with distilled water yields a white precipitate (CuCl?) and the solution has a very light greenish tint

Heating the liquid seems to do nothing, then when its diluted it gives the white precipitate but the brown remains

Adding ascorbic acid makes the brown color much lighter, heating this makes it darker again but not to the original darkness

Adding KNO3 seems to do nothing but then when this is heated it bubbles for a moment, releasing a brown gas the same color as the solution (NO2?) and the solution turns the dark green color of the solution before I added the pipe, when this is diluted it becomes bluish-green

Heating a small amount of the liquid until dryness yields a black solid, which further dehydrates to a tan color, adding water back makes it green

What is this stuff? I'm still a bit of a chemistry newbe, but I've looked around and havent found any brown soluble copper compounds with the stuff I started out with.

12AX7 - 5-2-2008 at 16:32

http://www.sciencemadness.org/talk/viewthread.php?tid=9882

blogfast25 - 6-2-2008 at 07:16

CyrusGrey:

Yes, that looks talks and walks very much like what is suspected to be a Cu+ - Cu2+ - Cl- complex discussed on that thread, no doubt about it.

Interestingly enough, you seem to have actually isolated the compound in "pure" (i.e. crystalline) form. I've been meaning to do that too but haven't gotten round to it yet. It would confirm what I've suspected for some time: the complex is remarkably stable.

woelen - 6-2-2008 at 14:18

Quote:

Originally posted by CyrusGrey
Dilution with distilled water yields a white precipitate (CuCl?) and the solution has a very light greenish tint

Yes, this is CuCl, contaminated with some blue copper(II) compound. This contamination makes it appear exceptionally white. Any attempt to isolate and dry the white material makes it green or even brown, due to partial oxidation.

Quote:

Heating the liquid seems to do nothing, then when its diluted it gives the white precipitate but the brown remains

This is understandable. There is an equilibrium between CuCl and the complex CuCl2(-) and a dark brown copper(I)/copper(II) complex. Your solution contains excess copper(I). On dilution, the complex dissociates into CuCl and chloride. The CuCl precipitates.
I did the reverse. I have some CuCl (commercial sample). This dissolves in conc. HCl very easily, while it is insoluble in water. The solution is almost colorless, but when air is allowed to reach the solution, then it quickly turns deep brown. When much more air is allowed to act upon the solution, then it turns lighter again and it becomes green/yellow, due to formation of CuCl4(2-).

Quote:

Adding ascorbic acid makes the brown color much lighter, heating this makes it darker again but not to the original darkness.

Understandable. The ascorbic acid reduces any copper(II) in solution, giving less of the copper(I)/copper(II) complex and making more colorless CuCl2(-) in solution.

Quote:

Adding KNO3 seems to do nothing but then when this is heated it bubbles for a moment, releasing a brown gas the same color as the solution (NO2?) and the solution turns the dark green color of the solution before I added the pipe, when this is diluted it becomes bluish-green

The nitrate ion oxidizes any copper(I) in solution, itself being reduced to NO2/NO. This makes the solution green/yellow, because only copper(II) remains in solution. However, the solution becomes dark green, because another new complex plays a role in this. Copper(II), chloride and NO2 form a dark blue complex, which with the yellow/green of the CuCl4(2-) gives a dark green color.

The copper(II)/chloride/NO2 complex also is not a well-known thing. I have written a web-page on that as well:

http://woelen.scheikunde.net/science/chem/riddles/copper+nit...

Quote:

Heating a small amount of the liquid until dryness yields a black solid, which further dehydrates to a tan color, adding water back makes it green

What is this stuff? I'm still a bit of a chemistry newbe, but I've looked around and havent found any brown soluble copper compounds with the stuff I started out with.

The tan material is mainly CuCl2, anhydrous. CuCl2.2H2O is a cyan/blue compound, but the anhydrous material is dark brown, when in the form of big crystals, when it is finely powdered it is much lighter brown. On addition of some water, it becomes green, due to complex formation.

-----------------------------------------------------------------

Copper is my favorite element. It has such a rich aqueous chemistry and there really is a lot to be discovered and explained. Even modern textbooks do not mention all these remarkable complexes, while they can be obtained with such common and simple reagents. Maybe modern chemists seem to find this kind of chemistry too common, possibly overlooking some interesting aspects of it.

[Edited on 6-2-08 by woelen]

blogfast25 - 7-2-2008 at 09:11

Hi Woelen:

I saw your page on the Cu2+/NO2 complex, obtained by adding NaNO2 to Cu2+ solutions.

But what you (and CyrusGrey) describe in relation to NO3 (-) does not sit well with my own observations.

At first, I've been using mixtures of Cu(NO3)2 (0.4 M) and NaCl, to simulate CuCl2 (I didn't have any CuCl2 at the time) and although I've always obtained a dark green solution as well as quite a bit of CuCl precipitating out of the solution when reacting it with copper metal.

Later I did use CuCl2 and the same dark green appeared: in fact the system Cu(NO3)2 = 0.4 M + NaCl = 0.8 M and the system CuCl2 = 0.4 M react almost in the same way, with Cu or even with sodium metabisulfite as a reducing agent: I always get a dark green solution and precipitate of CuCl. In Cu(NO3)2 solutions of higher NaCl concentration, the CuCl forms also but on cooling or dilution of the solution. Never have I noticed significant amounts of NO/NO2 developing using these copper nitrate solutions, even at BP.

Is it possible that the dark green solution is merely due to concentration and that on much more prolonged boiling the solution would darken further and further?

An experiment very similar to (and inspired by) your reduction by means of Zn experiment seems to confirm this. Here I've dissolved a pinch of CuCl2 in a test tube in a few ml of 32 w% HCl and divided it into equal amounts over two test tubes. To the first a small ball of copper wire was added and both tubes were properly stoppered. The colour of the copper wire containing solution gradually darkens over time. From a slight darkening at first, the solution, now about three days into the experiment, has turned almost black but is still transparent and without any precipitate (Cl- does indeed complex CuCl into CuCl2 (-) very well).

Interestingly, up to and including yesterday, the bottom of the test tube (where the copper wire rests) sometimes cleared up completely, presumably where Cu2+ had been reduced to Cu+ completely and without agitation two layers formed.

The shades of green (and now brown) in this experiment are different from the ones obtained with Cu(NO3)2 but then I never used such high concentrations of Cl- and H3O+ before either.

It now remains to be seen if in the coming days the solution will actually clear when all the Cu2+ has been reduced to Cu+ and the presumed Cu+ - Cu2+ - Cl- complex can no longer exist.

You're right about copper chemistry being very interesting and perhaps parts of it being a little overlooked. And this complex has got me thoroughly hooked, so much so that I'm thinking of carrying out another experiment.

It would be interesting IMHO to try and correlate the colour of the solution to the Cu2+/Cu+ molar ratio.

Using a larger version of the test tube experiment and using small, thin copper plates or copper foil as copper substrate, it should be possible to determine the Cu2+/Cu+ molar ratio (as well as the total amount of copper in solution) by means of the weight loss of the copper substrate and correlate this with the colour of the solution. Even better of course would be to quantify the colour by means of UV/VIS absorption spectrometry...

Have you ever contemplated this?

[Edited on 7-2-2008 by blogfast25]

[Edited on 7-2-2008 by blogfast25]

woelen - 7-2-2008 at 11:24

Blogfast25, one experiment which you could do is the following:

Dissolve some Cu(NO3)2.xH2O in dilute sulphuric acid and add a small pinch of NaCl. Only a small amount of chloride must be added.
Take another test tube with Cu(NO3)2.xH2O and heat, such that you get some brown NO2 gas.
Pour the acidic solution with copper nitrate, a little chloride and the dilute sulphuric acid into the test tube with the NO2. If you do this, I expect that you will get the deep blue complex.

Of course, if you have another easy way of making NO2, then you can use that. I just gave the 'recipe' above, in order to minimize the number of different reagents needed (only copper nitrate, sodium chloride and dilute sulphuric acid).

In your description I read that you did the copper nitrate + NaCl experiment, but not at high acid concentration. Try dissolving some copper nitrate in conc. HCl and then heat. I really expect formation of some orange/brown gas.

The molar ratio experiment sounds interesting, and I am willing to try some of that, but it is really difficult. Solutions of copper (I) in conc. HCl absorb oxygen from air EXTREMELY easily. That really is a big problem when you want to do quantitative work. I'll think about this and see what I can do. A collaboration between different people over here is a good thing and possibly could result in new interesting results

blogfast25 - 7-2-2008 at 12:38

Oh, regards the Cu2+/NO2 complex, I can produce that very easily, I'm sure. I make my Cu(NO3)2 by dissolving copper scrap into 35 w% HNO3, so plenty of NOx generated there (LOL).

But I'm simply not convinced the green colour is a mix between CuCl4 (2-) and the NO2 complex: how otherwise to explain it does also appear (in my conditions) with CuCl2, even with CuCl2 and metabisulfite? An alternative explanation is that concentration of the Cu+/Cu2+/Cl- complex, [CuCl4 (2-)] and the pH affect the colour of the solution. For the most part, I have worked close to pH = 7. Adding a large excess of HCl prevents CuCl dropping out of course (but my original purpose was to harvest the CuCl).

Quote:

Originally posted by woelen
The molar ratio experiment sounds interesting, and I am willing to try some of that, but it is really difficult. Solutions of copper (I) in conc. HCl absorb oxygen from air EXTREMELY easily. That really is a big problem when you want to do quantitative work. I'll think about this and see what I can do. A collaboration between different people over here is a good thing and possibly could result in new interesting results

Careful stoppering, allowing as little headspace between liquid and air as possible and a few other simple measures should be quite effective against oxygen contamination, IMHO.

Another trick I used to use when titrating Ti3+ solutions (with Fe3+, if I recall well) is that of a CO2 blanket: adding a small amount of bicarbonate causes CO2 to form (in acid conditions of course), preventing oxygen from entering the solution. This way you can titrate the highly oxygen sensitive Ti3+ with completely normal equipment (conical flask + burette). Very effective.

Alternatively, one could set up a range of tests, all containing the same amount of the same starting solution (of known [Cu2+]) and all containing an accurately weighed (and more or less equal) amount of copper metal (say wire). The solutions would then all be allowed to run 1, 2, 3, ..., n days and each day the amount of digested copper could be determined, after photographing the solution (or UV/VIS). From there determining the total amount of Cu in solution and the ratio Cu2+/Cu+ could be estimated, without oxygen interference. Copper is heavy, so it provides good gravimetric leverage.

woelen - 7-2-2008 at 12:56

Quote:

But I'm simply not convinced the green colour is a mix between CuCl4 (2-) and the NO2 complex

Yep, I've been thinking about this and did another simple test with the blue copper complex with NO2. You're right, the green color is not a mix with the NO2-complex. The NO2-complex cannot even exist under those conditions, it only exists at moderately low concentration of chloride.

You raised another interesting point. Time to go back to the lab and try the copper(II)/HCl/KNO3 mix myself

12AX7 - 7-2-2008 at 14:22

In my experience, "dilute" solutions of the brown complex appear dark green. This can be seen by making a strong CuCl2 solution and adding a reducing agent (e.g., Cu), maybe with a little acid. The color goes from green (CuCl4(2-)) to dark green to brown.

Tim

woelen - 7-2-2008 at 14:35

Back from the lab

I now have come to the conclusion that the dark green color is due only to increase of concentration. The nitrate does not really matter.

I did the following experiment:
- dissolve 100 mg CuCl2.2H2O in 3 ml of conc. HCl (30%). This results in a yellow/green solution. On heating, this solution becomes a little bit more yellow, less green.
- dissolve 500 mg of CuCl2.2H2O in 1 ml of conc. HCl (30%). This results in a very dark green solution. On heating, this solution becomes very dark brown. On cooling down, it becomes dark green again.
So, it all is a matter of concentration, but the color also depends on temperature.

To both solutions, I added some NaNO3. This does not result in a drastic change. On heating, the color again changes, it becomes brown, and some gas is produced (strong smell of Cl2, but also ONCl, which has a peculair odour). The color after heating and cooling down again is somewhat different than without the nitrate added, but that can be explained. ONCl has a strong orange color in solution, this affects the total observed color.

-----------------------------------------------------------

Finally, I did the following experiment:
- take 1 ml of HNO3 (52%)
- take 2 ml of HCl (30%)
- mix both chemicals
- add 1 spatula full of fine copper powder (< 50 um particle size)
- heat

On heating, the liquid quickly turns dark green, but soon, it becomes totally black and opaque. Probably the copper (I)/copper(II)/chloride complex is formed. I added another 2 ml of the HNO3/HCl mix. This did not really change things. It only became a little bit lighter and deep green again. On heating, again it became totally black.

Then I put aside the black, still hot solution. After half a minute or so, suddenly, a very violent reaction started. The material in the test tube was spewn out in an eruption of brown gas and a dark green liquid. It was quite scary

. I had the test tube in a sink, and inside the sink, I could see a brown/yellow color of NO2. A lot of this gas was produced. I ran away from the lab not wanting to breathe such a high concentration of this gas, such that one can see it

. Here endeth my experiment

.

More will follow, but now I know about the possibility of violent runaway, I will have to be more careful. I think that the dark black material, dissolved in the HNO3/HCl mix suddenly is oxidized at once. I saw the reaction starting near the surface and within two seconds a 'reaction front wave' propagated through the liquid. I must not think of what could have happened if this was 50 ml or so. Now I only had 5 ml of liquid.

CyrusGrey - 7-2-2008 at 14:55

Woelen: I also observed this reaction when I added the KNO3. I only used maybe 50mg of KNO3 though so it didnt spew out so much NO2. It seems to turn from brown to green with no change in concentration on addition of nitrates or heating in air. I'm going to try some other oxidisers and see if I get the same reaction.

The copper in my beaker of the brown solution continues to dissolve, but only slowly now. There is also a layer of white CuCl precipitate in the bottom now. Also at the interface between the air and the solution where the copper pipe comes out of it there is a green and white solid covering it.

CyrusGrey - 7-2-2008 at 15:31

I put some of the brown solution into three test tubes, about 2mL each.

In the first one I put about 50-100mg of potassium permaganate, it immediatly turned dark green, and the test tube was warm.

In the second I put about 150mg of Ca(OCl)2. It immediatly started bubbling and frothed up. The smell of chlorine gas was strong, and the test tube was again very warm for the amount of solution.

In the third I put a few drops of 17-18% H2O2 solution. I got a very nice seperation of the solution into two layers. The top was the dark green and the bottom was brown. I gently shook the test tube a little, and the layers just did not want to mix. I added about a dozen more drops and I got three layers of solution, the topmost one was a light blue-green. There was a perfectly clear boundry between each layer. Again I shook the test tube and the brown layer slowly dissapeared after about 30 seconds of shaking. The blue-green and dark green layers didnt want to mix either. During this time the solution was bubbling slightly and the tube was warm.

12AX7 - 7-2-2008 at 20:00

As I reported here,
http://www.sciencemadness.org/talk/viewthread.php?tid=9831
copper(I) solutions can act as a reasonable oxidation-absorber, acting catalytically between anything oxidizing (chlorate, chlorine, H2O2, etc.) and anything reducing (copper(0), etc.). Excess oxidizer will result in an excess of its active species (HClO3, O2, etc.).

Tim

woelen - 8-2-2008 at 05:26

CyrusGrey, the reaction with H2O2 indeed is very nice. I made a picture of that, quite some time ago already:

[img]http://woelen.scheikunde.net/science/chem/riddles/copperI+copperII/exp0004/exp0004-1.jpg[/img]

A nice multilayer system is formed, which only very slowly degrades.

EDIT: ????
please copy the url between the [img] tags. I don't understand why the image is not included in the message itself.

[Edited on 8-2-08 by woelen]

blogfast25 - 8-2-2008 at 08:05

Yep, no doubt the colour of these complexes depends on concentration, other spectator ions (such as H3O+) and probably temperature. I was reading up on the VIS absorption spectra of coordination complexes in Shriver and Atkins' 'Inorganic Chemistry' and the quantumchemical explanation of these spectra is quite something. I'll need to do some re-studying to come to grips with that stuff again!

My test tube experiment (HCl, CuCl2 and Cu) now looks completely black and doesn't seem to change anymore (but it's hard to tell). There's plenty copper left. I'll let this run on to see if eventually this solution clears or whether the Cu2+ - Cu+ - Cl- (my tentative formula is Cu2Cl4 (-) ) resists further reduction in these conditions.

For the Cu+/Cu2+ molar ratio experiment I now have a tentative design. I'll be using an old jam jar of more or less cubical shape (clear glass), about 100 - 200 ml volume, loaded with a CuCl2/HCl solution and a few clean (cleaned with HNO3) pieces of copper gas tubing. It'll be easy to remove (and reload them after weighing) the copper pieces, without too much oxygen contamination (although some is inevitable).

I haven't decided on the CuCl2 concentration yet. A 0.4 M solution will discolour quickly and the colour will probably 'saturate' quite fast. A lower concentration will perhaps allow a larger range of colours but the experiment will take longer.

Another thing that would be interesting to do is to check what kind of precipitate is obtained when at various degrees of reduction when you drop some of the solution into dilute NaOH. You [Woelen] obtained a brownish yellow but I've obtained a whole plethora of colours (unfortunately not in systematically controlled conditions), including this kind of green, more kakhish green, brown and occasionally even blue.

In one case I diluted the solution (after removing the copper) quite a bit and obtained the same kind of precipitate from the more concentrated and from the diluted solution. This seems to suggest the complex is indeed one well-defined species.

It's puzzling nonetheless...

[Edited on 8-2-2008 by blogfast25]

[Edited on 8-2-2008 by blogfast25]

smuv - 9-2-2008 at 13:44

Woelen, I was interested by the riddle proposed on your website. I think I may have figured out why the complex is so short lived and changes color.

NO2 has linkage isomerism; It can coordinate with its oxygen lone pair (nitrito ligand) or its nitrogen lone pair (nitro ligand). For example there are two [Co(NO2)(NH3)5]++ complexes. The nitrito compound is red; the nitro compound is yellow. Over time the unstable nitrito complex forms the yellow nitro complex. It seems very possible that you are observing the same effect with a different complex.

Complexation of CuCl - Estimate of Kf

blogfast25 - 10-2-2008 at 07:04

Assume the following system in equilibrium:

An aqueous solution of solid CuCl, a small amount of a strong, monoprotonic acid HA (but not HCl) to suppress OH(-) (so, [OH(-)] = 0), and a concentration c (mol/L) of a strongly dissociated chloride, MCl.

MCl --> M(+) + Cl(-) [M(+)] = c (eq. 1)

HA + H2O --> H3O(+) + A(-) [H3O(+)] = [A(-)] (eq. 2)

CuCl <--> Cu(+) + Cl(-) Ks = [Cu(+)] . [Cl(-)] (eq. 3)

CuCl + Cl(-) <--> CuCl2 (-) Kf = [CuCl2 (-)] / [Cl(-)] (eq. 4)

Neutrality:

[H3O(+)] + [M(+)] + [Cu(+)] = [CuCl2 (-)] + [Cl(-)] + [A(-)] (eq. 5)

From these equations [CuCl2 (-)] and the other concentrations can be derived:

[CuCl2 (-)] = Kf .{[ c + √(c^2 + 4.(1+Kf).Ks) ] / [2.(1+Kf)]} (eq. 6)

A few special cases:

a. c = 0, then [CuCl2 (-)] = Kf. √[Ks/(1+Kf)] (eq. 7)

[Cl(-)] = √[Ks/(1+Kf)] (eq. 8)
and [Cu(+)] = Ks/[Cl(-)]

b. c large and Ks << Kf

[CuCl2 (-)] ≈ c.Kf/(1+Kf) (eq. 9)

Eq. 6 is a positively sloped, continuous function with eq. 7 as intercept and tending to eq. 9 asymptotically (a practical determination of Kf would probably boil down to determining the slope of this asymptote).

An estimate of Kf can be obtained as follows. Ks = 1.7 10^-7 (mol/L)^2 and from Wiki the solubility of CuCl in water at 20 DC is S = 0.0062 g/100 mL or 6.3 10^-4 mol/L. Note that this value of S does not specify method or pH and is a single data point.

From S = {[CuCl2 (-)] + [Cl(-)]} and eq. 7 and 8, Kf can then be estimated:

Kf (est.) = 1.34 and pKf (est.) = - 0.13

Although this is a very rough estimate, it sits quite well with the observation that in concentrated solutions of HCl or NaCl, CuCl is quite soluble. For 2 M NaCl, [CuCl2 (-)] ≈ 1.15 M for example.

++++++

Unfortunately this also shows (by another route) that the proposed Cu(2+)/Cu(+) molar ratio experiment is unlikely to shed much light on the composition of the unknown Cu(2+)/Cu(+)/Cl(-) complex, as in the presence of excess HCl or NaCl, much of the Cu(+) is present as CuCl2 (-).

In the absence of excess HCl or NaCl, the overall Cl to Cu ratio remains 2, no matter how far the reduction has been carried, which could tentatively suggest Cu2Cl4 (-): perhaps {Cl(-) - Cu(+) - Cu(2+) - 3 Cl(-)}(-) ...???

[Edited on 10-2-2008 by blogfast25]

woelen - 10-2-2008 at 08:02

Some time ago I found a text on Internet, which also was about a dark copper(I)/copper(II) complex, but I can't find it back anymore

(stupid me, that I did not save the link).

The text made mention of a neutral Cl-Cu-Cl-Cu-Cl complex, with all chlorines in formal oxidation state -1, and both coppers in oxidation state +1.5 (both coppers were mentioned as being equivalent in the complex). The middle Cl is bound in a special way, such that it forms a bridge between the two copper-atoms. The bond also is fairly weak/labile, such that it easily is broken down (e.g. by oxidizers, but also by dilution).

I'll do my best to find it back again. I really hope to find back the text. The text already was somewhat older (IIRC from the 1960's), it was a scanned document.

I can imagine such a structure. A resonance structure over a larger length may indeed explain its strong color.

blogfast25 - 10-2-2008 at 10:02

Possibly... Cl- is a Lewis base and can form bridges between various metal centres.

My test tube experiment has used up all copper (I've now added some more) it seems to go darker still.

But I haven't seen a great deal of evidence that this complex breaks down on dilution...

[Edited on 10-2-2008 by blogfast25]

12AX7 - 10-2-2008 at 12:43

It does. Try diluting 10x or more. You'll see a blue solution with white precipitate, CuCl, suspended.

Tim

CyrusGrey - 10-2-2008 at 19:12

Ah, Woelen, your photography skills are better than mine then. Heres the photo of my mixture:

[Edited on 11-2-2008 by CyrusGrey]

photo.JPG - 14kB

497 - 10-2-2008 at 23:57

Today I accidentally found out an easy way to make what I'm pretty sure is CuCl (yellow insoluble copper compound). I had a pretty concentrated NaCl solution with copper electrodes supplied by an old printer DC supply. The anode had a fairly high surface area (a spiral of about 6 feet of wire). When the cathode (3 4inch sections of copper pipe wired together) was fully submerged there was very little visible action at all even with the full 20 volts on it for an hour. The interesting part is when the cathode was mostly pulled out of the solution with just a few millimeters submerged it bubbled strongly and within minutes there was a mass of very fine yellow precipitate floating at the surface. The supply got warm indicating there was actually current (maybe 3 amps) while there was little before. More info when I have a chance. Any idea what exactly is going on here?

blogfast25 - 11-2-2008 at 06:18

Quote:

Originally posted by 12AX7
It does. Try diluting 10x or more. You'll see a blue solution with white precipitate, CuCl, suspended.

Tim

Well, as luck would have it, yesterday I had to stop my test tube experiment because the cork stopper had suffered too much from the HCl fumes. By then the solution was basically black.

I diluted it 1:1 first and that caused the black to shift to dark brown. No precipitate formed. Not after standing overnight either.

I then diluted that solution another 1:1 time (by then it was 1 volume original : 3 volumes deionised water). It shifted to that characteristic green (distinctly different from the cyan/greenish of CuCl2). Still no precipitate formed. Not after standing overnight either.

That the complex hydrolyses in high dilution is very well possible but I think in most cases the precipitate is causes by hydrolysis of CuCl2 (-). If your solution contains much free Cl(-) (from HCl or NaCl) then much of the Cu(+) will be complexed as CuCl2 (-) and diluting it pushes CuCl + Cl(-) <---> CuCl2 (-) to the left.

The solution used here was of course in 32 w% HCl, so even at 1:4 dilution it's still about 2.5 M in HCl, perhaps strong enough to keep Cu(+) complexed.

Interestingly, adding some of the (1:4) diluted solution to the original control (CuCl2 in strong HCl) caused the control to shift to a colour slightly darker (green) than the two separate solutions. Maybe some more complex is being formed from the excess Cu(2+) in the control and some Cu(+) in the treated solution.

Having read your post, I've just poured a fairly small amount of the (1:4) diluted solution into a relatively large volume of dionised water and I see no CuCl suspended at all. The colour seems simply diluted.

I'll check what kind of precipitate these solutions yield with diluted NaOH.

I'm setting up three more test tubes: one with pure CuCl2, one with CuCl2 + HCl, one with CuCl2 + NaCl. All in comparable concentrations, all with Cu present.

497:

CuCl is actually white, not yellow. It's possible you're forming CuCl and that it's hydrolysing to Cu2O, which when very fresh and finely divided would look yellowish. Can you check the pH in the area where the yellow is appearing?

woelen - 11-2-2008 at 07:41

@497: As mentioned above, the yellow material most likely is Cu2O, finely divided. I know this yellow/ochre material you get. You also get it by simply taking two electricity wires and using that as electrode set with a 9V power supply. For me, it is not really clear why Cu2O is formed, even in a neutral solution of NaCl. I would expect formation of CuCl2(-) in the concentrated solution of NaCl. Yes, welcome to the complex world of copper chemistry

.

@CyrusGrey: That picture you show resembles what I have, the main difference being that you start with an already very dark solution, and I start with a solution, which only contains a small amount of the dark complex. In my situation you can see that there is a dark band, which is caused by the Cu(2+), formed just below the H2O2-layer, which reacts with Cu(+)/Cl(-) in the lower part of the liquid.

@blogfast25: The CuCl-precipitate is obtained if you dissolve a lot of copper metal (as much as possible) in a mix of equal volumes of HCl (30%) and H2O2(20%). Add a lot of copper wire (the thin wires from electricity cord is best) to conc. HCl and then slowly with stirring add H2O2 (20%) in small steps. Each time, when bubbling stops and the liquid goes from green to dark brown/black, then add some new H2O2 (20%). Continue, until all H2O2 is added. Then wait, until really no more copper is dissolved. This takes some time.

Now, pour the solution in a large excess of water. An amazingly large amount of very white CuCl is precipitated in the form of small crystals. The precipitate is not slimy, it is a compact crystalline precipitate, which quickly settles. I tried to isolate this precipitate, but as soon as some of this is exposed to air, it discolors, it becomes green/brown. I did not succeed in obtaining a nice dry, still white sample.

blogfast25 - 11-2-2008 at 09:17

Quote:

Originally posted by woelen
@blogfast25: The CuCl-precipitate is obtained if you dissolve a lot of copper metal (as much as possible) in a mix of equal volumes of HCl (30%) and H2O2(20%). Add a lot of copper wire (the thin wires from electricity cord is best) to conc. HCl and then slowly with stirring add H2O2 (20%) in small steps. Each time, when bubbling stops and the liquid goes from green to dark brown/black, then add some new H2O2 (20%). Continue, until all H2O2 is added. Then wait, until really no more copper is dissolved. This takes some time.

Now, pour the solution in a large excess of water. An amazingly large amount of very white CuCl is precipitated in the form of small crystals. The precipitate is not slimy, it is a compact crystalline precipitate, which quickly settles. I tried to isolate this precipitate, but as soon as some of this is exposed to air, it discolors, it becomes green/brown. I did not succeed in obtaining a nice dry, still white sample.

Yep, I'm definitely trying that. My H2O2 is less concentrated (I believe, haven't looked for a while) but it should work nonetheless.

I have seen CuCl drop out of Cu(NO3)2 solutions reduced with Cu or metasulfite many times before but I then end up with a dark green solution which seems to resist hydrolysis very well and is quite stable.

Have you tried filtering the CuCl quickly on a Buchner? Washing carefully with freshly boiled (but cooled, of course) water? My experience is that, yes, it does oxidise easily but perhaps not as quickly as you describe.

According Wiki, it's prepared industrially by direct chlorination of copper metal. There's a rainy day project, I guess...

The three solutions (1:4 diluted complex, further diluted complex and the mix of 1:4 and control), all yielded precipitates with varying hues of green. Again to me this suggests not a great deal of hydrolysis has taken place upon dilution, because then I'd expect some clean blue from Cu(OH)2. From very concentrated solutions of the complex things may of course be very different.

My three test tubes are now in place:

#1: a few mL of 0.1 M CuCl2, plus aq. dist + 0.4 g of copper wire - cyan colour

#2: same amount of 0.1 M CuCl2, saturated with NaCl + 0.4 g copper wire - green colour

#3: same amount of 0.1 M CuCl2 in 32 w% HCl + 0.4 g of copper wire - green colour, slightly darker than #2.

If I don't get any CuCl, I might still run the molar Cu(2+)/Cu(+) experiment with 0.1 M CuCl2.

I'm taking some pics too...

blogfast25 - 14-2-2008 at 05:31

Well, the only one out of the three that's actually producing the complex is #3, the one with concentrated HCl. Clearly (at least at RT) the reduction/formation of the complex needs protons or hydronium ions to proceed. #3 is already quite dark and the copper metal keeps clearing the solution due to the reduction. On agitating the solution then darkens again, gradually getting darker and darker. #1 and #2 are basically as they were on day 1.

I tried a little paper chromatography on #3 and the result is slightly unexpected: the eluate (eluent was 32 w% HCl) smears out as a long blob, greenish-yellow in colour. I may have over-loaded the paper (too much solution)

. Might try that gain with proper TLC and a smaller spot (Here's a few pages on how to make your own TLC plates, easier than one might think).

CyrusGrey - 16-2-2008 at 19:43

Did a few more tests today:

Adding sodium carbonate makes it fizz a bit, dissolving most of the carbonate and leaving a very small amount of green precipitate.

Adding sodium hydroxide gives a yellow, bronze, and black precipitate, depending on how much it touched the solution. (Black??)

Adding a pellet of activated charcoal made a very faint fizzing sound, but didnt appear to reduce the solution like the ascorbic acid.

Finally adding aluminum metal gave alot of fizzing, releasing steam and probably hydrogen. It also got so hot that I could feel the heat through 5 or 6 inches of glass test tube! The bottom of it was hot enough that I could almost burn myself on it. This was all from half a mL of the black solution and about 100mg of aluminum! A very cheap method of producing large amounts of hydrogen.

I also noticed that small droplets turn green in a few minutes, probably from atmospheric oxygen. I'm trying to slowly dry some of this on a watch glass to get crystals, maybe some of the copper compounds will seperate out into different crystals.

[Edited on 17-2-2008 by CyrusGrey]

12AX7 - 16-2-2008 at 20:28

It's well known (at least here) that copper chloride solutions have a catalytic effect on aluminum. It's quite striking. Woelen has an example on his website.

BTW, it's not a cheap way to produce hydrogen: it requires the expense of a copper salt, plus the immense energetic demand to produce aluminum. Almost all of that energy is given up in the reaction, as you've noticed; very, very little energy value is present in the hydrogen. Such is often the case with non-industrial hydrogen production.

Tim

[Edited on 2-16-2008 by 12AX7]

CyrusGrey - 16-2-2008 at 21:19

I was actually thinking about using the solution to strip the oxide layer and then let hydrochloric acid do the rest. Cheap because this was leftovers from dissolving copper off of welding rods, before I started testing it I was thinking about how I was going to have to dispose of 550 mL of very concentrated copper solution. It wouldnt be hard to adapt a copper salt and acid method to a more efficient production of hydrogen. HCl and aluminum are cheap for the hobby chemist and only a small amount of copper is needed.

I tested out a similar method awhile back and produced 39.6mL of hydrogen gas from 120mg of aluminum foil.

Looking back at my pictures of this experiment, I just noticed that the solution I used looks a bit similar to the black one I am testing now.

not_important - 16-2-2008 at 21:35

Quote:

Now, pour the solution in a large excess of water. An amazingly large amount of very white CuCl is precipitated in the form of small crystals. The precipitate is not slimy, it is a compact crystalline precipitate, which quickly settles. I tried to isolate this precipitate, but as soon as some of this is exposed to air, it discolors, it becomes green/brown. I did not succeed in obtaining a nice dry, still white sample.

Typically the CuCl precipitate would be washed with dilute sodium bisulfite (NaHSO3) solution, followed by alcohol then ether (acetone work OK I think) The traces of Cu(II) formed dissolve in the alcohol, which drys the CuCl enough that oxidation is slowed down. Some preparations used apparatus that allowed the exclusion of oxygen through a flow of inert gas; to me these mostly looked like much fussing and hassle.

blogfast25 - 17-2-2008 at 06:41

Quote:

Originally posted by CyrusGrey
I also noticed that small droplets turn green in a few minutes, probably from atmospheric oxygen. I'm trying to slowly dry some of this on a watch glass to get crystals, maybe some of the copper compounds will seperate out into different crystals.

[Edited on 17-2-2008 by CyrusGrey]

My own test tube experiment has now come to the end of the line, with all the copper (0.4 g) dissolved in #3 (CuCl2 + HCl) and the solution basically black. Tubes #1 and #2 appear not to have reacted at all.

The amount of copper in solution and the molar ratio of Cu(2+)/Cu(+) can now be roughly estimated for #3. I'll do this later.

"Yes" to the green droplets: drops that stuck to the wall of the test tube (#3) do indeed tend to turn green.

"Yes" also to the various colours obtained when precipitating.

As regards trying to isolate the material (at least for a short while before it air-oxidises), I believe that may well be possible. In my book the substance is more stable that some give it credit for. I will try it myself.

woelen - 17-2-2008 at 23:57

Isolating the stuff is not that easy. I once tried, but on heating it lost a lot of HCl in my case. What remained was a black (very dark brown with a green tinge) solid, which did not dissolve in water. This is more difficult than dehydrating pure CuCl2.2H2O. The latter results in a brown solid, which completely dissolves in water, giving a clear green (or blue when diluted) solution. The dark material, however, forms some insoluble gunk. When this stuff is added to concentrated HCl, it dissolves, giving an almost green solution. Apparently, most of the copper also is oxidized to the +2 state during the heating.

I might retry this, I now have much better equipment, at that time I only had some test tubes and an open beaker, heated above a propane torch. Please tell us your result, but also your procedure, on how you heated the material and how you prevent aerial oxidation of the material during heating.

blogfast25 - 21-2-2008 at 08:01

Well, the 0.4 g of Cu dissolved in about 6 ml of 0.1 M CuCl2 (in 32 w% HCl), means the molarity of Cu in solution is now about 1 M.

That presents in itself a slight mystery, as there wasn't enough Cu(2+) in solution for all this copper metal to dissolve via Cu + Cu(2+) --> 2 Cu(+)! The molar ratio of Cu(2+)/Cu(+) would also appear to have gone past 10 and then to ∞ !

I haven't tested the properties (dilution, precipitate etc) of this batch yet.

It would appear though, simply put, that although Cu is generally considered insoluble in HCL, it is soluble in HCl with some CuCl2 added to it!

As regards isolating the complex, I'm no better equipped than anyone else and cannot protect that stuff from oxygen either. Will have a bash though...

woelen - 21-2-2008 at 11:09

This is REALLY remarkable!! Are you sure that no oxygen from the air could enter the test tube?

Next week, I hope to find more time and then this is one of the things I definitely will try myself!

blogfast25 - 21-2-2008 at 12:30

Well, Woelen, I have thought of that too and I can only say the tube was properly stoppered. But even if some oxygen did leak through, how to still explain the intensely dark colour of the solution, normally attributed to the complex? I'll run a test, guaranteed free of O2 next week...

I'm telling you this for free

: to get to the bottom of all this "strange complex" stuff, we need to work more quantitatively...

woelen - 23-2-2008 at 10:41

I now have 5 ml of 35% HCl (analytical reagent grade) with 100 mg of CuCl2.2H2O (general lab grade) dissolved in it, and put 400 mg of copper wire (from electricity wire) in it. To this, I also added 50 mg of NaHCO3, in order to drive out most air, and in order to keep slight overpressure in the test tube. I do not expect that the NaHCO3 will affect the reaction between the copper and the copper(II) in solution.

The test tube is a screw-type test tube, it is tightly closed, and I have it inverted, with the screw-cap in a layer of water. With the slight overpressure from the NaHCO3 amd the immersion under water I really expect that no air will enter the test tube.

http://woelen.scheikunde.net/science/chem/riddles/copperI+co...

Now I will wait and see how it develops over the next few days.

[Edited on 24-2-08 by woelen]

woelen - 23-2-2008 at 15:44

Approximately 3 hours later, it looks as follows:

http://woelen.scheikunde.net/science/chem/riddles/copperI+co...

The piece of copper wire has reduced all of the dark copperI/copper(II) species to copper(I), and around the wire, the solution is colorless, above the wire, it still is quite dark.

I carefully have taken the test tube out of the erlenmeyer and made a close-up picture:

http://woelen.scheikunde.net/science/chem/riddles/copperI+co...

After taking this picture I shaked a little bit, and I put it back in the water, and now I'll see how it looks like tomorrow.

[Edited on 24-2-08 by woelen]

woelen - 2-3-2008 at 02:00

I cannot confirm the results of blogfast25 with my experiment. The solution becomes totally colorless and the copper becomes nicely red, but it does not dissolve any further.

This is the situation, one day after taking the previous picture:

http://woelen.scheikunde.net/science/chem/riddles/copperI+co...

The final stable result is shown in the picture below:

http://woelen.scheikunde.net/science/chem/riddles/copperI+co...

This stable situation is reached after approximately 1 week. I now have a perfectly colorless solution, with the copper wire in it. I am quite sure that as soon as I open the test tube, it will darken again. I'll wait a little longer, but if things do not change anymore, then I will open the test tube and make a small video of that, seeing what happens.

But I am inclined to think that oxidation of the copper does not go below the +1 stage. As long as the total oxidation state is somewhere between 1 and 2, the liquid is dark green/brown, but when all is at +1 the liquid is perfectly colorless and nothing changes anymore. I think that blogfast25 had some leak in his system, such that oxygen is absorbed slowly from the air. I see that as the only explanation now for dissolving all of the copper while just a small amount of CuCl2 is used at the start.

[Edited on 2-3-08 by woelen]

blogfast25 - 2-3-2008 at 12:49

Quote:

But I am inclined to think that oxidation of the copper does not go below the +1 stage. As long as the total oxidation state is somewhere between 1 and 2, the liquid is dark green/brown, but when all is at +1 the liquid is perfectly colorless and nothing changes anymore. I think that blogfast25 had some leak in his system, such that oxygen is absorbed slowly from the air. I see that as the only explanation now for dissolving all of the copper while just a small amount of CuCl2 is used at the start.

[Edited on 2-3-08 by woelen]

I think Woelen is right. Picture one is exactly as my situation after a few days.

But mine continued to get darker and darker.

By sheer coincidence I've kept tube #3 because I wanted to do some more work on it. Having looked at it now, it's gone to a deep emerald green, the type I've seen appear over and over when reducing Cu2+ salts in acid conditions in the presence of air.

Considering Woelen's set up is more airtight than mine, I believe the end of my result would be due to air infiltration.

I will repeat my experiment with tighter bounds on air.

If I'm right that air was the problem then it shows nonetheless that HCl bearing Cu2+ in the presence of air is a solvent for Cu, albeit a slow one.

In your case, we have to assume the solution is colourless CuCl2 (1-).

Woelen, could you dilute it to confirm this? CuCl (s) should precipitate upon dilution...

Thanks for your input!

[Edited on 3-3-2008 by blogfast25]

Nicodem - 2-3-2008 at 13:11

Quote:

Originally posted by blogfast25
If I'm right that air was the problem then it shows nonetheless that HCl bearing Cu2+ in the presence of air is a solvent for Cu, albeit a slow one.

I have not read the entire thread so this might already have been said. A mixture of diluted HCl and CuCl2 is one of the standard etching solutions for the copper in the making of electronic circuits. When I was a teenager electronics was my main hobby and I used just one such solution for making plates for electronic circuits. It is one of the least effective etching solutions, but I liked it because it is self regenerating. After the copper is dissolved in it, the previously green blue solution turns dark green (due to the dark [CuCl2]- complex anion). I regenerated it by simply adding a bit more conc. HCl and letting it stand on air for a couple of days (to oxidize all Cu(I) to Cu(II)) until it regained the original green blue color. Found the instructions in one old electronics book.

blogfast25 - 3-3-2008 at 11:30

Quote:

Originally posted by Nicodem
[...] blue solution turns dark green (due to the dark [CuCl2]- complex anion). I regenerated it by simply adding a bit more conc. HCl and letting it stand on air for a couple of days (to oxidize all Cu(I) to Cu(II)) until it regained the original green blue color. Found the instructions in one old electronics book.

The darker colour is due to the complex of Cu+/Cu2+ or an intermediate oxidation state [+I/+II] complex of Cu, which we've been discussing here all along.

[CuCl2]- (Cu [+I]) is colourless, [CuCl4]2- (Cu [+II]) is green.

woelen - 8-3-2008 at 12:40

Quote:

Woelen, could you dilute it to confirm this? CuCl (s) should precipitate upon dilution...

I stiull have the sample under water, and it still is perfectly colorless. I am willing to do a final thing with it (but then it will be gone), so could you be a little more specific about what you want me to do with it. You now mean simply diluting it with plain water? If that is done, I expect it to become dark again (as soon as it comes in contact with air) and there might be some white precipitate. Only _might_ be, because the concentratin is not that high.

Nicodem - 9-3-2008 at 08:26

Quote:

Originally posted by blogfast25
The darker colour is due to the complex of Cu+/Cu2+ or an intermediate oxidation state [+I/+II] complex of Cu, which we've been discussing here all along.

[CuCl2]- (Cu [+I]) is colourless, [CuCl4]2- (Cu [+II]) is green.

What "complex of Cu+/Cu2+" are you talking about? I never heard about two cationic species forming complexes. As far as I know, complexes are cations liganded with Lewis bases and not to other cations. I'm not an expert in inorganic chemistry so I would appreciate if you could be more specific since I find the chemistry of copper quite interesting.
Also, given that copper(I) disproportionates in equilibriums containing Cu(0), Cu(I) and Cu(II), wouldn't it be most logical that the dark color comes from colloidal copper? This should be easily checked by testing for colloids with the light dissipation test.

12AX7 - 9-3-2008 at 09:33

Cu(I) disproportionates in H2O solution. This entire thread is specifically discussing a Cu(I)-chloride complex, where it is stable in solution. If there were colloidial copper, that would suggest it could deposit from solution in any size, granular particles or even as a mirror, in any case removing copper from solution, but this is not observed.

There is more info in the rest of this thread and in related threads.

Tim

Nicodem - 9-3-2008 at 10:21

Quote:

Originally posted by 12AX7
Cu(I) disproportionates in H2O solution. This entire thread is specifically discussing a Cu(I)-chloride complex, where it is stable in solution.

I thought the thread was about the reduction of copper chloride aqueous species by elemental copper in which case we can not talk of Cu(I) chloride complex only – more like of a Cu(I) chloride in acidic media with a bunch of electrons added making things a bit more complex (or should I say making more complexes?).

Quote:

If there were colloidial copper, that would suggest it could deposit from solution in any size, granular particles or even as a mirror, in any case removing copper from solution, but this is not observed.

I admit I'm a bit ignorant on inorganic and colloidal stuff, but one thing I know of colloids is that they do not deposit from solution unless aggregation is induced. I don't understand how could copper deposit from such a colloidal solution where other Cu species coexist. This would require the particles to aggregate or grow which is hard to imagine in one such equilibrium where the surface of the particles is in constant Cu(0)<->Cu(I) dynamic exchange. In simple words, would not such particles be too "slippery" to aggregate?
Also, the deposition of more copper than there could be present in colloidal form (surely an incredibly tinny amount, less than what is required to form a precipitate or mirror) would mean that the concentration of Cu(II) species in solution would have to increase above the equilibrium level which is impossible since this would cause the copper precipitate to redisolve.

woelen - 9-3-2008 at 12:21

@Nicodem: This has nothing to do with colloids, but a complex is formed.

Copper (I) chloride and the CuCl2(-) complex are colorless.
The copper(II) complex CuCl4(2-) is green/yellow.

When a colorless copper(I) solution in conc. HCl is exposed to air, then it first becomes very dark, almost black, and on longer standing it becomes yellow/green. What is interesting here, is the intermediate stage, where the copper is between +1 and +2. Apparently, multiple copper ions and chloride ions combine to form a very dark mixed oxidation state complex.

Mixed oxidation state complexes are not that rare. Prussian blue is another example and the molybdenum blues are yet another example. I myself also discovered a mixed complex of titanium(III) and titanium(IV). A lot of this stuff I have written on my website in the riddles section:

http://woelen.homescience.net/science/chem/riddles/copperI+c...
http://woelen.homescience.net/science/chem/riddles/titanium+...

Edit(woelen): Made links work again.

[Edited on 30-7-16 by woelen]

Nicodem - 10-3-2008 at 00:59

Aha! But "solutions" of Prussian blue are colloidal!

OK, forget about colloidal copper. Your experiment of mixing solutions of CuCl in HCl(aq) with CuCl2 in HCl(aq) clearly indicates that the dark coloration only occurs when the average Cu oxidation state is between 1 and 2 (while colloidal copper could only exist at the average oxidation state of <1).

So the first thing to do in order to get an idea of what kind of complex we are dealing with is to check Gmelin. Woelen, on your page you say this complex is not described in the literature, but are you sure you checked thoroughly?

Then there is the color. We do not get a specific color, but more like dark green to almost brown, which in my opinion indicates several absorption bands. I would say this indicates the cooexistance of several complexes each with a specific absorbtion band rather than a single species with several absorbtion bands (isn't it more common for complexes to have only one absorbtion maxima?). An experiment that could be designed to give a good answer would be to measure UV/Vis spectra of solutions containing various ratios of Cu(I)/Cu(II) in HCl. Then measuring the same at a diffent concentration of HCl would give an answer about the competing of the chloride ligands in the complex(es).

Also, what if these compex(es) is/are only solution species? They might not be isolable compounds. I can easily imagine [CuCl2]- being able to compete with Cl- as a ligand for Cu(II) cations, yielding a plethora of possible complex species with low stability constants (stable only in solution resulting in the color change), too low to exist as discrete compounds. They would also have to have high visible light absorption due to intervalence charge transfer trough the Cu-Cu coordination bond. See the scheme for better understanding what kind of complex(es) I mean. Hopefully someone with a better understanding of inorganic chemistry can dispute these hypotheses, but at the moment this is all I can add to the discussion.

copper.gif - 2kB

Nicodem - 10-3-2008 at 02:58

A simple google search on this topic yielded a few papers on this topic, but unfortunately the studies were done on solid state chemistry (the putative complexes in solutions of Cu(I,II) chlorides in HCl are indeed deemed to be only solution phenomena and apparently do not form isolable compounds).

See (and references therein):

On Some Dark-colored Chlorocuprates(I, II) and Related Compounds. I. The Method of Preparation and Some Properties (freely available)
Masayasu Mori
Bulletin of the Chemical Society of Japan, 33 (1960) 985-988. DOI:10.1246/bcsj.33.985

On the possibility of homonuclear intervalence charge-transfer bands in non-stoichiometric cuprous chloride (attached)
A. Goltzene, C. Schwab, S. Nikitine
Physica Status Solidi A, 22 (2006) 465-471. DOI: 10.1002/pssa.2210220212

Charge transfer in mixed-valence solids. Part I. Crystal spectra of chlorocuprates(I,II) (attached)
P. Day and D. W. Smith
J. Chem. Soc. A, (1967) 1045-1046. DOI: 10.1039/J19670001045

There are a couple more papers, but I guess this will do for the moment.

Attachment: copper(I,II) chloride complexes.rar (573kB)
This file has been downloaded 426 times

blogfast25 - 10-3-2008 at 08:43

Quote:

Originally posted by woelen

Quote:

Woelen, could you dilute it to confirm this? CuCl (s) should precipitate upon dilution...

I meant dilution with plain water: as the Cu+ is now mainly in the CuCl2- species, on dilution white CuCl(s) should form by pushing CuCl2- <--> CuCl(s) + Cl- to the right.

This what I've experienced many times when reducing cupric chloride solutions with an excess of Cl-present and diluting them: white CuCl then forms. It depends on concentration, sure, but since as you used more or less the same amounts as me, CuCl2- should be around 1 M! It would be simply confirmation of the complete reduction to Cu [+I]. We should really try this reduction in the absence of air but at higher temperature and with agitation.

[edit]

Ooops! Woelen, as your final picture clearly shows, most of your copper hasn't dissolved, so the total molarity will be much lower than 1 M, in which case no solid CuCl may form, ecept on quite extreme dilution...

Quote:

Originally posted by 12AX7
Cu(I) disproportionates in H2O solution.

I disagree: CuCl is quite stable in the absence of air and does oxidise in the presence of it, but rather slowly. Solid CuCl can be kept under water for days and the discolouration (due to Cu2+ formation) is very gradual. Never have I seen it disproportionate. And dry CuCl is stable, period.

Quote:

Originally posted by Nicodem
Then there is the color. We do not get a specific color, but more like dark green to almost brown, which in my opinion indicates several absorption bands. I would say this indicates the cooexistance of several complexes each with a specific absorbtion band rather than a single species with several absorbtion bands (isn't it more common for complexes to have only one absorbtion maxima?). An experiment that could be designed to give a good answer would be to measure UV/Vis spectra of solutions containing various ratios of Cu(I)/Cu(II) in HCl. Then measuring the same at a diffent concentration of HCl would give an answer about the competing of the chloride ligands in the complex(es).

Yes to UV/VIS of course but the varying colour has IMHO a more mundane explanation: present in the reducing solution are a number of species, including CuCl2- (colourless), CuCl42- (blueish green) and the elusive Cu+,2+ complex, at varying concentrations during the various stages of reduction: one cannot expect in these circumstances to see only one typical colour. But it's surprising how quickly the solution turns a very dark brown, almost black, which we believe is typical of the mixed oxidation state complex.

The Cu-Cu structure you propose is one possibility (not sure about your ligands though...) But a structure with a bridging Cl atom (Cu-Cl-Cu) is also plausible. Cl is a bridging atom in several well know transition complexes.

Another thread on this complex can be found here.

[Edited on 10-3-2008 by blogfast25]

[Edited on 10-3-2008 by blogfast25]

[Edited on 10-3-2008 by blogfast25]

Nicodem - 10-3-2008 at 09:39

Quote:

Originally posted by blogfast25
The Cu-Cu structure you propose is one possibility (not sure about your ligands though...) But a structure with a bridging Cl atom (Cu-Cl-Cu) is also plausible. Cl is a bridging atom in several well know transition complexes.

A bridging Cl ligand is a possibility. However, I'm not sure if such a complex would allow for the intervalence charge transfer. On the other hand, even the tiniest concentration of any species containing a Cu-Cu coordination bond would induce an intense visible light absorption so characteristic for intervalence charge transfers. After all, that is what is being observed, an intense darkening of the solutions of chlorocuprates(I,II).

woelen - 10-3-2008 at 09:51

In personal communications, some years ago, it was suggested to me that the soluble complex could be something like Cl-Cu-(μ-Cl)-Cu-Cl. In the other thread, guy also makes mention of such a structure. The Cl-Cu-(μ-Cl)-Cu-Cl then is symmetric, with copper in the +1½ oxidation state (oxidation states need not be integer numbers, another example of non-integer oxidation state is the superoxide ion).

Unfortunately, I have no reference at all, so it remains speculation.

CyrusGrey - 11-3-2008 at 13:25

Wow. I know very little about these things like how complexes form, so you all have somewhat lost me on the exact chemistry of these things.

I did perform another experiment.

I put some of the brown liquid in a watch glass and some in a test tube. In the watch glass I placed a small bit of copper wire. This morning the test tube had turned from the brown solution to the green [CuCl4]2- complex. However the liquid in the watch glass had not. The copper had a little blue corrosion on it where it was sticking out of the liquid.

I then placed some copper wire in the green solution in the test tube and put mineral oil over it (to keep air out). A few hours later when I checked back it was brown again.

I had tried heating this same solution to dryness, or adding H2O2, or adding permanganate, with no smell of Cl2 gas or white HCl fumes. Only with hypochlorites did the solution evolve chlorine. This makes me think that there is no HCl left in solution to just form new brown complex from the copper, and overpowering the green color.

So I think that adding excess copper to the [CuCl4]2- complex will form the brown complex. This would be a good method for storing it too.

woelen - 11-3-2008 at 13:43

Initially it indeed does form the brown complex, but reaction goes on. Finally, the solution becomes colorless and CuCl2(-) is formed.

blogfast25 - 13-3-2008 at 05:07

Well, I had about half a liter of green Cu2+,1+ solution left from various reductions and decided to convert it to oxide with NaOH and boiling.

The precipitate formed was dark-khaki green, not blue as you would expect from Cu2+ plus hydroxide.

The solution had been standing in my shed in a cut-in-half pop bottle for probably more than a month. Clearly much of the Cu+ had not been oxidised to Cu2+.

The dark-khaki green precipitate is also much less floccular than the blue Cu(OH)2.n H2O and on boiling hardly changes colour. Altough this could simply be a physical mix of CuO and Cu2O, potentially it might be a mixed oxide: Cu3O2.

woelen - 13-3-2008 at 06:04

The precipitate you get has an interesting color profile.

If your solution is purely based on copper(I), then you get a bright yellow/orange precipitate.
If it is mostly copper(I) with only a tiny amount of copper(II), then you get a dirty brown/yellow precipitate.
If it is an approximately equal mix of copper(I) and copper(II), then it becomes dark green/brown.
If it is mostly copper(II), then it becomes dirty green or dirty/blue.
If it is copper(II) only then it becomes bright blue in the absence of chloride, or green in the presence of chloride.

The absence or presence of chloride also has a marked effect. Try making a solution of copper(I) in ammonia (can be done by dissolving copper(II) sulfate in ammonia and then adding a solution of sodium dithionite in ammonia). When a concentrated solution of NaOH is added, then bright yellow/orange copper(I)oxide/hydroxide is precipitated. When the same experiment is repeated with chloride present, then the color hardly is different. But as soon as copper(II) is present as well, then the presence of chloride does have a visible effect, which is stronger for higher concentrations of copper(II).

I think that the precipitates you get are mixed (hydr)oxides and chlorides, with also mixes of copper(I) and copper(II) of indeterminate stoichiometry. In simple words: complex stuff!

blogfast25 - 13-3-2008 at 07:37

Yes, I've seen all these colours during various experiments with the [+I, +II] complex.

What I found interesting about this latest batch of precipitate is that, despite the solution having been exposed to air for so long, still considerable Cu [+I] remains. The solution was an emerald kind of green and around 0.5 M in Cu (guesstimate). I thought the colour may by now be entirely due to CuCl42-, i.e. Cu [+II], because there was plenty Cl1- present too but clearly that was not the case: green solutions of CuCl42- precipitate as blue Cu(OH)2.n H2O with NaOH.

12AX7 - 13-3-2008 at 09:01

http://webpages.charter.net/dawill/tmoranwms/Chem_Cu.html#Ox...

The picture is reasonably representative of the color of copper oxychloride.

Mixed with some Cu2O, you'd get a drab color.

Tim

CyrusGrey - 14-3-2008 at 19:13

Quote:

Initially it indeed does form the brown complex, but reaction goes on. Finally, the solution becomes colorless and CuCl2(-) is formed.

Well, I dunno Woelen. It has been 2 or 3 days and the brown color hasn't gotten any lighter, though I do have a layer of CuCl dust building up at the bottom of the test tube.

Is the CuCl2(-) complex unstable?

woelen - 17-3-2008 at 01:16

I dumped the colorless solution of CuCl2(-) in 50 ml of cold water. No precipitate was formed at all. The solution remained colorless and clear, also after stirring.

blogfast25 - 17-3-2008 at 02:20

No oxidation either? Have you tried adding ammonia?

CyrusGrey - 17-3-2008 at 21:11

Sorry Woelen, I take my comment back. It is turning clearer. It took about a week for it to be noticeable, but now I can see my lab lights through the solution in the test tube. I'm guessing that the copper was passivated somewhat by the layer of CuCl formed on it. Either that or the reaction is just very slow.

When it becomes completely clear I'll take the solution and expose it to the atmosphere without copper and see if it changes back to the brown complex.

My hypothesis is that the brown solution is stable on exposure to both metallic copper and atmospheric oxygen at once, but not one or the other.

[Edited on 18-3-2008 by CyrusGrey]