| Pages:
1
2
3 |
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
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
Hazard to Others
Posts: 123
Registered: 20-1-2007
Location: USA
Member Is Offline
Mood: Oooh! Shiny!
|
|
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
Post Harlot
Posts: 4803
Registered: 8-3-2005
Location: oscillating
Member Is Offline
Mood: informative
|
|
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
Hazard to Others
Posts: 123
Registered: 20-1-2007
Location: USA
Member Is Offline
Mood: Oooh! Shiny!
|
|
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
International Hazard
Posts: 3873
Registered: 21-7-2006
Member Is Offline
Mood: No Mood
|
|
| 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
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
| 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
Super Administrator
Posts: 8027
Registered: 20-8-2005
Location: Netherlands
Member Is Offline
Mood: interested
|
|
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
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
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
Super Administrator
Posts: 8027
Registered: 20-8-2005
Location: Netherlands
Member Is Offline
Mood: interested
|
|
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
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
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
Super Administrator
Posts: 8027
Registered: 20-8-2005
Location: Netherlands
Member Is Offline
Mood: interested
|
|
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
Super Administrator
Posts: 8027
Registered: 20-8-2005
Location: Netherlands
Member Is Offline
Mood: interested
|
|
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
Super Administrator
Posts: 8027
Registered: 20-8-2005
Location: Netherlands
Member Is Offline
Mood: interested
|
|
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
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
| 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
Super Moderator
Posts: 4230
Registered: 28-12-2004
Member Is Offline
Mood: No Mood
|
|
| 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 [CuCl<sub>2</sub>]<sup>-</sup>
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.
…there is a human touch of the cultist “believer” in every theorist that he must struggle against as being
unworthy of the scientist. Some of the greatest men of science have publicly repudiated a theory which earlier they hotly defended. In this lies their
scientific temper, not in the scientific defense of the theory. - Weston La Barre (Ghost Dance, 1972)
Read the The ScienceMadness Guidelines!
|
|
|
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
| Quote: | Originally posted by Nicodem
[...] blue solution turns dark green (due to the dark [CuCl<sub>2</sub>]<sup>-</sup> 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.
[CuCl<sub>2</sub>]<sup>-</sup> (Cu [+I]) is colourless, [CuCl<sub>4</sub>]<sup>2-</sup> (Cu [+II]) is
green.
|
|
|
woelen
Super Administrator
Posts: 8027
Registered: 20-8-2005
Location: Netherlands
Member Is Offline
Mood: interested
|
|
| 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
Super Moderator
Posts: 4230
Registered: 28-12-2004
Member Is Offline
Mood: No Mood
|
|
| 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.
[CuCl<sub>2</sub>]<sup>-</sup> (Cu [+I]) is colourless, [CuCl<sub>4</sub>]<sup>2-</sup> (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.
…there is a human touch of the cultist “believer” in every theorist that he must struggle against as being
unworthy of the scientist. Some of the greatest men of science have publicly repudiated a theory which earlier they hotly defended. In this lies their
scientific temper, not in the scientific defense of the theory. - Weston La Barre (Ghost Dance, 1972)
Read the The ScienceMadness Guidelines!
|
|
|
12AX7
Post Harlot
Posts: 4803
Registered: 8-3-2005
Location: oscillating
Member Is Offline
Mood: informative
|
|
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
Super Moderator
Posts: 4230
Registered: 28-12-2004
Member Is Offline
Mood: No Mood
|
|
| 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.
…there is a human touch of the cultist “believer” in every theorist that he must struggle against as being
unworthy of the scientist. Some of the greatest men of science have publicly repudiated a theory which earlier they hotly defended. In this lies their
scientific temper, not in the scientific defense of the theory. - Weston La Barre (Ghost Dance, 1972)
Read the The ScienceMadness Guidelines!
|
|
|
woelen
Super Administrator
Posts: 8027
Registered: 20-8-2005
Location: Netherlands
Member Is Offline
Mood: interested
|
|
@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
Super Moderator
Posts: 4230
Registered: 28-12-2004
Member Is Offline
Mood: No Mood
|
|
Aha! But "solutions" of Prussian blue are colloidal!
OK, forget about colloidal copper. Your experiment of mixing solutions of CuCl in HCl(aq) with CuCl<sub>2</sub> 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
[CuCl<sub>2</sub>]<sup>-</sup> being able to compete with Cl<sup>-</sup> 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.
…there is a human touch of the cultist “believer” in every theorist that he must struggle against as being
unworthy of the scientist. Some of the greatest men of science have publicly repudiated a theory which earlier they hotly defended. In this lies their
scientific temper, not in the scientific defense of the theory. - Weston La Barre (Ghost Dance, 1972)
Read the The ScienceMadness Guidelines!
|
|
|
Nicodem
Super Moderator
Posts: 4230
Registered: 28-12-2004
Member Is Offline
Mood: No Mood
|
|
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
…there is a human touch of the cultist “believer” in every theorist that he must struggle against as being
unworthy of the scientist. Some of the greatest men of science have publicly repudiated a theory which earlier they hotly defended. In this lies their
scientific temper, not in the scientific defense of the theory. - Weston La Barre (Ghost Dance, 1972)
Read the The ScienceMadness Guidelines!
|
|
|
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
| Quote: | Originally posted by woelen
| 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. |
I meant dilution with plain water: as the Cu+ is now mainly in the CuCl<sub>2</sub><sup>-</sup> species, on dilution white
CuCl<sub>(s)</sub> should form by pushing CuCl<sub>2</sub><sup>-</sup> <--> CuCl<sub>(s)</sub>
+ Cl<sup>-</sup> to the right.
This what I've experienced many times when reducing cupric chloride solutions with an excess of Cl<sup>-</sup>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,
CuCl<sub>2</sub><sup>-</sup> 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 Cu<sup>2+</sup> 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 CuCl<sub>2</sub><sup>-</sup> (colourless), CuCl<sub>4</sub><sup>2-</sup> (blueish green) and
the elusive Cu<sup>+,2+</sup> 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
Super Moderator
Posts: 4230
Registered: 28-12-2004
Member Is Offline
Mood: No Mood
|
|
| 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).
…there is a human touch of the cultist “believer” in every theorist that he must struggle against as being
unworthy of the scientist. Some of the greatest men of science have publicly repudiated a theory which earlier they hotly defended. In this lies their
scientific temper, not in the scientific defense of the theory. - Weston La Barre (Ghost Dance, 1972)
Read the The ScienceMadness Guidelines!
|
|
|
| Pages:
1
2
3 |
![[*]](images/xpblue/default_icon.gif){kind=icon}
