Vargouille
Hazard to Others
Posts: 380
Registered: 16-4-2012
Member Is Offline
Mood: No Mood
|
|
Copper Reaction in Ethanol
So today has been an interesting day.
A while ago, I was attempting to make a solution of CuCl2 by dissolving copper wire in 31% HCl and 3% H202, which made a pretty blue solution. Being
impatient, and knowing that high heats would not seriously impact the CuCl2, I evaporated the solution half-way over an alcohol lamp and left it to
evaporate over time. Between this point and the next point, a whitish precipitate formed. Then, still impatient, I heated the solution further, and
IIRC, changing both the precipitate and the solution to a green color. Not worried, I continued heating until the precipitate (seemingly insoluble)
appeared to become solvated. I left it for a good while, and when I returned and agitated the mixture, it solidified in much the same way a
supercooled solution will, forming a green solid.
For some reason I don't recall, I dissolved this solid in denatured alcohol after testing its usefulness as a wash for the solid, lacking a good
non-polar solvent like hexane or xylene. A dark green solution was created, and after a day or two of leaving it to dry, I removed half of the
solution to attempt to form Cu(OH)2 from it with solid NaOH added directly to the solution. It created an apparent green precipitate which quickly
turned to a dark, almost chocolate brown. I had apparently underestimated the heat released, as the solution quickly reached boiling temperatures, at
which point I backed away and allowed it to percolate. Once I returned, I filtered it off, and it is currently drying. It may be noted that it appears
to be such a dark brown as to be nearly black in certain lighting.
Not entirely sure what happened, I took off another sample of the CuCl2 from the stock, and diluted it with water, returning to a blue solution. On
addition of NaOH, now as a solution, the characteristic Cu(OH)2 precipitate was formed. At this point, thinking that the high temperatures may have
decomposed whatever precipitate was formed, I repeated my first trial with the NaOH solution, diluting the CuCl2 solution to aid visibility. With
addition of the NaOH, a greenish precipitate was formed that appeared to be thick, almost like plant matter. Upon stirring, however, this quickly
filled the solution, turning the same chocolate brown in the process. Looks like that's supposed to happen. That solution is currently filtering, as I
am unable to do any filtration that is not a gravity filtration at the moment.
So, what happened? Did the copper form a complex with the ethanol and the hydroxide, creating the precipitate?
EDIT: After some serendipity, I found that the NaOH in EtOH forms sodium ethoxide, which reacts to form the insoluble copper ethoxide, which accounts
for the green precipitate. Is the brown precipitate that results due to moisture in the air?
[Edited on 16-5-2012 by Vargouille]
|
|
|
Lithium
Hazard to Others
Posts: 104
Registered: 25-2-2012
Location: Australia
Member Is Offline
Mood: Thinking!
|
|
copper hydroxide turns into copper (II) oxide when left long enough, here is a video where it happened while trying to synthesise sweizers reagent:
http://www.google.com.au/url?sa=t&rct=j&q=sweizers%2...
Li
|
|
|
weiming1998
National Hazard
Posts: 616
Registered: 13-1-2012
Location: Western Australia
Member Is Offline
Mood: Amphoteric
|
|
Quote: Originally posted by Vargouille  | So today has been an interesting day.
A while ago, I was attempting to make a solution of CuCl2 by dissolving copper wire in 31% HCl and 3% H202, which made a pretty blue solution. Being
impatient, and knowing that high heats would not seriously impact the CuCl2, I evaporated the solution half-way over an alcohol lamp and left it to
evaporate over time. Between this point and the next point, a whitish precipitate formed. Then, still impatient, I heated the solution further, and
IIRC, changing both the precipitate and the solution to a green color. Not worried, I continued heating until the precipitate (seemingly insoluble)
appeared to become solvated. I left it for a good while, and when I returned and agitated the mixture, it solidified in much the same way a
supercooled solution will, forming a green solid.
For some reason I don't recall, I dissolved this solid in denatured alcohol after testing its usefulness as a wash for the solid, lacking a good
non-polar solvent like hexane or xylene. A dark green solution was created, and after a day or two of leaving it to dry, I removed half of the
solution to attempt to form Cu(OH)2 from it with solid NaOH added directly to the solution. It created an apparent green precipitate which quickly
turned to a dark, almost chocolate brown. I had apparently underestimated the heat released, as the solution quickly reached boiling temperatures, at
which point I backed away and allowed it to percolate. Once I returned, I filtered it off, and it is currently drying. It may be noted that it appears
to be such a dark brown as to be nearly black in certain lighting.
Not entirely sure what happened, I took off another sample of the CuCl2 from the stock, and diluted it with water, returning to a blue solution. On
addition of NaOH, now as a solution, the characteristic Cu(OH)2 precipitate was formed. At this point, thinking that the high temperatures may have
decomposed whatever precipitate was formed, I repeated my first trial with the NaOH solution, diluting the CuCl2 solution to aid visibility. With
addition of the NaOH, a greenish precipitate was formed that appeared to be thick, almost like plant matter. Upon stirring, however, this quickly
filled the solution, turning the same chocolate brown in the process. Looks like that's supposed to happen. That solution is currently filtering, as I
am unable to do any filtration that is not a gravity filtration at the moment.
So, what happened? Did the copper form a complex with the ethanol and the hydroxide, creating the precipitate?
EDIT: After some serendipity, I found that the NaOH in EtOH forms sodium ethoxide, which reacts to form the insoluble copper ethoxide, which accounts
for the green precipitate. Is the brown precipitate that results due to moisture in the air?
[Edited on 16-5-2012 by Vargouille] |
First NaOH partially reacts with C2H5OH to form Na ethoxide.
Formula: NaOH+C2H5OH<===>NaC2H5O+H2O.
Source: http://en.wikipedia.org/wiki/Sodium_ethoxide
Then, a precipitation reaction might occur between CuCl2, which is soluble in ethanol, and Na ethoxide, forming Cu ethoxide. The precipitation would
drive the reaction forward, which consumes Na ethoxide and shifts the equlibrium to the right. The Cu ethoxide then reacts with water to form Cu(OH)2.
Source:http://pages.towson.edu/ladon/chemeq.html
Now, what happened to the Cu(OH)2? It decomposed. Cu(OH)2 decomposes spontaneously in the presence of water, and the heat produced by your reaction
probably triggered the decomposition into black CuO.
http://en.wikipedia.org/wiki/Copper(II)_hydroxide
So that's why a black-brown precipitate is formed.
[Edited on 16-5-2012 by weiming1998]
|
|
|
woelen
Super Administrator
Posts: 8027
Registered: 20-8-2005
Location: Netherlands
Member Is Offline
Mood: interested
|
|
If the color is really brown, then most likely you also had reduction of some of the copper(II) to copper(I). Copper(I) oxide can have a color ranging
from bright yellow/orange to brick red, depending on particle size and how strongly it is hydrated. If you have black CuO, mixed with some
reddish/orange material then it may have a color like chocolate.
|
|
|
Vargouille
Hazard to Others
Posts: 380
Registered: 16-4-2012
Member Is Offline
Mood: No Mood
|
|
Weiming, I'm having some trouble with that conclusion. The decomposition of copper ethoxide to Cu(OH)2 should have created a significant blue color
that wasn't seen, and in the second trial, when very little heat was given off, it still rapidly turned brown with no blue intermediate. While Cu(OH)2
decomposes spontaneously in the presence of water, that is to say nothing of the rate, as the sample of Cu(OH)2 that I made in water and left
overnight just has small specks of CuO contaminant.
Can you elaborate?
|
|
|
weiming1998
National Hazard
Posts: 616
Registered: 13-1-2012
Location: Western Australia
Member Is Offline
Mood: Amphoteric
|
|
| Quote: Originally posted by Vargouille | Weiming, I'm having some trouble with that conclusion. The decomposition of copper ethoxide to Cu(OH)2 should have created a significant blue color
that wasn't seen, and in the second trial, when very little heat was given off, it still rapidly turned brown with no blue intermediate. While Cu(OH)2
decomposes spontaneously in the presence of water, that is to say nothing of the rate, as the sample of Cu(OH)2 that I made in water and left
overnight just has small specks of CuO contaminant.
Can you elaborate? |
No intermediate blue colour at all? I'm not totally sure+ I cannot find anything, but I'll give my speculations here:
Maybe, The Cu(OH)2 stage is skipped. So instead of the reaction being Cu(C2H5O)2+2H2O===>Cu(OH)2+2C2H5OH, it is Cu(C2H5O)2+H2O==>CuO+2C2H5OH,
but I doubt that because water is not acidic enough to deprotonate twice in a media like ethanol. Another possible reaction mechanism is the two-stage
reaction of Cu(C2H5O)2+2H2O===>Cu(OH)2+2C2H5OH, but then the Cu(OH)2 was instantly dehydrated by the ethoxides/the NaOH.
Of course, this is just speculation with no citations, but you can test my claims by adding a small amount of water (10%) to your ethanol and use
small amounts of NaOH+CuCl2 in a relatively large amount of ethanol. If what I say is correct, limited or no CuO should form, and you will be able to
see some blue precipitate.
Edit: I did the test. I mixed together dry NaCl and CuSO4 hydrate, then added ethanol, to form CuCl2 on site. NaOH was then added to the reaction
vessel. There was a clear intermediate blue Cu(OH)2 stage, which lasted for about 10 seconds, and then turned into the brown you described. The
ethanol that I was using contains water, and the vessel itself has traces of water in it, so maybe that's why my results were slightly different.
I then added small amounts of water to the ethanol and did the same thing. The precipitate lasted a lot longer (This is a few minutes later after I
mixed them together, and the precipitate is only starting to turn black).
[Edited on 16-5-2012 by weiming1998]
|
|
|
|
![[*]](images/xpblue/default_icon.gif){kind=icon}
