Yes, the element copper is intriguing me again. This element has such a remarkably rich chemistry. This time I did some experimenting with copper
sulfate and potassium thiocyanate (the sodium salt and ammonium salt work equally well).
This webpage is far from exhaustive, even more different results can be obtained by simply mixing thiocyanate and copper(II). Other variations are
dependency on pH and the presence of other coordinating ions like chloride. The webpage is just a first step in a direction of research where lots of
things can be discovered, even for home chemists. Unfortunately, the amount of accessible literature on these (iso)thiocyanato complexes is very
limited.Taoiseach - 5-1-2009 at 00:30
Very nice
How stable is the black Cu(NCS)2?
Also I wonder if mixed species can be formed, like
[(Cu2+)(NCS-)]Cl-
Cl- might be a bad choice because it easily coordinates to the copper. ClO4- might work as it rarely every acts as a ligand.Nerro - 5-1-2009 at 01:38
Are you sure the ligand complexes via the N rather than S? Are polymers formed?JohnWW - 5-1-2009 at 03:10
S usually bonds to transition metal cations preferentially to N or Cl- or O. That is why H2S is so poisonous, and also a catalyst poison. An exception
would be the most highly electropositive transition metals, i.e. the rare earth metals (group III-B) , and the group IV-B metals.
[Edited on 5-1-09 by JohnWW]kmno4 - 6-1-2009 at 00:36
There is large amout of literature about Cu - SCN complexes and their structures.
For example:
Bonding properties of thiocyanate groups in copper(II) and copper(I) complexes.
(Coordination Chemistry Reviews,1995 )
An old (1916) article - The action of water on cupric thiocyanate (RSC, CT9160900597) - describes stoichiometry of decomposition of Cu(SCN)2 in H2O.
These are only two from hundred articles about Cu/SCN - - serching machines on chemical databases are for free.
BTW.
Cu(SCN)2 reminds me, similar in colour, K2Pb[Cu(NO2)6]*. Introducing Cu(II) to NaNO2 solution also gives green complex (but green colour = yellow+blue
ones ).
* some books wrongly describe this salt as K2Cu[Pb(NO2)6].
X-ray investigations of crystal structure clearly show existence of Cu(NO2)6 (4-) anion, not Pb(NO2)6 (4-) .
[Edited on 6-1-2009 by kmno4]kmno4 - 12-2-2009 at 15:44
Recently I have made few experiments with Cu(II) in acetone solutions. Starting materials were CuSO4 (anhyd.), KSCN and KBr etc.
Results are interesting:
with KSCN deep red-brown solution (stable) and adding a few drops of pyridine or morpholine gives light green solution.
KBr gives deep green solution and acetone is slowly converting to bromoacetone.
NaCl gives brown-yellow sol., Pb(NO3)2 gives (slowly) nice blue sol. (both experiments in progress).
All proportions was 1-1,5 g CuSO4 + 20 ml acetone+ stoichiometric amout (counting for Cu) of "countersalts".
Really interesting..... try it woelen - 13-2-2009 at 01:13
If you have anhydrous CuCl2, then you can even make an "acetoned" solid salt.
Simply dissolve some CuCl2 in acetone and heat to 40 C or so. Try to dissolve as much of CuCl2 as possible. You get a very dark green/brown solution.
Decant the solution from any remaining solid and let the solution stand for one day. After a day, the solution has become much lighter. A copper-red
solid has settled at the bottom. The solid actually looks a little bit like metallic copper, except that it is slightly more yellow and if the
crystals are large, then it is slightly transparent as well.
This red/brown solid has formula CuCl2.xCH3COCH3. I do not know the value of x though.
A nice further experiment is to decant the liquid from the crystals and let them dry (which only takes a few minutes in luke-warm air). Next, heat the
crystals and light the vapor coming from the crystals. The vapor is acetone. The crystals become dark brown: free anhydrous CuCl2.ChemSys - 13-2-2009 at 12:28
You can try another experiment with CuSCN...
Heat white dry CuSCN in a test tube. It produces CS2 vapors and some other substances. Carbon disulphide easily condenses into liquid state by
cooling. IMHO, it's a quite simple way to obtain 1-2 ml of impure CS2.kmno4 - 13-2-2009 at 16:18
Quote:
Originally posted by woelen
If you have anhydrous CuCl2, then you can even make an "acetoned" solid salt.
Simply dissolve some CuCl2 in acetone and heat to 40 C or so. Try to dissolve as much of CuCl2 as possible. You get a very dark green/brown solution.
Decant the solution from any remaining solid and let the solution stand for one day. After a day, the solution has become much lighter...
Sounds interesting. I do not have CuCl2 (even x H2O) but CuCl2 is generated from reaction CuSO4 and NaCl in acetone. I will repeat this with much
higher concentrations.
I will see if this method works at all....
ps. have you got any good reference about this reaction (procedure) ? I mean not just quotation from a book, but journal reference.
[Edited on 14-2-2009 by kmno4]woelen - 14-2-2009 at 05:29
I have no reference at all. I found out myself. The prep is very easy and you can also easily show that the red/brown stuff contains acetone (on
heating, it turns into anhydrous CuCl2 and you get a strong smell of acetone).Jor - 14-2-2009 at 09:29
If you want some anhydrous CuCl2, do the following:
-prepare the dihydrate, for example by dissolving the oxide or basic carbonate in hydrochloric acid, and evaporating/crystallising.
-to the CuCl2.2H2O (With probably some mixed HCl in it, giving it the green color), add concentrated sulfuric acid or heat SLOWLY. This will get you
the anhydrous form in both methods. The former is problematic, because I'm not sure how to remove the sulfuric acid completely from the CuCl2, the
latter is problematic as you have to do it really slowly, or it will just hydrolyse forming a basic chloride and HCl. This therefore takes time.panziandi - 14-2-2009 at 13:38
Jor wouldn't your former method make copper sulphate and release hydrogen chloride? You could put the CuCl2.xH2O in a shallow dish OVER H2SO4 in a
dessicator, perhaps that would work particularly if its under vacuum and left for a while.
(could always use SOCl2 too... hehe )Jor - 14-2-2009 at 14:02
Thionyl chloride ...
Im not sure if sulfuric will convert the CuCl2 to HCl and CuSO4. Are you sure? This happens with NaCl, but I remember that doing the experiment with
CuCl2.2H2O (I once made 1g of it, was not very pure, because I didn't know how to make it in early days ), that a brown powder formed, with hardly any bubbling. It was about 15C.
More copper
Paddywhacker - 1-3-2009 at 05:20
Cupric tartrate is soluble in excess alkaline tartrate to give Fehling's Solution, but is Cupric Tartrate itself soluble in water? What about Cupric
Hydrogen Tartrate? What about the citrates?
What is this white material? It looks as if a copper(I) compound is formed. Or is this white color an artefact of the photography?Paddywhacker - 1-3-2009 at 11:44
Quote:
Originally posted by woelen
What is this white material? It looks as if a copper(I) compound is formed. Or is this white color an artefact of the photography?
The mixtures were all initially black from the copper oxide, but reacted with the acids when the vials were heated in a pressure cooker for an hour.
The light coloured material is the product, and you can guess its solubility from both the amount of deposit and the blue colour of the water.
The copper tricitrate may have all reacted, and it may be that dark brown colour, or the colour might arise from unreacted copper oxide.
The colours are genuine. I don't think the copper was reduced to Cu(I), because the same-looking precipitates can be made by mixing solutions of
CuSO4 with solutions of tartrates or citrates.... but I could easily be mistaken.
[Edited on 2-3-2009 by Paddywhacker]woelen - 3-3-2009 at 10:33
Another explanation might be that anhydrous precipitates are formed, in which the copper-ions are not coordinated to water, but to citrate/tartrate.
But this does surprise me somewhat, because normally these coordination complexes are deep blue. It might also be that the precipitates simply are
from the acid itself, or some mix of the the acid and a copper tartrate/citrate.
It is interesting to know what happens when a lot of water is added to the precipitates. Do they dissolve and what color do you get?woelen - 6-3-2009 at 12:27
I tried the experiment with a large excess of citric acid and copper(II) oxide:
The result of my experiment differs a lot from the one, reported by Paddywhacker. I end up with a nice green crystalline powder of cupric citrate,
somewhat resembling cupric acetate, but a little lighter and greener.Paddywhacker - 6-3-2009 at 12:46
A photo would be nice to see the comparison. You found the CuO not so easy to dissolve. I solved that by capping the vials and heating them in a
pressure cooker for an hour... which was probably excessive.
That is one of the intreaguing things about these compounds.
It seems that with excess citrate, or excess tartrate, you get complexes that are different from the plain salt.
The experiments I setu up tried to get the molar ratios right to give me exactly Cu(H2Citrate)2, Cu(HCitrate), and Cu3(Citrate)2 ... and the
corresponding tartrate salts.
My idea had been to find a sparingly soluble copper salt and grow nice big blue crystals of it via the hydrothermal process, but I never set up the
gear.
[Edited on 7-3-2009 by Paddywhacker]kmno4 - 11-9-2009 at 14:20
Just something nice - luminescence of Cu4I4Py4
Preparation is extremely simple* and effect is interesting: seledine luminescence of this complex under UV (BLB lamp).
When complex loses some pyridine in the air, it becomes completly odorless (in my case it is colored gray because some impurities) and luminescence
changes to yellow-orange.
Some pictures below (UV pictures are bad quality but I did not want to play with camera settings). Yellow-white powder is freshly prepared complex,
gray powder is dried for a few days in the air powder.
This complex can be recrystallized from DCM, acetone, acetonitrile... etc
Very nice kmno4! So the complex is quite unstable? Also if stored in a closed container, over pyridine vapour?
Rather stable, but stable form seems to be this with yellow luminescence.
Partial dried complex looks funny, with yellow-blue luminescence:
In a daylight it is practically white.woelen - 12-9-2009 at 06:09
Kmno4, this is a great find. Beautiful! Of course I could not resist trying it myself. I have all required chemicals.
I once made some CuI (a long time ago and I still have it). It is an off-white solid. I think pure CuI is white like snow, but mine is good enough for
most experiments.
I dissolved some KI in water and boiled for a while to get rid of dissolved oxygen. To this I added some CuI. Only part of it dissolves, I obtained a
pale yellow solution, almost colorless with an off-white precipitate at the bottom.
When this precipitate of CuI under a solution of KI/CuI is illuminated with a black light no special effect can be observed.
In another test tube, I took 3 ml of acetone and added four drops of pyridine (appr. 0.2 ml). The pyridine dissolves in the acetone without problem.
Next, I added this solution to the test tube, in which there is the solution of KI+CuI and solid CuI at the bottom. At first all material seems to
dissolve, but almost immediately, a white precipitate is formed. The pale yellow color becomes even paler, and what remains is a white and fairly
compact precipitate under a nearly colorless liquid.
This is a picture of the material in a dimly lit room.
In the picture you can see a black light tube. Next, I switched on the tube. It hardly gives off any visible light and the digital camera also makes
little of the UV-light. The solid material, however, beautifully lights up, showing yellow fluorescence.
woelen - 12-9-2009 at 06:23
Here I have a few pictures of the same material, after the precipitate has settled at the bottom. First a picture in a beam of sunlight in the
otherwise dimly lit room and then the same precipitate placed somewhat above and behind the blacklight tube.
Here is the material in sunlight:
And this is near the blacklight:
kmno4 - 12-9-2009 at 11:51
Another complex Cu2I2, this time with morpholine.
In daylight this is white powder with weak green shade, odorless.
In UV it showes orange luminescence (on the picture).
These complexes were described long time ago. My fellow has sent me a picture of pyridine complex recently and interested me with this. Some time ago
I tried urotropine complex of CuCN but because of unknown (to me) reasons it did not show any luminescence (but should do).
Many amines can be used for preparations these nice things - have a nice experimenting itchyfruit - 12-9-2009 at 14:02
This all looks rather interesting,is copper thiocyanate as toxic as it sounds?woelen - 12-9-2009 at 14:14
Thiocyanates are not really toxic and do not have to do much with cyanide. Potassium- or sodium thiocyanate and ammonium thiocyanate are easy to
obtain chemicals and they can be purchased without any eyebrows raised. They are used extensively in hobby photography, darkroom processes.ammonium isocyanate - 12-9-2009 at 14:18
Depends. The thiocyanate anion is much less toxic than the cyanide anion, so much so that back in adv. chem in 10th grade they let us work with it.
Copper ions are somewhat toxic, but again nothing compared to cyanide. Sigma Aldrich lists it as Xn and N (harmful and hazardous to the environment),
but they don't list any serious safety phases, other than dispose properly, keep away from food, etc.
In summary, I wouldn't be too worried about working with it.
Edit: whoops, looks like I got beaten to it.
[Edited on 12-9-2009 by ammonium isocyanate]itchyfruit - 12-9-2009 at 14:40
Cool !!! i'm quite interested in copper salts myself and i think i have some ammonium thiocyanate.JohnWW - 12-9-2009 at 15:32
The soluble (mainly alkali-metal and ammonium) and volatile (capable of being inhaled, particularly light organic esters) thiocyanates should be more
toxic (and more immediately so) than the heavy-metal ones, due to the latter being much less soluble and so less reactive. The latter may be
decomposed by stomach acid if ingested, but this would be a relatively slow process (enabling detoxification at the same time) compared to the rapid
effects of soluble ionic thiocyanates.
[Edited on 13-9-09 by JohnWW]woelen - 12-9-2009 at 22:44
With the latter I disagree, e.g. lead thiocyanate and mercury thiocyanate are MUCH more toxic than e.g. potassium thiocyanate, simply because of the
toxic metal ion, despite their low solubility. Many insoluble thiocyanates easily dissolve in stomach acid (contains dilute HCl), due to formation of
HSCN and the consequent release of the toxic heavy metal ion. Jor - 13-9-2009 at 02:09
Talking about mercury thiocyanate, something really fun can be done with this substance:
Not recommended though due to the mercury vapours released, wich are bad for health and environment.itchyfruit - 13-9-2009 at 02:39
Cool, any idea what you end up with? compound wise, i know it looks like a big piece of tree bark. DJF90 - 13-9-2009 at 04:45
Thats the basis for the "snakes in the grass" experiment (basically that combined with an ammonium dichromate volcano). Quite interesting, although
I'm not sure what the products of the reaction is.woelen - 8-11-2015 at 13:20
I updated my experiments with copper/pyridine/iodide. I now was able to make the complex with green/cyan fluorescence and again made the complex with
yellow fluorescence and I made two durable display samples:
woelen do you have copper cyanide? mixing potasium cyanide with copper sulfate makes copper cyanide but also cyanogen and cyanogen is toxic.will
cyanide dissolve copper sponge as easy as dissolving silver sponge?copper cyanide is supposed to be a pale yellow color which precipitates after
adding sodium sulphide to copper sulfate but also makes cyanogen,i think.
[Edited on 11-8-2015 by cyanureeves]woelen - 9-11-2015 at 23:56
I do not have copper cyanide. I see no relation with the subject of my previous post. Can CuCN be used to make interesting fluorescent complexes? cyanureeves - 10-11-2015 at 15:45
dont know if if copper cyanide glows with a fluorescent light but i know scorpions do.careysub - 10-11-2015 at 16:12
I do not have copper cyanide. I see no relation with the subject of my previous post. Can CuCN be used to make interesting fluorescent complexes?
Possibly he is thinking of the use of potassium cupric cyanide in the luminol clock reaction?cyanureeves - 10-11-2015 at 16:54
no i just saw copper complex and jumped right in because i was curious as to how copper cyanide was made.this thread is about glowing copper only and
i just know about the glowing scorpions.at least i know they glow under a black light but fluorescent light might be different maybe.a luminol clock
huh?well i'll just go over to google to see what all that is about. woelen - 20-11-2015 at 07:15
With morpholine, also interesting fluorescence is exhibited. I tried in a similar way as with pyridine, add a solution of morpholine in water to a
suspension of CuI in a solution of KI. A pale yellow precipitate is obtained, which gives bright orange light when irradiated with black-light.
I also tried with 2,2-bipyridyl, but this does not give fluorescence. A nice, bright orange/brown precipitate is fomed though. On dilution with water,
part of this dissolves, but the color of the remaining material becomes even brighter, red/brown.
).
)
), that a brown powder formed, with hardly any bubbling. It was about 15C.
