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jamit
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When I find time next week I'll show the chemicals used and pictures of each step.
Btw do you find it hard to take pictures of green copper II chloride solution? It always looks blue when you take the pictures when in fact it's deep
dark green solution. Is it just my bad camera.. I'm using iPhone 4.
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jamit
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@Woelen
This might be an important factor which I forgot to mention but I did the experiment in my garage and its about 3-5C. So everything was done in a
cold environment. And the solution of copper II chloride was concentrated without heating. I just added copper carbonate into concentrated HCI until
it stoppped fizzing and allowed it to fully react by leaving it overnight. Next day I filtered twice and got a deep and dark and transparent green
solution of copper II chloride. You can tell visually that it was concentrated by the deep dark color... if it was dilute it would be a blue with a
slight green solution.
Anyway, I wanted to update what I did and hope you and others can duplicate my results.
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woelen
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I have now tried one experiment with labgrade chemicals.
I took reagent grade 35% HCl and added basic copper carbonate until the solution stopped fizzling. At that point I obtained a very dark green solution
and also some solid green material. I had to add some water and had to heat the liquid in order to obtain a clear solution. To this (still warm)
solution I added more copper carbonate until the solution became just a little turbid and it no more reacts anymore. To this I added then a single big
drop of 30% HCl in order to make it really clear.
This clear solution was put aside in a room at 12 C (colder than that is not available to me anymore in this time of year) and after one night I had a
lot of needle-like crystals at the bottom and still a fairly dark green solution above it. Everything was green though.
Before proceeding further I now want to ask you a question. What if I add just enough water to dissolve all these needle-like crystals and then put
the resulting solution in a fridge (keeping it at 6 C or so)? I now see that my solution really is too concentrated.
---------------------------
Another thing you could try to do is dissolve some of your 'normal' green CuCl2.xH2O in as little of water and then let it crystallize again. Does
this yield green crystals or blue crystals?
[Edited on 31-3-12 by woelen]
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jamit
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@woelen
I have done what you reported above and I agree that you do get green needle like crystals.
However, don't heat or boil down the solution. Do the experiment in the coldest temperature possible. Add cuco3 into conc. HCl until it no longer
fizzes then filter any insolubles. this should create a conc. solution of copper II chloride. Now make sure there are no visible particles at the
bottom... That's why I filtered seceral times. Once this is done put it into the fridge and wait. Several days later, I noticed these crystals at
the bottom of the green solution... Once they were removed from solution, I noticed that they were blue with a slight green.
Your results do make me wonder what I might have done differently... I mean this is a simple inorganic reaction, although the water ligands makes
things interesting.
[Edited on 1-4-2012 by jamit]
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jamit
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Here's a comparison of two different color crystals of copper II chloride. the blue crystal is what I harvested first.
Several days later I harvested another batch of crystals from the same solution of copper II chloride and the results are different -- green crystals
but not needle like.
Question? did I get a different hydrate of copper II chloride?
Is it not the water molecules ligands or the lack thereof that is causing the different colors? This is really puzzling!
[Edited on 1-4-2012 by jamit]
[Edited on 1-4-2012 by jamit]
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blogfast25
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Quote: Originally posted by jamit  |
Question? did I get a different hydrate of copper II chloride?
Is it not the water molecules ligands or the lack thereof that is causing the different colors? This is really puzzling!
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The contrast is striking and puzzling.
Yes, the degree of hydration could be a possible explanation, also because in general different types of hydrates can form at different temperatures.
But it’s strange that only the dihydrate ever seems to be mentioned and no others.
Chloride and water ligands are responsible for the different colours of the copper (II) cation in solution: tetrachloro cuprate
(CuCl<sub>4</sub><sup>2-</sup> is green,
Cu(H<sub>2</sub>O)<sub>6</sub><sup>2+</sup> is blue (add NaCl or HCl to CuSO4 to see the colour change). Whether
that is the cause of the difference in colour of your crystals remains to be determined.
Hydrates with a higher number of crystal water molecules tend to form at lower temperatures. Is it possible that your blue product is a higher
hydrate, with water molecules replacing at least one of the chloride ligands? If the degree of hydration plays a part then determining the water
content of the blue and green crystals would give clues. Copper (II) chloride is easy to dehydrate without any special precautions (to prevent
hydrolysis) required.
The shape of the crystals could also provide clues: if a different hydrate is at play here, then both products may not have the same crystal structure
(lattice). Try isolating a 'large' crystal of each, wash with a little iced water (to remove smaller crystals), then look at them with a magnifying
glass. Do they look to have the same shape?
[Edited on 1-4-2012 by blogfast25]
[Edited on 1-4-2012 by blogfast25]
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Nathaniel
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If you ask me the blue crystals are almost 100% CuSO4 pentahydrate crystals, judging by the colour and of course crystal structure (some of
the blue crystals you have there are quite nicely shaped and they have the same structure as CuSO4*5H2O). It is very unlikely
that some weird copper chloride crystals would crystalize out at those conditions.
So I suppose some sulfate ions got into your solution and because copper sulfate is less soluble than chloride it crystalized out before
CuCl2 (it's crystals are needle-shaped and it can be seen that your green crystals are not as nicely shaped as blue ones).
Do you think there might be sulfate ions in your solution? Did you filter and rinse your copper carbonate well enough before adding HCl?
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jamit
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@Nathaniel
The thought that those crystals are copper sulfate has crossed my mind but that doesn't seem possible. If you read what I did, I'm not sure where the
copper sulfate would have been formed.
Sulfate ions might have contaminated my production of copper carbonate, as sodium sulfate. But that would have been in small amount, as I washed the
copper carbonate at least 10x with distilled water.
When I added HCI, all the copper should have formed into the chloride. I suppose there might have been small amount of sulfate ions (but that would
be very very small amount) that might have reacted with the copper ions but that does not explain how the crystals which are forming right now in my
beaker keeps getting bigger?
Anyway, thanks for sharing your thought. I'll test the purity of my copper carbonate.
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Nathaniel
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Oh in that case I'll definately try making those blue crystals myself next weekend... I originally noticed this topic because of the pictures (I like
coloured crystals) so thanks for the idea 
I have quite a lot copper chloride so I'll make the conc. solution simply by dissolving it in cold water...
Another guess would be that other ions are interfearing with the crystalization process... I would never really thought so, but only yesterday one of
my teachers (for "Structure of solids" subject) mentioned that addition of urea can cause NaCl to crystalize in alum-like structure instead of
expected cubic  .... I'm suspecting carbonate ions since any heating is
obviously preventing the formation of blue crystals and CO2 is indeed driven out of the solution by heating... So I was thinking of saturating conc.
CuCl2 solution with CO2, but then remembered that I could simply dissolve CuCl2 in some "fizzy water", keeping the solution cold (ice bath)
I don't know whether this could be possible at all or is it another wrong guess (sorry for doubting your carbonate quality jamit ) but I'll try it anyway
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CHRIS25
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| Quote: Originally posted by jamit | I agree that the green crystals are needle shaped. But have you ever tried to do a quick cold wash with distilled water and then quickly dry it?
You get blue crystals like the one you see on the far right.
So here's my three forms of copper II chloride (from left to right): copper II chloride (+chloride), anhydrous copper II chloride, and copper II
chloride (washed in water).
If you dissolve the green crystals, it turns the water solution blue. So there must be a way to get blue copper II chloride crystals. I just don't
know how? Can someone help?
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Hi there, I have just produced Blue copper chloride by accident. Are you still wanting to know how?
[Edited on 7-4-2012 by CHRIS25]
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blogfast25
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| Quote: Originally posted by CHRIS25 |
Hi there, I have just produced Blue copper chloride by accident. Are you still wanting to know how?
[Edited on 7-4-2012 by CHRIS25] |
Yes. Please provide evidence of what you're asserting.
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CHRIS25
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Actually I was too hasty, my precipitate should be copper carbonate from the reaction that I posted as a separate topic a few minutes ago "Why so many
different colours fro same chemical". I confess confusion here, maybe you could confirm my reaction's results from this post and if it is copper
chloride then I will provide more accurate details. Thanks
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woelen
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I now have a solution of copper carbonate in conc. HCl standing here for over 1 week now, but still no blue crystals. I do get a very slight amount of
precipitate, but it is green. I think that green material is basic copper chloride. It is insoluble in water, but dissolves at once, as soon as I add
a tiny amount of HCl.
Might it be that the blue material is due to the presence of impurities? The material I use is general lab reagent grade and for most practical
purposes that is more than pure enough. Jamit made his copper carbonate from sodium bicarbonate and copper sulfate. I know that many copper
precipitates suffer from coprecipitation of other ions. It might be that the material of Jamit contains sodium ions and sulfate ions and then the blue
crystals might be copper chloride with some places in the lattice replaced by sodium and/or sulfate. Even a small fraction of foreign ions may lead to
a completely different appearance. I will try my experiments with a small amount of added Na2SO4 (appr. 5% of the amount of copper carbonate used) and
see what results I get. Another week of waiting ...
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Arthur Dent
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Aah, the miracle of copper chemistry... the chloride is particularly intriguing. A few months (years?) ago, I dissolved some telephone wire in
concentrated HCl and I was surprised to obtain what I could described as a dirty brown, almost copper-colored solution. Addition of a bit of 9%
Hydrogen Peroxide turned that solution almost instantly into a beautiful deep emerald green color. I have kept the solution in an airtight
Schott-Duran bottle and will eventually try to evaporate and crystallize it.
At the time of my experiment, our excellent colleague Woelen explained to me the transition from Copper (I) to Copper (II), and has described that
reaction in much detail on his excellent website. Here's my original thread:
http://www.sciencemadness.org/talk/viewthread.php?tid=14940#...
So the crystalized Copper Chloride color variations are also quite interesting. I'll have to pull out that CuCl2 bottle and experiment with the
resulting crystals, it'll be a good opportunity to test out the macro function of my camera! 
Robert
--- Art is making something out of nothing and selling it. - Frank Zappa ---
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jamit
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@ Woelen @Nathaniel
I finally did some test and further experiment and I found out what my problem was! Thank you all for all your hypothesis and stimulating thoughts.
As suggested by Nathaniel and by Woelen's last text above... the problem was impurities... specifically impurities in my copper II carbonate made by
sodium bicarbonate and copper sulfate. I must have done a bad job of removing the sulfate ions from the copper carbonate.
I did several tests.
As suggested by Woelen, I heated the blue (supposedly copper chloride) crystal and it turned white. Nathaniel suggested by the pictures that my
copper chloride looked more like copper sulfate pentahydrate based on color and crystal patterns. So I heated what I thought were copper chloride.
If it was copper chloride it should have turned brown into its anhydrous form, but instead it turned white, which is the anhydrous form of copper
sulfate. This was my first clue that what I had was not copper II chloride crystals but maybe copper sulfate. So how did sulfate ions get into my
solution?
This lead me to think that my copper carbonate was contaminated with sulfate ions. So I added my entire batch of copper carbonate (about 400g) which
I made before and added about 1000ml of water and allowed it to settle to the bottom. I took some of the water solution to test for sulfate ions
using barium chloride (with a few drops of hydrochloric acid). It produced a film of white ppt of barium sulfate indicating a sufficient amount of
sulfate ions contamination!!!!
I must have messed up in my synthesis of copper carbonate... I needed to wash it more thoroughly... which is what I'm doing now.
So here's what I think happened. When I added copper carbonate to hydrochloric acid to make copper II chloride, it was contaminated with sodium
sulfate which I failed to remove when I made copper carbonate.
So what made was copper II chloride mostly and some copper sulfate. And since copper sulfate has a lower solubility in cold water than copper II
chloride, it crystallized out in my cold garage. The solution was green but the copper sulfate which crystallized at the bottom was blue once I
removed it from the solution.
Lesson learned: Make sure to purify the synthesis of your chemicals and test them for purity!! I'm going to be doing that with all my chemical
synthesis from now on!
In conclusion, I want to thank everyone who participated in his topic especially Woelen, Nathaniel, and blogfast, etc. I love Sciencemadness... thank
guys for all the help.
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woelen
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Precipitation reactions in fact are not the preferred types of reaction when it comes to synthesis. The school textbooks always present precipitate
reaction as nice well-defined reactions, but in reality there always is coprecipitation which leads to considerable impurities of your product.
Washing helps to remove some of the coprecipitated ions, but in many cases the ions simply are part of the precipitate and you cannot remove them.
Especially if the precipitate forms at once, there is a lot of coprecipitation. There also are precipitating reactions in which the precipitate forms
slowly (minutes). In such cases, the precipitate usually is more compact and looks a little bit more crystalline (miniature crystals). This kind of
precipitates has better purity and it usually is produced when slightly soluble salts are formed (e.g. mixing dilute solutions of NaClO4 and KCl,
which leads to slow formation of KClO4 and a 'snow' of glittering crystals from the solution).
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mr.crow
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Time to revive an old thread.
I made some CuCl2 starting from 1 mol CuSO4, turning it into carbonate with Na2CO3, filtering, washing, then adding 2 mols of HCl, then boiling down
and cooling. I filtered one crop of green CuCl2 small needle like crystals so far
The mother liquor was diluted with some water due to washing then left to stand. Some beautiful crystals appeared! They look like CuSO4 but are a
completely different color.
What are they?
Double, double toil and trouble; Fire burn, and caldron bubble
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jamit
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They are copper sulfate with chloride impurity. Heat it to see if it turn brown or white? If brown, mainly copper II chloride. If white, mainly
copper sulfate.
[Edited on 23-7-2012 by jamit]
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mr.crow
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I heated a tiny bit in a tube and it turns greenish white. So mostly CuSO4. Its strange I made CuSO4 of a different color
Double, double toil and trouble; Fire burn, and caldron bubble
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triplepoint
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But for the slight green tinge, this looks identical in shape and color to the CuSO4 crystals I have seen. (I know CuSO4 can be grown into big,
beautiful crystals, but it starts off with crystals very much like these if you simply evaporate off a small volume of solution.
Can't write much now, I am in a vacation cabin with painfully slow internet. Reminds me of my dial-up days. 
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Rogeryermaw
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| Quote: Originally posted by mr.crow | | I heated a tiny bit in a tube and it turns greenish white. So mostly CuSO4. Its strange I made CuSO4 of a different color |
when i actually used copper sulfate as root killer, outdoors, and after being exposed to the sun and humidity for a day or two, that is the appearance
it took. blue-greenish-white and powdery.
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Vargouille
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I believe he was referring to the light color of the CuSO4 crystals. The technical grade CuSO4 I have from a root killer is a much darker blue.
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tetrahedron
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i'm trying to synthesize some CuCl2 by fractional crystallization of CuSO4 + NaCl. what i noticed so far is that the blue
species predominates at lower temperature and low concentration. a partially dissolved stoichiometric mixture of CuSO4 + 2NaCl in some
water became intensely green after heating, and increasingly so as the water boils off, whereupon the suspended solids increased in size and took on a
whitish color (arguably Na2SO4). cooling and/or dilution cause the color to shift again to blue. thus the
Na2SO4 has to be removed while hot/concentrated. the remaining sulfates (probably a mix of Cu and Na salts) should crystallize
next, while the chlorides (mostly CuCl2) stay in solution.
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mr.crow
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The intensity of the green color does increase with temperature. I also got those CuSO4 crystals to form upon cooling. If you get it to work it might
be a better method than the CuCO3 route.
Oh yeah, an update. After leaving the CuCl2 in a desiccator for many weeks with NaOH and CaCl2 the small green crystals turned light blue. Success!
Double, double toil and trouble; Fire burn, and caldron bubble
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tetrahedron
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after several boiling>cooling>decanting cycles i finally got a solution that stayed green and had only green crystals in it. no CuSO4
crystals were observed. however, the soln didn't crystallize nicely. it emits an acid smell, but i see no needle-like crystals. the anhydrous form is
brown, as expected.
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