Sciencemadness Discussion Board - copper II chloride crystals

copper II chloride crystals

jamit - 31-1-2012 at 02:45

I find copper II chloride fascinating, particularly, its various stages of hydration.

I made copper chloride using copper carbonate and hydrochloric acid but here's my dilemma: Its so hard to make transparent crystals out of them. I was able to get transparent blue crystals out of them once but I haven't been able to duplicate the result since.

After filtering to get a clear solution of copper chloride. I allow it to evaporate slowly at room temp and all I get is this bright green power at the bottom.

If you wash this green powder in water, for a moment, it turns blue, and you can quickly separate it and dry to get blue crystals. It's not as blue as copper sulfate but clearly the copper chloride can have a transparent blue color.

Has anyone had any similar experience working with copper chloride?

I've dried the green copper chloride in an oven and obtains the anhydrous brown color.

So my experience has been that copper chloride exist in at least three forms:

1. Anhydrous copper chloride (brown) - made from hydrous copper chloride (green)

2. various hydrous forms of green copper chloride - I've read that this green is the result of excess chloride ions (see wiki)

3. blue crystalline (transparent) copper chloride that was washed with distilled water.

I would love to hear from anyone who has actually done some work on copper chloride and if you could share your findings... I would appreciate it very much. thanks.

[Edited on 31-1-2012 by jamit]

[Edited on 31-1-2012 by jamit]

MrHomeScientist - 31-1-2012 at 06:23

I did one experiment a while back to test wiki's claim of the green color, and found that it is in fact true. I made copper(II) chloride basically the same way you did, going from CuSO4 -> CuCO3 -> CuCl2. The final CuCl2 solution was a deep green. I then took a small amount of the solution and diluted it with some water, and the color changed to light blue as the concentration decreased. Then, I added a small amount of ammonium chloride to this solution and stirred to dissolve, and the green color returned!
I made another light blue solution, then poured both that and the ammonium chloride-ed green solution onto watch glasses to let them evaporate. I now have two batches of copper chloride crystals, one green and one light blue!

LanthanumK - 31-1-2012 at 09:04

When I dry a highly acidic solution of copper(II) chloride, it dries as green crystals which smell like HCl.

bbartlog - 31-1-2012 at 17:15

It's possible that the blue form is some sort of nonstoichiometric oxychloride. CuCl2 is a fairly weak lewis acid but under the right conditions I think you could still see a little bit of CuCl2 + H2O -> Cu(OH)Cl + HCl.

LanthanumK - 1-2-2012 at 06:50

I think the green form is an acidic species of CuCl2, and the blue salt is the normal salt. Hydrolysis would probably result in an insoluble residue, which is not seen with blue CuCl2 (I possess both forms).

jamit - 4-2-2012 at 15:45

Correct me if I'm wrong but the green crystals of cucl2 (more acidic) are long needle like crystals, whereas the blue colored cucl2 is more cubical... Which is quite surprising. Does any concurr with my description of the crystals structure of copper II chloride?

LanthanumK - 8-2-2012 at 06:49

All copper(II) chloride crystals, AFAIK, are needle-shaped.

The picture is a comparison of green copper(II) chloride with normal copper(II) sulfate.

Copper(II) chloride and copper(II) sulfate small.jpg - 43kB

jamit - 8-2-2012 at 16:13

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?

Poppy - 8-2-2012 at 18:10

Smash de green crystals very finely. then add a small ammount of water, keep the salt and the water in an near to equilibrium by heating the mikxture until near total dissolution, this might syntherize the crystals into bigger ones, but this would take time, and probably pH control (anionic strenght since thats what differs the colors)

woelen - 8-2-2012 at 23:45

Pure copper(II)chloride dihydrate is blue/green, more like cyan. It has needle-like crystals.

The picture shows a commercial sample of reagent grade CuCl2.2H2O.

The grass-green compound is an acidic compound, which contains HCl as well. It contains complexes like HCuCl3.xH2O and these cause the green color.

jamit - 9-2-2012 at 13:03

Thanks Woelen! I was waiting for your input on this subject knowing that you have done a lot of work on copper based compounds.

But could you comment on what you think I have above in the three vials. So the green vial on the far left is HCI contaminated and has various complexes? What about the middle vial? Is it anhydrous of the vial on the left with HCI contaminants?

And more importantly, what do you make of the vial on the far right? These are not needle like crystals. When you add the green crystals to water they turn the solution blue. Have you ever tried a quick wash of the green crystals and then dried them on a paper towel? They turn blue.

My question is, 'what do I have on the far right vial since I've wash it with water and the crystals are blue'?

Your commercial sample of blue/green crystal -- how do you make that at home?

I don't think you have any experiment on your website on copper II chloride synthesis or do you?

anyway, i would love to hear back from you on your thoughts! thanks.

zoombafu - 9-2-2012 at 14:51

I do believe that the middle is the anhydrous CuCl2, as it is the right color and the crystals look right from my experience. However I have not made the needle like crystals, those must be the super pure from. To get them you would (most likely) have to do a recrystallization. Ill try this in the near future, and post my findings on this thread.

[Edited on 9-2-2012 by zoombafu]

jamit - 10-2-2012 at 00:54

Thanks Zoombafu for your input. However, I'm looking for answers from those who have actually attempted to make copper II chloride and its anhydrous and hydrous forms.

I look forward to your post when you do the recrystallization. Making copper II chloride is a simple inorganic procedure but its not easy making pure copper II chloride crystals... at least not the purity that Woelen showed in his picture.

woelen - 10-2-2012 at 01:30

The left vial in your picture almost certainly is quite pure CuCl2.2H2O with some HCl in it. It is a very well-known color for me. If I look at the picture (assuming color balance on your camera is correct) , then it must be quite pure already, the amount of HCl in it is not large. Just try heating it a little bit. An acrid smell of HCl almost certainly will be given off.

The brown material definitely is the anhydrous compound. I have made that as well. Add a little of the brown material to water. Does it dissolve to a green or blue solution? Is the solution 100% clear or somewhat turbid? If the CuCl2 anhydrous is prepared by heating, then it might be that there is a small amount of oxychloride or even oxide in the sample, due to driving off HCl. A carefully prepared sample, however, will dissolve without any turbidity.

I cannot make anything of the bright blue compound you have. I have never seen such a compound when based on copper chloride. My sample of copper chloride becomes green when a drop of water is added to it. Concentrated solutions are green (chloro complexes) and dilute solutions are blue (aqua complexes). Just a few suggestions:
1) Heat a small sample of the bright blue compound. What color do you obtain? Brown like the anhydrous CuCl2?
2) Dissolve some of the bright blue compound in water. Do you get a clear solution?
3) Add some NaOH solution (just a few drops) to your solution, prepared in (2). Do you get a bright blue precipitate, or a dirty green one? The latter is obtained if you have high concentration of chloride in solution.
If indeed the blue compound is CuCl2 as well, then I have learned something new.

jamit - 10-2-2012 at 23:44

Thanks woelen for your valuable input! I'll go and test the blue bottle of what I think is copper II chloride. I'll report back.

zoombafu - 11-2-2012 at 23:27

Ok I think that I have made copper II chloride by chlorinating copper, with H2O2 acting as a the oxidizing agent:

Cu + HCl + H2O2 --> CuCl2 + H2O

I will be posting a full lab write up eventually, but here are my initial findings. I was able to obtain the needle like crystals, however I was unable to get a good picture, Ill post one when I make my final purification, and once I scraped them off of the glass the crystals got jumbled and broken (so I don't think I got the 'true' needle crystal). The First picture Below is the sample of CuCl2.2H2O. The second is the anhydrous form. The third is a highly concentrated aqueous solution. And the fourth is a lightly concentrated solution. The color of the solution also translates down to the color of the crystal, which is why I think your samples are different colors.

Kola - 21-2-2012 at 06:29

I've once prepared (impure) CuCl2 by electrolysis of a mixture of NaCl and Al2(SO4)3 using Cu electrodes. The aluminium sulphate i added was to ensure the solution remains acidic. The CuCl2 formed was first green, but later as some Cl- was escaping as CuCl (a side reaction), the solution turned blue

zoombafu - 24-2-2012 at 23:04

I posted this synthesis on my blog if you want to check it out.
http://thehomechemist.com/blog/cucl2synthesis/

jamit - 29-3-2012 at 23:25

Quote: Originally posted by woelen

The left vial in your picture almost certainly is quite pure CuCl2.2H2O with some HCl in it. It is a very well-known color for me. If I look at the picture (assuming color balance on your camera is correct) , then it must be quite pure already, the amount of HCl in it is not large. Just try heating it a little bit. An acrid smell of HCl almost certainly will be given off.

The brown material definitely is the anhydrous compound. I have made that as well. Add a little of the brown material to water. Does it dissolve to a green or blue solution? Is the solution 100% clear or somewhat turbid? If the CuCl2 anhydrous is prepared by heating, then it might be that there is a small amount of oxychloride or even oxide in the sample, due to driving off HCl. A carefully prepared sample, however, will dissolve without any turbidity.

I cannot make anything of the bright blue compound you have. I have never seen such a compound when based on copper chloride. My sample of copper chloride becomes green when a drop of water is added to it. Concentrated solutions are green (chloro complexes) and dilute solutions are blue (aqua complexes). Just a few suggestions:
1) Heat a small sample of the bright blue compound. What color do you obtain? Brown like the anhydrous CuCl2?
2) Dissolve some of the bright blue compound in water. Do you get a clear solution?
3) Add some NaOH solution (just a few drops) to your solution, prepared in (2). Do you get a bright blue precipitate, or a dirty green one? The latter is obtained if you have high concentration of chloride in solution.
If indeed the blue compound is CuCl2 as well, then I have learned something new.

I've confirmed that there is a blue crystalline solid form of copper II chloride. Just last week I added HCl to copper II carbonate until fizzing stopped and filtered and got this deep green solution. If you let it evaporate you will get the green solid. However, if you concentrate it to supersaturation and then filter again to get rid of any undissolved solids and allow this to stand in a cold fridge, you will get these beautiful blue crystals from the deep green solution.

I've done this twice and I still don't understand why sometimes it produces green solids and others it crystallizes into a blue diamond shaped crystals??
See the picture below:

[Edited on 30-3-2012 by jamit]

woelen - 30-3-2012 at 00:32

This is something I am going to try. I really am surprised by this. Nowehere on the internet one can find info on these blue crystals. Did you boil down the liquid? Please give some more detail on the exact process and then I can try to reproduce your results.

jamit - 30-3-2012 at 03:08

Quote: Originally posted by woelen

This is something I am going to try. I really am surprised by this. Nowehere on the internet one can find info on these blue crystals. Did you boil down the liquid? Please give some more detail on the exact process and then I can try to reproduce your results.

Thanks Woelen for your quick reply. I know through your website that you have done a lot of copper related experiments. What I did was simple. I used copper II carbonate, which I made with copper sulfate and sodium hydrogen carbonate. I added this to a 31% hardware brand of HCI until it stopped fizzing. I allowed the reaction to sit for a day. Next I filtered twice using a fine filter paper to remove any copper carbonate or other insolubles. The solution I collected was about 250ml, deep green and transparent solution, which seems quite concontrated. When I allowed it to sit for a day, you can see fine particles or dust that has settled to the bottom -- it was missed by the filter paper. I filtered again but only the clear section of the solution and discarded the dust/insoluble impurities.
Next, I transferred the solution to a clean beaker and left it in my garage for several days. Its winter where I live so its about 5C. I checked yesterday and notice diamond like crystals. I could have allow it to grow even more but I decanted the solution and collected the crystals to show off picture above. I'm going to use the seed to grow bigger crystals. I'll post the picture after it gets bigger.

After your previous comments, I just assumed that I must have done something wrong to get "blue" crystals of copper II chloride and that maybe I confused it with copper II nitrate. But after this result, I'm sure that copper II chloride also exist in this blue crystal form. The solution from which it came, is deep green, from which I expected green solids.

I'm going to try this experiment again in the next several weeks/months... at least several times. Like you, I've this fixation with copper compounds. Give it a try... in fact I'm surprise that you've never got this result. Good luck

[Edited on 30-3-2012 by jamit]

jamit - 30-3-2012 at 03:23

More pictures of my results

woelen - 30-3-2012 at 10:15

I'll try the experiment with commercial grade CuCO3 and 30% HCl, just to assure that no impurities like sodium ions or sulfate ions are in the compounds I use. As soon as I have results I'll let you know.

Dennis SK - 30-3-2012 at 10:27

Beautiful pictures in this topic!
Thanks for sharing.

blogfast25 - 30-3-2012 at 13:02

This quite interesting. An ambiguous colour ('blue-green') is of course always debatable. Different people may even perceive it differently. The Wiki photo for CuCl2.2H2O for instance appears more blue than green to me but not as deep blue as Jamit's photos.

jamit - 30-3-2012 at 18:16

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.

jamit - 30-3-2012 at 20:31

@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.

woelen - 31-3-2012 at 10:19

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]

jamit - 31-3-2012 at 20:16

@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]

jamit - 1-4-2012 at 04:35

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]

blogfast25 - 1-4-2012 at 09:04

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!

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 (CuCl42-

is green, Cu(H2O)62+ 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]

Nathaniel - 5-4-2012 at 06:40

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?

jamit - 5-4-2012 at 12:08

@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.

Nathaniel - 6-4-2012 at 03:26

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

CHRIS25 - 7-4-2012 at 05:18

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?

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]

blogfast25 - 7-4-2012 at 05:21

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.

CHRIS25 - 7-4-2012 at 05:36

Quote: Originally posted by blogfast25

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.

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

woelen - 7-4-2012 at 06:25

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 ...

Arthur Dent - 7-4-2012 at 06:36

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

jamit - 9-4-2012 at 23:06

@ 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.

woelen - 10-4-2012 at 01:13

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).

mr.crow - 22-7-2012 at 14:50

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?

jamit - 23-7-2012 at 03:43

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]

mr.crow - 23-7-2012 at 14:31

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

triplepoint - 23-7-2012 at 16:52

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. :mad:

Rogeryermaw - 23-7-2012 at 18:24

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.

Vargouille - 24-7-2012 at 04:25

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.

tetrahedron - 7-11-2012 at 02:37

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.

mr.crow - 7-11-2012 at 20:25

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!

tetrahedron - 8-11-2012 at 07:49

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.

tetrahedron - 9-11-2012 at 16:10

Quote: Originally posted by mr.crow

After leaving the CuCl2 in a desiccator for many weeks with NaOH and CaCl2 the small green crystals turned light blue. Success!

failure! i left my CuCl₂ solution on a large petri dish in a cool room to dry. as previously mentioned, the stuff smells acidic, so i guess the hydrated chloride hydrolyzes and gives off HCl even in the cold. the leftover must be basic, probably Cu2(OH)3Cl. indeed, today i noticed some of the crystals had turned light blue. i was reminded of your post and reckoned you must have been using a base as a desiccant (otherwise the HCl would soon have reached an equilibrium in the closed space of a desiccator, and the decomposition of the CuCl₂ would have stopped). my conclusion proved correct.

Vargouille - 9-11-2012 at 17:30

I am somewhat skeptical of that analysis of the results. According to this study on the hydrolysis of copper ions, *β1 = 7.7 x 10^-9 and *β2 = 2.1 x 10^-17, where *β represents the equilibrium constant of the hydrolysis. Compare this to the *β1 of iron (III), which this article claims is 10^(−2.18 ± 0.01). I have made copper chloride previously, and they turned out green due to formation of CuCl4^-2 ions. Some time later, after leaving them out, they turned blue, which I attributed to a return to CuCl2 as opposed to more complex tetrachlorocuprates. Moreover, with much less access to water, the hydrolysis of copper is expected be lower as a hydrated salt than for a solution.

tetrahedron - 10-11-2012 at 01:36

that might explain the color change, but not the HCl given off, which necessarily leads to a stoichiometric imbalance between copper and chloride ions in the residue.

Vargouille - 10-11-2012 at 10:20

After reading up on the thread again, I found woelen's post that contained this line:

Quote:

The grass-green compound is an acidic compound, which contains HCl as well. It contains complexes like HCuCl3.xH2O and these cause the green color.

This implies that the crystals themselves are not hydrolyzing, but rather, giving off HCl already contained within them, and in doing so, forming the blue CuCl2 sans tri- and tetrachlorocuprate complexes. The HCl would have been formed by the hydrolysis of copper in solution. Assuming [Cu⁺²] is at a maximum, according to the *β1 and *β2 given above, around 0.0021 M of HCl would be formed, and the hydrolyzed copper components precipitated. They would have been filtered or contained within the green copper chloride crystals. The small amount of HCl produced by the hydrolysis accounts for the acidic smell, since the odor threshold of HCl is only about 7 mg/m³.

tetrahedron - 11-11-2012 at 06:53

thanks for pointing that out, that explains many things. BTW can anhydrous copper sulfate precipitate from aqueous solutions?

my product has lost the smell and is now turquoise:

a closer look shows distinct crystal species admixed, namely the colored CuCl₂.2H₂O needles plus clear NaCl cubes (no noticeable sulfate crystals):

Vargouille - 11-11-2012 at 07:56

I would be quite nonplussed if anhydrous copper sulfate precipitated from an aqueous solution. Double salts? Sure. But not anhydrous forms of salts that have hydrates.

As for the sodium contamination, that's pretty par for the course. From what I read on this forum, it's astoundingly difficult to completely remove sodium contamination. It's one of the benefits of using copper carbonate to make soluble copper salts, since sodium contamination tends to be lower, provided that the CuCO3 is washed well.

12AX7 - 11-11-2012 at 09:04

I would suppose CuSO4 can be grown from a high temperature solution (i.e., in an autoclave), since the decomposition point of CuSO4.5H2O is relatively high. Compare to NaSO4.10H2O, which melts at 80C or so and therefore the anhydride can be grown from solution at atmospheric pressure.

Concerning mixtures: I seem to recall I have a baggie of about a pound of undefined copper mixture. IIRC, it seems to have small sulfate crystals, lots of chloride, probably sodium chloride as well, and maybe others. I recall also I had dried this material over paper towels, which partially decomposed due to the acidity, and I assume the cellulose fragments (including some glucose) reduced some copper, putting brown Cu2O crystals into the mix as well! I suppose some day I should dissolve it, dump in Na2CO3 and do the carbonate workup.

Tim

tetrahedron - 11-11-2012 at 15:52

turns out that CuCl₂ is very soluble in ethanol, NaCl only slightly, whereas the sulfates not at all. this is what the product looks like after recrystallization from warm 95% EtOH (in shades of grey):

the needles formed several star-shaped centers on the petri dish. the ethanolic solution was again intensely green, yet no HCl smell was observed. the insoluble sediment gives only a pale blue solution when dissolved in water, a sign that most copper was removed. no reduction to the monovalent cation was noticed. note that some NaCl passes through the filter and falls out upon cooling; this was removed by decantation.

the desire to remove all sodium is usually confined to applications in pyrotechnics. a bit of contamination is no big deal, as long as it's not so blatant as in the crude product.

Vargouille - 11-11-2012 at 16:01

When purifying copper compounds, I use bicarbonate, since there's not much hydroxide formed, and copper bicarbonate decomposes into copper carbonate. When you use carbonate, you get a mixture of hydroxide and carbonate. If you use it immediately and you don't mind the uncertainty, it doesn't matter, but since I rarely use mine immediately, I dislike the uncertainty of the inclusion of hydroxide and oxide, from the decomposition of hydroxide, I use bicarbonate.