Sciencemadness Discussion Board - Copper Compound Reactions

Copper Compound Reactions

MrHomeScientist - 20-5-2011 at 08:35

Recently I've been trying to make a video for my YouTube channel about different copper compounds. The overall scheme was the following. In between each step, the precipitate was filtered off and dried (except with CuCl2, that was left in solution).

1) CuSO4 + 2NaHCO3 == CuCO3 + Na2SO4 + CO2 + H2O
2) CuCO3 + 2HCl == CuCl2 + CO2 + H2O
3) CuCl2 + 2NaOH == Cu(OH)2 + 2NaCl
4) Cu(OH)2 + H2SO4 == CuSO4 + 2H2O

I thought it would be nice to end back at copper sulfate, which is where I started.

Anyways I went through this whole scheme but something went wrong on the final step. My copper hydroxide was filtered, washed, and dried so that I could weigh it before adding the sulfuric acid. However, I seem to have trouble keeping Cu(OH)2 stable when drying and it had at least partially decomposed to make the dry product a dark green instead of the turquoise blue it should be. I suspect it reacted with air into black CuO (which I've seen happen before) and green CuCO3.
I wasn't too worried about this because sulfuric acid should react with any of these compounds to still produce copper sulfate, so I went ahead and added the acid.

Here's what should have happened: Everything reacts and I am left with a blue solution of copper sulfate with no precipitate left over.
Here's what did happen: The solution is now green (similar to my CuCl2 soln) with a light blue precipitate remaining insoluble in the acid.

Any ideas to what might have been in my air-dried Cu(OH)2 that caused this? I thought that anything in there should have reacted away to CuSO4 in the final step, but clearly something else is happening. I can post a picture of it when I get home tonight if that is useful.

--Edit--
One thought I had was the NaCl from step 3 might have reacted with sulfuric acid to make HCl, reacting with the Cu(OH)2 to bring me back to CuCl2, which would explain the green solution. I washed my Cu(OH)2 a few times with distilled water though, so I can't imagine there was very much salt contamination left in it.

[Edited on 5-20-2011 by MrHomeScientist]

Neil - 20-5-2011 at 08:49

Have you tried washing the remaining solids with strong HCl to see if it is copper chloride?

The WiZard is In - 20-5-2011 at 08:57

Quote: Originally posted by MrHomeScientist

Recently I've been trying to make a video for my YouTube channel about different copper compounds. The overall scheme was the following. In between each step, the precipitate was filtered off and dried (except with CuCl2, that was left in solution).

1) CuSO4 + 2NaHCO3 == CuCO3 + Na2SO4 + CO2 + H2O
2) CuCO3 + 2HCl == CuCl2 + CO2 + H2O
3) CuCl2 + 2NaOH == Cu(OH)2 + 2NaCl
4) Cu(OH)2 + H2SO4 == CuSO4 + 2H2O

Anyways I went through this whole scheme but something went wrong on the final step. My copper hydroxide was filtered, washed, and dried so that I could weigh it before adding the sulfuric acid. However, I seem to have trouble keeping Cu(OH)2 stable when drying and it had at least partially decomposed to make the dry product a dark green instead of the turquoise blue it should be. I suspect it reacted with air into black CuO (which I've seen happen before) and green CuCO3.

Tooo much sodium hydroxide perhaps?

http://tinyurl.com/3v6d8sd

LanthanumK - 20-5-2011 at 14:41

Is your H2SO4 pure? How long did you notice the light blue precipitate remaining insoluble?

MrHomeScientist - 20-5-2011 at 16:41

I performed a test on my unknown compounds, and here are the results (with pretty pictures!)

The beaker on the right is my concoction of unknown copper compounds. You can see the emerald green solution, with blue precipitate. There also appears to be a thin layer of a whitish precipitate settled above the blue. It's been a few days, and everything remains insoluble. The solution is still very much acidic with excess H2SO4.
The test tube on the left is the result of taking about 1mL of the green liquid and diluting it with a few mL of distilled water. Notice the color change to blue - this is a strong indication that it is indeed copper chloride.

For further confirmation, I dissolved a very small amount of silver nitrate in 2mL distilled water, and added that to the test tube. Immediately a white precipitate of silver chloride was observed, confirming the presence of the chloride ion. The solution is now a lighter blue, suggesting copper nitrate (as it should be).

Conclusion: the green solution is definitely copper(II) chloride. The question remains though, how did this happen? Did the sulfuric acid react with stray chloride ions from previous steps to somehow form HCl? That seems unlikely to me. There shouldn't have been much of that around.

And the second question remains: What is(are) the precipitate(s)?

===========Edit===========

I performed some tests on the precipitate, with interesting results!

I scooped some of the blue solid out of the bottom, placed it in a filter paper, and rinsed it with water to remove any of the green solution. Surprisingly, all the solids dissolved!

My blue crystals going easily into solution like this lead me to believe that this must be the copper sulfate I was looking for. So I transferred some into two test tubes for testing.

To the left tube I added a small piece of iron wire. If this is indeed copper sulfate, copper metal will plate out on the iron as it displaces copper in solution.
To the right test tube, I will perform an anion test, similar to what I did with silver nitrate above. This time I am looking for the sulfate ion, so I made up a solution of barium chloride. If sulfate is present, a white precipitate of barium sulfate will form. Here's the two tests taking place:

You can see the wire is visibly reddish, and a white precipitate immediately forms when BaCl2 is added to the right tube. Here's what everything looked like after a few minutes reacting:

On the left, lots of copper plated out and the blue color of the solution cleared. On the right, white barium sulfate is observed.

Conclusion: The precipitate is indeed copper sulfate.

So, I ended up with a concentrated solution of copper(II) chloride with a precipitate of copper(II) sulfate. Apparently, the sulfate is not soluble in a solution of the chloride. I'm still not really sure why this happened though. Adding sodium hydroxide in step 3 should have neutralized any remaining hydrochloric acid and yielded NaCl, and washing the precipitate in step 4 should have carried away most or all of this salt. I would not expect that there would be enough chloride left over to form such a concentrated solution of copper chloride.

That was pretty long-winded, but I enjoyed doing it. Anyone have any theories as to why this all happened?

[Edited on 5-21-2011 by MrHomeScientist]

MrHomeScientist - 20-5-2011 at 16:47

Edit: Sorry for the double post, I wanted to split up my comment replies from my experimental results.

Neil:
The copper chloride was soluble, and what was causing the green solution, so I don't think it'd be in the precipitate. I can try adding HCl to a sample to see what it does.

The WiZard is In:
Interesting reference. Amazing what was around in the 40s! The article says that excess hydroxide shouldn't matter in the case of Cu(OH)2 though. When I repeat this sequence of reactions it's probably smart to add the NaOH slowly anyway, so thanks for leading me to that bit of advice.

LanthanumK:
Pure enough I'd think, it's battery acid grade fresh from the bottle. I've never had any problems with it before. The precipitate has been sitting in there like that for 3 or 4 days now, with no noticeable change.

Second edit: I've updated my post above to reflect new experimental results.

[Edited on 5-21-2011 by MrHomeScientist]

LanthanumK - 21-5-2011 at 08:06

Did you produce http://en.wikipedia.org/wiki/Dicopper_chloride_trihydroxide instead of copper carbonate or copper hydroxide like you thought? This would explain the results. The result after dissolving this in H2SO4 would be a mixture of CuCl2 and CuSO4. The common ion effect may precipitate the CuSO4 in a concentrated H2SO4 and CuCl2 solution because CuCl2 is more soluble than CuSO4.

Too little sodium hydroxide may be your problem, rather than too much like the Wizard suggested above. Too much may make cuprates, though.

[Edited on 21-5-2011 by LanthanumK]

[Edited on 21-5-2011 by LanthanumK]

MrHomeScientist - 23-5-2011 at 19:04

That's really interesting LanthanumK, I hadn't heard of that compound before. I could easily see that happening with poor mixing of the solutions - I bet some NaOH was trapped unreacted in the goopy Cu(OH)2 product. It's an interesting balance, can't have too much or too little. I think slow addition might be the answer. I'll also probably just skip the drying step when I redo this, to ensure it doesn't oxidize in the air. Thanks for the tip!

Random - 17-7-2011 at 18:27

Is there some interesting copper compound that can be made? I already made copper acetate crystals, have copper sulfate, made CuCl2 solution and half a kilo of CuCO3. Now I don't know where to continue, any tips?

Mixell - 17-7-2011 at 19:03

You can try organo- metallic salts of copper, for example: Formate, oxalate, benzonate,citrate and much more.
Also compounds of the following type: Cs(x)CuM(2+x). Where X equals 1 or 2, and the M is a halogen anion (usually chloride, but sometimes iodide or bromide). Also, you can try replacing the halogens with some other ligands, although it may be tricky.

kryss - 18-7-2011 at 10:55

Copper Iodide is an interesting one as well as it instantly forms cuprous Iodide, if you have any Iodide of course.

sternman318 - 18-7-2011 at 12:21

Tetraaminecopper(II) Sulfate

Eye catching to make! Adding small amounts of an ammonia solution to CuSO4 ( or any copper salt) turns dark blue-purple where the ammonia is poured, then turns to pale copper hydroxide when stirred. Adding more ammonia eventually redissolves the Cu(OH)2 into the ammonia complex. Add some sort of acid and you will get your CuSO4 back plus the ammonium salt of the acid, or adding NaOH will precipitate out the hydroxide again ( I believe)

[Edited on 18-7-2011 by sternman318]

Grenn Solution

ChemCrazy - 27-7-2011 at 07:56

Green solution maybe is not copper chloride, maybe it is copper base carbonate, which is famous mineral malachite. I dont knew about it too. Once I made experiment: electrolysis copper sulfate, i wanted to knew, that happened when SO4 is produced and the copper base carbonate produced, because of high humidity and lot of CO2 is pruduced this salt. Actually copper carbonate is not dissolved in water, but copper base carbonate - Cu2CO3(OH)2, malachite is dissolved and when heated it dissolved and produced carbon dioxide, water and copper oxide, so you can test your green solution: take not many green solution, boil until powder will settle. Then heat powder and if it's colour will change to black, this water is with mineral malachite. If not, it is copper chloride.

Copper Salts

ChemCrazy - 27-7-2011 at 08:09

You can make: copper hydoxide - Cu(OH)2 with Copper Sulfate electrolysis in right conditions. Copper Oxide - CuO with heating Copper Carbonate or copper heating under 100 Celsius. Copper Nitrate with reaction Copper to Dinitrogen tetroxide - N2O4 or copper to concentrated nitric acid.

Chemistry Alchemist - 11-9-2011 at 08:52

What would be the best way to make Copper Nitrate without Nitric Acid?

blogfast25 - 11-9-2011 at 09:54

Probably calcium nitrate (an agricultural product) + copper sulphate:

Ca(NO3)2(aq) + CuSO4(aq) === > Cu(NO3)2(aq) + CaSO4(s)

But CaSO4 is still slighty soluble in water, so the copper nitrate will need recrystallising.

Copper salts can be obtained quite easily from copper (conductor) scrap, conc. H2SO4 ("Draino" et al) and hydrogen peroxide, search for Nurdrage's CuSO4 Utoob vid, e.g.

kryss - 11-9-2011 at 10:28

You will find Copper nitrate very hard to crystallise, it tends to hydrolyse as it gets concentrated, heated.

MrHomeScientist - 11-9-2011 at 10:39

Quote: Originally posted by kryss

You will find Copper nitrate very hard to crystallise, it tends to hydrolyse as it gets concentrated, heated.

I've had that same problem. I had some copper nitrate left over from doing the "silver tree" demo and wanted to crystallize it, so I left it in a Petri dish to evaporate. It never did, and I'm left with a few drops of extremely concentrated solution that refuses to crystallize.

cyanureeves - 11-9-2011 at 10:57

i have quite a bit of copper compound left over from when i separate copper from sterling using nitric acid. i use the sodium chloride method to get silver chloride in precipitate and leave the copper in solution. so what copper compound do i end up with? copper chloride? or both copper nitrate and chloride? i started mixing copper compounds left over from both silver and gold extractions with aqua regia and i got a bottle full of a bright green salt.if aqua regia makes gold chloride from 14kt. then the copper solution left over is also copper chloride isnt it?or is it copper sulfate if i drop the gold with sodium metabisulfate? adding the sodium chloride to the nitrated sterling would also leave me with a copper chloride solution?

[Edited on 11-9-2011 by cyanureeves]

blogfast25 - 11-9-2011 at 12:57

Quote: Originally posted by kryss

You will find Copper nitrate very hard to crystallise, it tends to hydrolyse as it gets concentrated, heated.

Wiki's solubility data on Cu(NO3)2 shows a very large difference in solubility between 100 C and 0 C. It should be possible to obtain crystals by cooling a hot, saturated solution.

kryss - 12-9-2011 at 14:05

To be fair i never tried seeding it, its been sitting in our shed for the last 5 years, still liquid and strangely enough with no dead spiders in!

TheNaKLaB - 19-9-2011 at 03:54

@MrHomeScientist
I had the same problem when I was trying to make tetraaminecopper (II) Hydroxide.
I found that it worked well when I dissolved the Sodium Hydroxide in water then cooled it in the fridge for about an hour. I also cooled the Copper Sulfate for the same time until they were the same temp. I reacted them together and precipitated Copper Hydroxide, which is the noticeable turquoise colour.

blankets - 19-9-2011 at 12:42

Hi all, long time lurker first time posting, blah, blah. I gotta say i luv these Cu complex's. Any one new to this should def start tinkering with the Cu. Any how on to my question i made a saturatedsol of CuCl2 and evap'd down to get some damp evergreen colored crystals, after a lot of time n effort i checked wiki and it says "CuCl2 may be purified by crystallization from hot dilute hydrochloric acid, by cooling in a CaCl2 ice bath" so it appears i took the wrong route.
Does any one hear have exp forming CuCl2 crystalls? What do they mean when they say dilute HCl acid, how dilute?
I've attached an image (hopefully).
Wiki also describes the appearence as blue-green, im assuming mine are evergreen in color because their still moist? I started with CuSO4 then CuCo, to the CuCl2, btw. thx

bbartlog - 19-9-2011 at 12:50

Looks a lot like my CuCl2. I think the issue that is being addressed by crystallization from dilute HCl is possible contamination with oxychlorides. Like many other transition metal chlorides, CuCl2 will give off some HCl if you try to just heat to dryness from aqueous solution, i.e. you get a small amount of Cu(OH)Cl mixed in with your CuCl2. I really think that in the case of CuCl2 this contamination is pretty minor, though; unlike ZnCl2 and some other strong Lewis acids, it can't attack water that easily.

AndersHoveland - 19-9-2011 at 21:57

Copper(II) nitrate trihydrate melts at 114 °C.
The anhydrous form cannot be obtained by heating the trihydrate, as it decomposes at the same time that the water is liberated.
The anhydrous form is a surprisingly reactive oxidizer.

I have tried reacting Cu metal foil with a solution of Cu(NO3)2, but there was not any reaction, even after several days.
Apparently Cu2O reacts with dilute HNO3 to form Cu(NO3)2 and copper metal.
http://www.ucc.ie/academic/chem/dolchem/html/comp/nitric.htm...

[Edited on 20-9-2011 by AndersHoveland]

rstar - 23-9-2011 at 00:12

Great Forum. Well this is my first post.

I tried to make Na2CuCl4 by adding NaCl to CuCl2 solution, and the solution went from bluish to green. But I got problems, to get that stuff out of the solution.

Chemistry Alchemist - 23-9-2011 at 00:17

what is Na2CuCl4 used for? well are there any byproducts from the reaction? if not, then cant you just evaporate the solution?

rstar - 26-9-2011 at 22:10

Na2CuCl4 = Sodium Tetrachlorocuprate(II)
Honestly I don't know any of its uses and I don't think that the reaction has any byproducts.
2NaCl + CuCl2 = Na2CuCl4

497 - 26-9-2011 at 22:21

I don't think that reaction exists. Do you have a reference for it?

Chemistry Alchemist - 26-9-2011 at 22:26

If it does exist and does work, then you should be able to simply evaporate away the solution to form Sodium Tetrachlorocuprate(II)

I guess if its a nice looking salt then it doesnt matter if it hasnt got any uses, may make a colour effect in a flame... Yellow + Green/Blue

rstar - 26-9-2011 at 23:09

A Wikipedia page on CuCl2,

<a href="http://en.wikipedia.org/wiki/Copper(II)_chloride#Properties_and_reactions"> Here </a>

It tells that CuCl2 reacts with chloride sources to form CuCl3- and CuCl42- ions.

Chemistry Alchemist - 26-9-2011 at 23:29

So in other words if would make copper 3 or 4 chlorides insteaspd of reacting to make Sodium Tetrachlorocuprate(II) ?

497 - 26-9-2011 at 23:37

Yeah it might work. Every preparation I've seen uses concentrated HCl though. The color change you observed could be the result of a small fraction of tetrachlorocuprate forming.

rstar - 26-9-2011 at 23:38

Maybe yes

[Edited on 27-9-2011 by rstar]

rstar - 26-9-2011 at 23:56

Here's what i got after evaporation, sorry for bad picture quality :

Chemistry Alchemist - 27-9-2011 at 00:19

Yeah I was going to say, where would the Sodium go if the copper gained another chlorine, would the copper 3 and 4 chloride make a better flame co our then copper 2 Chloride? Or doesn't the extra chlorine make any Difference?

Chemistry Alchemist - 27-9-2011 at 00:57

Quote: Originally posted by rstar

Here's what i got after evaporation, sorry for bad picture quality :

Is this ur suposibly Sodium Tetrachlorocuprate(II)?
Because that kinda looks exectly like my copper (II) chloride

The picture on the Left is no flash and the one on right is with the flash... Copper (II) Chloride in a concerntrated solution is one colour, and once dilute it turns another colour (cant remember if its blue to green or green to blue)

Chemistry Alchemist - 27-9-2011 at 01:13

ok for some reason my pictures didnt attach :/ prob coz if shit internet

Arthur Dent - 27-9-2011 at 03:46

@ rstar: A nugget of the purest green: http://www.youtube.com/watch?v=TkZFuKHXa7w

Robert

rstar - 27-9-2011 at 04:09

Quote: Originally posted by Chemistry Alchemist

Is this ur suposibly Sodium Tetrachlorocuprate(II)?
Because that kinda looks exectly like my copper (II)

I also think like that's just a mixture of CuCl2 and NaCl. The complex might have decomposed as I got all solvent evaporated.

Maybe the complex is only stable in solution form.

rstar - 27-9-2011 at 04:16

@Arthur Dent
Thanx for the video link, that man has done the "Green" gold, it might be eco friendly

rstar - 27-9-2011 at 04:21

Quote: Originally posted by Chemistry Alchemist

So in other words if would make copper 3 or 4 chlorides insteaspd of reacting to make Sodium Tetrachlorocuprate(II) ?

They are not Copper 3 or 4 chlorides
they are:
CuCl3- = Trichlorocuprate(II) ion
CuCl42- = Tetrachlorocuprate(II) ion

LanthanumK - 27-9-2011 at 09:22

When copper(II) is crystallized out of an acidic solution, it forms green crystals, such as my CuCl2. This is probably due to the presence of moisture, which allows CuCl4(2-) ions to be in there. When dried further, it turns bluer as the rest of the HCl is driven off, leaving aqua ions behind.

Mixell - 27-9-2011 at 10:06

NaCl and CuCl2 have different solubility, and I don't know about the solubility of NaCl in concentrated HCl, what you got is most certainly a mixture of NaCl and CuCl2, may be a small amount of CuCl3 -1/CuCl4/ 2- in there...
Those complexes are usually made with cesium/rubidium, mostly cesium. CsCl in soluble in concentrated HCl, whereas the complex salt is not.
Here is a detailed preparation of those salts and a bit of information about them: http://woelen.homescience.net/science/chem/exps/CsCuCl3/inde...

Precipitation of Copper Phosphate

Chemistry Alchemist - 29-9-2011 at 01:34

Sorry to change the subject but i made Copper Phosphate by reacting Phosphoric Acid with Copper Oxide, i now hve a deep Blue solution of copper phosphate, ive wanted the solution to evaporate to form the precipitate but for some reason it isnt evaporating... how can i precipitate it out? can i add more water to it? Copper Phosphate is insoluble in water but soluble in phosphoric acid... could this work?

Another method is reacting the solution with Sodium Carbonate, this will react with the acid to neutralize it and and precipitate the Copper Phosphate but because the Copper Phosphate is in solution, would the Sodium Carbonate also react with the Copper Phosphate to Form Copper Carbonate and Sodium Phosphate?

bbartlog - 29-9-2011 at 04:47

If you have a deep blue solution I would assume you have excess phosphoric acid, since copper phosphate is insoluble. What quantities of reagents did you use? Ideally it seems like you should add more CuO until all the phosphate is thrown down, but dilution might also work. It should be easy enough to draw off a few ml of your solution and dilute it with 20x its volume in water to see what happens.
As for the sodium carbonate idea, yes, I suspect you end up with copper carbonate and sodium phosphate.

Chemistry Alchemist - 29-9-2011 at 05:00

i just added CuO untill it hardly fizzed, so do you suggest keep adding the CuO untill no more dissolves? yeah ill try dilute a small proportion of it and see what happens