Sciencemadness Discussion Board

Stability of the Cu(I) compounds

Metallus - 17-5-2013 at 09:51

When asking my university professors, i always received different answers about the stability of Cu(I) compounds and, in particular, of the Cu+ ion.

According to one of my professors, the Cu+ ion isn't stable in solution and tends to give
2Cu+ --> Cu(s) + Cu2+

According to another one of my profs, the Cu+ is stable and has to be oxidized by an oxidant agent.

The reason for this question was that i wanted to try out the Copper(I) acetylide but i have no idea how to get the Cu+ if it isn't stable; is there a way to oxidize Cu to Cu+ without getting Cu2+? I remember we once mixed Cu(II) with iodide to get CuI, where Cu got +1, but it isn't soluble. Also, when i add Al to a solution of CuCl2, i immediately get copper.

Thanks for your attention

ScienceSquirrel - 17-5-2013 at 09:57

This is how you make copper acetylide;

http://en.wikipedia.org/wiki/Copper%28I%29_acetylide

Do it on a small scale following a known procedure and handle it with care.

blogfast25 - 17-5-2013 at 10:22

Copper (I) compounds do require an oxidant to push them to copper (II) but not much. Even the relatively insoluble CuCl needs to be stored under water because air oxygen will get to it quickly and oxidise it to Cu (II).

That said, CuCl can be stored indefinitely in the right and not very problematic conditions. I believe it is used as a catalyst in the production of some substituted silanes. Dry and in the complete absence of water it should also store well.



[Edited on 17-5-2013 by blogfast25]

chemcam - 17-5-2013 at 10:25

Copper(II) Acetylide can also be made, however, it is less powerful and more sensitive than Copper(I) Acetylide.

There was a time when I didn't feel like making or procuring cuprous chloride, CuCl, so I instead followed the same procedure for making (I) but I used cupric chloride, CuCl2, so I ended up with (II). Copper(II) acetylide It is a pretty weak energetic material, and if you do make the less stable (II) then make very small amount to test its sensitivity first! <0.50g is what I do for a first time synthesis.

I use (I) as an abbreviation for copper(I) acetylide
I use (II) as an abbreviation for copper(II) acetylide

// \\ // \\ // \\ // \\

BlogFast25, will dissolved oxygen in water not oxidize the CuCl?

[Edited on 5-17-2013 by chemcam]

Metallus - 17-5-2013 at 10:54

Thanks for the fast answers but i have few more questions:

1) Where should i find CuCl? CuSO4 is no problem but i have no idea where to get the Cu(I) nor what its uses should be.
2) Is it possible to make the acetylide by using another starting compound? If CuCl is strictly necessary, is there any not extreme way to get Cu(I) from metallic Cu or Cu(II) salts?

Thanks again for yor attention and ye, shouldn't the O2 present in water oxidize the CuCl to CuCl2?

chemcam - 17-5-2013 at 11:10

1.) You will either have to make it or buy from a chem supplier, it's not available OTC.
2.) Yes, any soluble Cu(I) compound

That is why I was asking blogfast25 why he recommended storing CuCl under water. I would think that would be bad idea but I don't know for sure.

DraconicAcid - 17-5-2013 at 11:29

Copper(I) isn't stable towards disproportionation in solution, but it can be stabilized by anions with which it forms an insoluble compound. So CuCl, CuBr, CuI and CuCN are stable, but CuNO3 and Cu2SO4 are not stable (although they may be made under anhydrous conditions). CuCl and CuBr will often be contaminated with copper(II) compounds due to exposure to the oxygen and moisture in air.

The copper(I) state can be stabilized in certain complex ions, notably with chloride ion, acetonitrile and ammonia. If you dissolve copper(II) chloride in concentrated HCl or acetonitrile, and stir it over copper shavings for an extended time, they will react to give copper(I) complexes (as long as it's not exposed to air while its reacting). Dilute it with a large amount of degassed water, and CuCl will precipitate. Collect it by filtration, and then you can dissolve it in ammonia for your acetylide reaction.

blogfast25 - 17-5-2013 at 11:46

Copper (I) iodide is easy: copper (II) sulphate solution+ iodide solution = insoluble copper (I) iodide precipitate+ iodine solution;

Cu2+(aq) + 2 I-(aq) === > CuI(s) + 1/2 I2(aq)


This reaction is used to create a standardised I2 solution, from a standard Cu2+ solution plus a small excess of KI, for purposes of titrometry (standardising sodium thiosulphate solutions for instance).


Tap water contains air (thus oxygen) but freshly boiled water doesn't, because gases are very poorly soluble in hot solvents. CuCl is indeed often sold stored under water, in a hermetically sealed container.



[Edited on 17-5-2013 by blogfast25]

papaya - 17-5-2013 at 12:17

Check out one of the PCB etching methods: it uses CuCL2 + HCL solution to dissolve the copper. If I remember correctly, equation seems like this :

H2[CuCl4] + Cu ---> 2 H[CuCl2]

Note that acid must be in a large excess and because it reaches equilibrium at some point you probably won't make pure complex. Also I think if resulting solution contacts with air, it'll be oxidized back to Cu2+.

[Edited on 17-5-2013 by papaya]

Adas - 17-5-2013 at 12:35

I've heard that CuCl can be made by reducing CuCl2 with ascorbic acid.

chemcam - 17-5-2013 at 12:47

Thats funny because right now I am recrystallizing ascorbic acid from vitamin c capsules. I will give this a shot if its an actual reaction. I don't understand how it would work. Care to explain?

12AX7 - 17-5-2013 at 22:28

IIRC, Brauer says Cu2SO4 can be isolated. Takes specific conditions though.

Cu2O and CuCl decompose to Cu(0) and Cu(II) in H2SO4, for example (in HNO3 I would imagine the Cu(0) is oxidized, with release of NOx).

Cu(I) is stable in concentrated Cl- solutions, as a clear (pure Cu(I)) or extremely dark brown (mixed Cu(I), Cu(II)) solution. Rapid dilution precipitates CuCl, so a complex is keeping it happy.

Cu(I), in the presence of moisture, is sensitive to oxygen in any form, gaseous or dissolved.

Tim

Adas - 18-5-2013 at 11:11

Quote: Originally posted by chemcam  
Thats funny because right now I am recrystallizing ascorbic acid from vitamin c capsules. I will give this a shot if its an actual reaction. I don't understand how it would work. Care to explain?


2 CuCl2 + C6H8O7 ----> 2 CuCl + 2 HCl + C6H6O7 (dehydroascorbic acid)

[Edited on 18-5-2013 by Adas]

Fantasma4500 - 20-5-2013 at 10:27

if youre looking for an easy route to Cu(I)Cl then i have perhaps something of interest for you..

take some concentrated HCl..
to this add CuCl2 (or make it by copper carbonate + HCl)
now add aluminium foil.. this will obviously heat up alot, and i believe the CuCl2 will speed up the reaction aswell as how quickly it starts up..
if you care you should try to light the hydrogen formed by this which burns lightblue.. very pretty.. (:

what happens:
CuCl2 + HCl + Al > Cu
CuCl2 + Cu > Cu(I)Cl
when you have gotten this thing highly boiling take a large amount of cold water and dump it into.. suddenly you get a nice bluewhite precipitate..

you can let this settle for a few minutes or just filter right away.. when this is filtered if you want this for acetylides you can just drop strong ammonia straight on the filter as there will be copper metal clumps stuck in the filter aswell, the tetraamminecopper chloride will go through the filter as a very strongly darkblue coloured liquid

it should be able to be stored like this.. 12AX7 mentioned this aswell.. i have about 100 mL of this stored and last time i checked after many months its still darkblue

if you want to dry it, as in the Cu(I)Cl as i remember you need to dry it in SO2 atmosphere (dont know if were talking pure SO2 completely anhydrous or just abit of SO2 in air or what)

anyways you can pretty much prepare the tetraammine in 10 minutes, compared to boiling copper metal with HCl and CuCl2 for several hours also wasting precious energy on heating it externally (:

blogfast25 - 20-5-2013 at 10:39

Antiswat:

CuCl can be prepared more simply by boiling a mixture of CuSO4 and NaCl with copper powder, filings or wire, overall reaction:

CuSO4(aq) + 2 NaCl(aq) + Cu(s) === > 2 CuCl(s) + Na2SO4(aq)

After separating off the supernatant solution replace it with ammonia solution and the colourless [Cu(NH3)2]+(aq) complex forms.

Adas - 20-5-2013 at 13:13

Quote: Originally posted by blogfast25  
Antiswat:

CuCl can be prepared more simply by boiling a mixture of CuSO4 and NaCl with copper powder, filings or wire, overall reaction:

CuSO4(aq) + 2 NaCl(aq) + Cu(s) === > 2 CuCl(s) + Na2SO4(aq)

After separating off the supernatant solution replace it with ammonia solution and the colourless [Cu(NH3)2]+(aq) complex forms.


Does it really work? To me, it just looks like the reaction of CuCl2 and Cu. Or does the sodium sulfate help the reaction somehow?

bbartlog - 20-5-2013 at 13:35

If you have CuCl2, then CuCl2 + Cu -> 2CuCl will work (but boiling would be advisable). CuSO4 is more widely available, though, and it is also likely that that reaction would proceed more quickly because the CuCl is rendered more soluble by the presence of NaCl.

Fantasma4500 - 20-5-2013 at 17:43

Quote: Originally posted by bbartlog  
If you have CuCl2, then CuCl2 + Cu -> 2CuCl will work (but boiling would be advisable). CuSO4 is more widely available, though, and it is also likely that that reaction would proceed more quickly because the CuCl is rendered more soluble by the presence of NaCl.


true, but you can skip the heating of stove / connecting gashoses or whatever by adding HCl to CuCl2 and then adding aluminium foil.. nearly instantly boils

Tdep - 20-5-2013 at 19:40

Did this yesterday

Beaker A: 30g CuSO4 and 9g NaCl
Beaker B: 4.5g NaOH and 7g NaHSO4

and it appear to work, got quite a bit of lightly blue precipitate (assume a bit of oxidation happened, tap water wouldn't have helped that) of CuCl.

I can't remember where I got this from, or how/why it works, but it seems to.

[Edited on 21-5-2013 by Tdep]

woelen - 20-5-2013 at 23:17

Copper(I) only is stable in combination with certain ligands. Compounds like Cu2SO4 and CuNO3 do not exist, they are unstable and would immediately disproportionate to CuSO4 and Cu or Cu(NO3)2 and Cu.

In CuCl, chloride ion is the ligand and it coordinates to two copper ions besides it (and conversely, the copper ion has two chloride ions coordinating to it). In solution, the complex CuCl2(-) can exist, two chloride ions being coordinated to the copper(I) ion. As soon as such complexes are broken (e.g. by precipitating the Cl(-) ligands with something else, highly insoluble) then the copper(I) disproportionates to copper(II) and metal.

You can also have colorless Cu(NH3)2(+) in solution, but this complex is VERY air sensitive. At the surface of such a solution you get a blue color at once, because it absorbs oxygen from the air very easily. The same is true for solutions of CuCl2(-). Such solutions look brown when oxygen is absorbed from air.

Some information can be found here: http://woelen.homescience.net/science/chem/solutions/cu.html

blogfast25 - 21-5-2013 at 05:02

Quote: Originally posted by Adas  

Does it really work? To me, it just looks like the reaction of CuCl2 and Cu. Or does the sodium sulfate help the reaction somehow?


Yes, it really works. I did this as a student some 30 years ago. The sulphate ions are just spectator ions, they don't actually do anything. But CuSO4 is easier to get than CuCl2. Back then, mine came from a chem set, these don't usually include CuCl2!

It works essentially because of:

Cu2+(aq) + Cu(s) < === > 2 Cu+(aq)
Cu+(aq) + Cl-(aq) === > CuCl(s)


[Edited on 21-5-2013 by blogfast25]

Fantasma4500 - 21-5-2013 at 05:19

i guess i misunderstood something..
i took a amount of CuSO4 and a amount of NaCl and dissolved
mixed
no ppt after pouring that after being boiled for approx. 10-20 mins in cold water

i dont see why you guys dont wanna go through carbonate into pure chloride??

CuSO4 + NaHCO3 >CuCO3*2Cu(OH)2(s) + Na2SO4
CuCO3*2Cu(OH)2 + HCl > CuCl2 + H2O + CO2

if you use excess you can pretty much perform this in one step, or well nearly

CuSO4 + CaCO3 > CaSO4 (s) + CuCO3*2Cu(OH)2 (s)

the HCl would by this way only react with the copper carbonate and CuCO3 can be bought in bulk quanities cheap

bbartlog - 21-5-2013 at 06:01

Quote: Originally posted by Tdep  
Did this yesterday
Beaker A: 30g CuSO4 and 9g NaCl
Beaker B: 4.5g NaOH and 7g NaHSO4

and it appear to work, got quite a bit of lightly blue precipitate (assume a bit of oxidation happened, tap water wouldn't have helped that) of CuCl.
I can't remember where I got this from, or how/why it works, but it seems to.
[Edited on 21-5-2013 by Tdep]


There doesn't seem to be any possibility of reduction here (of Cu(II) to Cu(I)), so I doubt that it worked, and I don't think you can take the presence of a light blue precipitate as evidence. My expectation is that what you have is more or less Cu(OH)2 via
CuSO4 + 2NaOH -> Na2SO4 + Cu(OH)2
...with the NaCl and NaHSO4 likely playing no role (though they might buffer the pH, I suppose).

You could run this reaction hot (say, 80C) as a test. If the precipitate is still pale, that's interesting and would suggest that what you have is at least not just some hydroxide of copper. If the precipitate turns dirty brown / black then you've got Cu(OH)2 or rather at that point CuO...

Tdep - 21-5-2013 at 06:08

I'm pretty confident it's not a straight copper hydroxide. When the solutions are first mixed the bright blue hydroxide is clearly formed but on stiring the colour fades and it really is quite a different compound formed. It's less of a blue, more of a pale green, distinctly different to the copper hydroxide.

I mean i'm open to suggestions, the chemistry didn't look right to me, I'll try to see if I can decompose the solid under solution but I don't think that's it.

blogfast25 - 21-5-2013 at 09:45

Quote: Originally posted by Antiswat  
i guess i misunderstood something..


Lets go back to basics:

CuSO4(s) === > Cu2+(aq) + SO4(2-)(aq) (dissolution of copper sulphate)

NaCl(s) === > Na+(aq) + Cl-(aq) (dissolution of sodium chloride)

Cu2+(aq) + Cu(s) === > 2 Cu+(aq)

Cu+(aq) + Cl-(aq) === > CuCl(s)

Overall: CuSO4 + 2NaCl + Cu === > 2 CuCl + Na2SO4

Not everyone has Cu carbonate lying around and preparing it for this purpose is a waste of time and chemicals.

[Edited on 21-5-2013 by blogfast25]

ElectroWin - 21-5-2013 at 10:25

Colour of copper (II) salts depends on pH; becoming green in more acidic conditions; and lovely deep blue in basic conditions.

[Edited on 2013-5-21 by ElectroWin]

blogfast25 - 21-5-2013 at 10:34

Quote: Originally posted by ElectroWin  
Colour of copper (II) salts depends on pH; becoming green in more acidic conditions; and lovely deep blue in basic conditions.

[Edited on 2013-5-21 by ElectroWin]


Are you sure you're not confusing with the colour of the tetrachloro cuprate anion [CuCl42-]? That is very green. Add HCl or NaCl to CuSO4 and it turns green, more intensely with chloride concentrations. But I've seen (and made) highly acidic copper nitrate solutions as blue as the ocean...

woelen - 21-5-2013 at 10:52

Color of copper(II) does not directly depend on pH, it depends on the nature of the ligands.

With water-ligands, the color is sky blue. You have this when you dissolve copper sulfate or copper nitrate or copper sulfamate.
When water ligands are replaced by chloride-ligands, then the color goes from blue to yellow. Mixed water/chloride complexes are green.
Copper(II) with bromide as ligands (e.g. CuBr4(2-)) gives deep purple color. Try dissolving some CuSO4.5H2O in a saturated solution of KBr or NaBr and you'll see that color.
With ammonia ligands copper(II) has a deep royal blue color, very beautiful to see this complex.

Complexes of copper with certain ligands only exist in certain ranges of pH, e.g. the ammonia complex only exists in weakly alkaline to moderately strongly alkaline solutions. At very high pH (> 14) it decomposes and leads to formation of copper hydroxide.

Complexes of copper with chloride or bromide only exist at low pH to neutral pH.

blogfast25 - 21-5-2013 at 13:07

Quote: Originally posted by woelen  
At very high pH (> 14) it decomposes and leads to formation of copper hydroxide.



... which at that pH dissolves to blue cuprate [Cu(OH)42-], another ligand causing blue colour of the complexed Cu2+ ion!



[Edited on 21-5-2013 by blogfast25]

Eddygp - 21-5-2013 at 13:39

Cu(NH3)4SO4 is an amazing colour. I strongly recommend its easy synthesis just to see it.