Sciencemadness Discussion Board

Anhydrous copper (II) nitrate

Zyklon-A - 9-12-2014 at 09:51

As many of probably know, copper nitrate hydrates cannot be dehydrated by heat or other conventional means. [1]
Anhydrous copper (II) nitrate is usually prepared by reacting copper or copper oxide with liquid or gaseous N2O4, which although I could do (not liquid N2O4), I would like to find an easier process.
Copper (II) nitrate is a volatile covalent compound, not a salt, which according to Wikipedia, sublimes in vacuum at 150-200°C[2]. I recently got a vacuum pump, but I don't know how much vacuum to use.

As for the actual process to make Cu(NO3)2, one idea I had is to distill a mixture of potassium nitrate and copper (II) chloride anhydrous:
2 KNO3(s) + CuCl2(s) → 2 KCl(s) + Cu(NO3)2(g)↑

Copper chloride melts at 498°C so it must be heated to that temperature at least.
Obviously, all this is useless unless my premise of the (above) reaction works. I can't think of any reason for it not to.
When trying to calculate the Gibbs free energy for the reaction, I was unable to find any information for the entropy of Cu(NO3)2.
The ΔH° of the above reaction = -16.7, slightly exothermic.
Since Cu(NO3)2 is not ionic, and thus contains more entropy, I assigned a S° value of 150 J/mol to it, just as a guess, since I couldn't find it's real value.
Anyway, at 498°C the ΔG = 28.86, slightly endothermic. However, this doesn't include the volatility of Cu(NO3)2, which will almost certainly drive the reaction to the <s>left</s> RIGHT.
Can anyone think of a reason this won't work? Blogfast25, I'm talking to you:D.
Any comments and constructive criticism are always welcome and appreciated.

[Edited on 9-12-2014 by Zyklon-A]

Amos - 9-12-2014 at 10:13

What kind of reaction vessel are you proposing for this distillation? Copper salts tend to react with metals more reactive than copper, especially at that temperature.

DraconicAcid - 9-12-2014 at 10:17

Chloride's a better ligand than nitrate is, generally, so I have my doubts.

Zyklon-A - 9-12-2014 at 10:51

No Tears Only Dreams Now, Is there a reason to not use ground glass? That's what I generally use.
DraconicAcid, Can you expound on that?

woelen - 9-12-2014 at 11:17

This will not work. At the temperatures involved, the nitrates will decompose, giving oxygen and nitrogen dioxide. You will end up with KCl, contaminated with CuO and copper oxychloride.

From what I understood, anhydrous copper(II) nitrate was made from N2O4, dissolved in an inert solvent (I don't know which one) and copper metal and the resulting compound (Cu(NO3)2.2NO2) was carefully heated to 80 C to drive off the NO2 and excess solvent and at 80 C the temperature was not so high that the anhydrous copper nitrate decomposes.

If making anhydrous copper nitrate were as simple as suggested in the first post, then why was it made for the first time only 50 years ago? The more recently a certain chemical was synthesized and isolated for the first time, the more difficult it will be. I myself think that any compound which for the first time was made after WW II will be very hard to make in a home/amateur setting. I do not say impossible, though!

[Edited on 9-12-14 by woelen]

Zyklon-A - 9-12-2014 at 11:58

Thanks for the explanation woelen, Don't know how I missed that!

DraconicAcid - 9-12-2014 at 12:35

Quote: Originally posted by woelen  
This will not work. At the temperatures involved, the nitrates will decompose, giving oxygen and nitrogen dioxide. You will end up with KCl, contaminated with CuO and copper oxychloride.

From what I understood, anhydrous copper(II) nitrate was made from N2O4, dissolved in an inert solvent (I don't know which one)


Ethyl acetate, according to Cotton and Wilkinson's text.

Amos - 9-12-2014 at 12:42

Glass could be used, but how are you going to get your glassware that hot in the first place?

Zyklon-A - 9-12-2014 at 13:07

Bunsen burner comes to mind. But since it won't work, I'm in no hurry to waist the methane.

blogfast25 - 9-12-2014 at 13:10

Quote: Originally posted by Zyklon-A  
Blogfast25, I'm talking to you:D.
Any comments and constructive criticism are always welcome and appreciated.


[Edited on 9-12-2014 by Zyklon-A]


I'm with woelen on this, 100 %. I'm not sure where you get the idea that anh. Cu(NO3)2 is a 'volatile covalent compound'? If it was, then your reasoning would be sound and the value for Delta H not very important: distilling off one of the reaction products would pull the equilibrium to the right, acc Le Chatelier.

Acc. wiki the MP of the anhydrous form is already 256 C, the BP would obviously be higher would decomposition not occur. Alas, it does.

[Edited on 9-12-2014 by blogfast25]

Zyklon-A - 9-12-2014 at 13:18

Quote: Originally posted by blogfast25  


Acc. wiki the MP of the anhydrous form is already 256 C.

From the second link in the OP post: " Anhydrous copper nitrate forms deep blue-green crystals and sublimes in a vacuum at 150-200 °C." -Wikipedia.
Quote:

I'm not sure where you get the idea that anh. Cu(NO3)2 is a 'volatile covalent compound'?

From the first link in the OP post:
Quote: Originally posted by woelen  


Anhydrous copper (II) nitrate is not a salt, but a volatile covalent compound.

[EDIT] "Left" was a typo, I meant "right"

[Edited on 9-12-2014 by Zyklon-A]

Oscilllator - 9-12-2014 at 23:22

I'm with woelen on this one - If heating copper nitrate to dehydrate it causes it to decompose, then dry distillation is just going to result in decomposition.

One other possible method is azeotropic distillation with toluene or similar. I've done this with nickel nitrate, and it certainly worked to an extent, but I'm not sure how far the dehydration went. The nickel nitrate initially dissolved in its water of hydration, and then solidified in a lighter green colour (see picture).

blogfast25 - 10-12-2014 at 05:54

Quote: Originally posted by Zyklon-A  
From the second link in the OP post: " Anhydrous copper nitrate forms deep blue-green crystals and sublimes in a vacuum at 150-200 °C." -Wikipedia.


Well, if that is entirely correct then your experiment would be worth doing. But I have my doubts re. that Wiki statement. Also, how deep is the vacuum mentioned? And how deep can you go?

Zyklon-A - 10-12-2014 at 09:19

I couldn't find anything on how much vacuum was used. I can go to about 0.1 torr easily.

blogfast25 - 10-12-2014 at 10:01

Quote: Originally posted by Zyklon-A  
I couldn't find anything on how much vacuum was used. I can go to about 0.1 torr easily.


Do it. That's an order! ;)

Glass will be attacked by molten salts but it will probably survive the one test.

DraconicAcid - 10-12-2014 at 10:10

Quote: Originally posted by blogfast25  

Glass will be attacked by molten salts but it will probably survive the one test.


Depends on the salt. Molten salts containing lithium will mess up glassware, and basic melts such as sodium hydroxide will chew them up like candy, but molten chlorides and nitrates will be fine, as long as it's not near the softening point.

(I once did dissolve a copper salt in NaNO3/KNO3 eutectic melt, and I think I put it under vacuum, too, but didn't notice any evaporation. But that was a long time ago, and there wasn't much copper in the mixture.)

ETA: I did dry some lithium chloride under vacuum at about 500 oC for that project, thinking that it was about 100 oC under the melting point of the chloride, so it should be fine. I failed to realize that the glassware softened at 500 oC, so I came back the next day to find my lovely long Schlenk tube flattened like a deflated balloon and my vacuum pump overworked. Watch the temperature under vacuum.

[Edited on 10-12-2014 by DraconicAcid]

Zyklon-A - 10-12-2014 at 10:19

It might take a few days, most of my lab is packed away. Also I'll need to buy a few adaptors and vacuum safe glass. I've shattered an RBF with less vacuum than 0.1 torr, and I doubt 500°C will help.

blogfast25 - 10-12-2014 at 10:23

I electrolysed molten pure PbCl2 some 30 years ago, in a wide borosil test tube. The attack was very noticeable but the tube was afterwards still usable, after about 1 h slightly above the MP. Nice pure lead at the bottom of the melt.

Molten alkalis really go w/o saying.

Zyklon-A - 10-12-2014 at 10:28

Actually, only one of the reactants needs to be molten right?
Potassium nitrate fuses at 334°C so I won't have to go higher than that. At that temperature, Cu(NO3)2 will probably decompose much slower.

DraconicAcid - 10-12-2014 at 10:46

Quote: Originally posted by Zyklon-A  
Actually, only one of the reactants needs to be molten right?
Potassium nitrate fuses at 334°C so I won't have to go higher than that. At that temperature, Cu(NO3)2 will probably decompose much slower.


The mixture will melt at a lower temperature than either pure substance. If you were to heat your KNO3 to 333 oC and then add the CuCl2, it would be like sprinkling salt on ice (but you're better off mixing them before heating them). The mp of CuCl2 is about 600 oC, so I wouldn't expect it to be terribly soluble in the melt at 340 oC, but 10% or so by mass should be fine (don't ask for a cite for that number- that was pulled out of thin air. I'd be happy to be proved wrong).