Sciencemadness Discussion Board

About to anhydrate Copper Nitrate, any suggestions ?

KonkreteRocketry - 14-2-2013 at 07:35

I am about to heat Cu(NO3)2(H2O)3 above the boiling point of water to take the water molecules away, so i will get the more powerful Anhydrous form, any tips or cautious I need to know before for it before I do this ?

Many thanks.

[Edited on 14-2-2013 by KonkreteRocketry]

hissingnoise - 14-2-2013 at 08:11

The trihydrate decomposes when heated:

2 Cu(NO3)2 → 2 CuO + 4 NO2 + O2

The anhydrous form is obtained by reaction of Cu with dry NO2!



KonkreteRocketry - 14-2-2013 at 08:15

Quote: Originally posted by hissingnoise  
The trihydrate decomposes when heated:

2 Cu(NO3)2 → 2 CuO + 4 NO2 + O2

The anhydrous form is obtained by reaction of Cu with dry NO2!






What if i heat it under the decoposition temperature ? like 101 degree celsius just above the boiling point of water

DraconicAcid - 14-2-2013 at 09:06

Quote: Originally posted by KonkreteRocketry  
Quote: Originally posted by hissingnoise  
The trihydrate decomposes when heated:

2 Cu(NO3)2 → 2 CuO + 4 NO2 + O2

The anhydrous form is obtained by reaction of Cu with dry NO2!






What if i heat it under the decoposition temperature ? like 101 degree celsius just above the boiling point of water


The water is actually part of the compound, so it doesn't really care how hot it is compared to the boiling point of liquid water. If you don't heat it enough, nothing will happen. If you heat enough that something will happen, it will decompose to copper(II) oxide.

(Actually, according to the Merck Index, the hexahydrate will decompose to the trihydrate at 26.4oC; the trihydrate will melt without decomposition at 114.5 oC. It doesn't give the decomposition temperature.)

(ETA: According to Cotton & Wilikinson's Advanced Inorganic Chemistry, "The hydrated nitrate cannot be fully dehydrated without decomposition. The anhydrous nitrate is prepared by dissolving the metal in a solution of N2O4 in ethyl acetate and crystallizing the salt Cu(NO3)2(dot)N2O4."

[Edited on 14-2-2013 by DraconicAcid]

KonkreteRocketry - 14-2-2013 at 09:32

Quote: Originally posted by DraconicAcid  
Quote: Originally posted by KonkreteRocketry  
Quote: Originally posted by hissingnoise  
The trihydrate decomposes when heated:

2 Cu(NO3)2 → 2 CuO + 4 NO2 + O2

The anhydrous form is obtained by reaction of Cu with dry NO2!






What if i heat it under the decoposition temperature ? like 101 degree celsius just above the boiling point of water


The water is actually part of the compound, so it doesn't really care how hot it is compared to the boiling point of liquid water. If you don't heat it enough, nothing will happen. If you heat enough that something will happen, it will decompose to copper(II) oxide.

(Actually, according to the Merck Index, the hexahydrate will decompose to the trihydrate at 26.4oC; the trihydrate will melt without decomposition at 114.5 oC. It doesn't give the decomposition temperature.)

(ETA: According to Cotton & Wilikinson's Advanced Inorganic Chemistry, "The hydrated nitrate cannot be fully dehydrated without decomposition. The anhydrous nitrate is prepared by dissolving the metal in a solution of N2O4 in ethyl acetate and crystallizing the salt Cu(NO3)2(dot)N2O4."

[Edited on 14-2-2013 by DraconicAcid]


I guess it can not decompose before the water molecules escape and become anhydrous and decompose, I heard you can force water molecules out with heat cus its not a strong bond.

DraconicAcid - 14-2-2013 at 09:37

Quote: Originally posted by KonkreteRocketry  

I guess it can not decompose before the water molecules escape and become anhydrous and decompose, I heard you can force water molecules out with heat cus its not a strong bond.


You heard incorrectly.

That may be true for other hydrates, but there are numerous ones that cannot be dehydrated with heat alone.

KonkreteRocketry - 14-2-2013 at 10:00

Quote: Originally posted by DraconicAcid  
Quote: Originally posted by KonkreteRocketry  

I guess it can not decompose before the water molecules escape and become anhydrous and decompose, I heard you can force water molecules out with heat cus its not a strong bond.


You heard incorrectly.

That may be true for other hydrates, but there are numerous ones that cannot be dehydrated with heat alone.


Umm can i dehydrate it with a vacuum pump ? where the boiling point of water automatically goes low and will water escape ?

Mercedesbenzene - 14-2-2013 at 10:43

I would have a little read through this
http://www.chem1.com/acad/webtext/chembond/cb09.html
website to get a better understanding on how "water of hydration" works. First of all it is better described as a coordination complex on a metal ion with water. Basically the water is bonded by it's lone pairs to the metal center with varying degrees of strength. If the bond is weak enough then the compound can be dehydrated by heating, ie magnesium sulfate. If the bond is strong then the compound can not be dehydrated by heating, eg aluminum chloride. If you try to heat hydrated aluminum chloride (Al(H2O)6Cl3) you will end up getting HCl vapours and some oxy/hydroxy salts of aluminum. So trying to remove water from a hydrated compound is not so much about adjusting the temperature or pressure to remove the water as it is a property of the compound. This simply means you can dehydrate some with heat, and some you cant. Off the top of my head I suspect copper nitrate would decompose, but I am not 100% sure, I need to have a look through some literature before I can answer that.

KonkreteRocketry - 14-2-2013 at 13:40

Ok so i did a test, i put copper nitrate on a alcohol flame, i diluted the alcohol to 70% so the flame is less hot.

Flame temp is around 180, sugar was decomped in the flame after around 10-15 sec.

Then, i put 3.3 gram of Cu(NO3)2(H2O)3 and heated in a steel container.

after around 10 seconds, steam started coming out..

after around 10 seconds, A small sign of decomp on the side of the container and i stopped the flame. I restarted it after a minute.

Again, steam coming out again, then i tried to stir it, after around 15 sec, another sign of decomp, and i stopped.

End : 3.1 grams.
Start : 3.3 grams.

I shall end up with 2.4 if all of it were to turn Anhydrous, so yeah, you can take some water out.

So In the end, resulted in 6% Anhydrous, or just took 6% water from the whole thing. More water were coming out, I will test with a cooler flame temperature from Acetone later, I will try to find whats the most amount of water you can get out of this trihydrate before decomp.

DraconicAcid - 14-2-2013 at 13:45

Quote: Originally posted by KonkreteRocketry  
Ok so i did a test, i put copper nitrate on a alcohol flame, i diluted the alcohol to 70% so the flame is less hot.

Flame temp is around 180, sugar was decomped in the flame after around 10-15 sec.

Then, i put 3.3 gram of Cu(NO3)2(H2O)3 and heated in a steel container.

after around 10 seconds, steam started coming out..

after around 10 seconds, A small sign of decomp on the side of the container and i stopped the flame. I restarted it after a minute.

Again, steam coming out again, then i tried to stir it, after around 15 sec, another sign of decomp, and i stopped.

End : 3.1 grams.
Start : 3.3 grams.

I shall end up with 2.4 if all of it were to turn Anhydrous, so yeah, you can take some water out.

So In the end, resulted in 6% Anhydrous, or just took 6% water from the whole thing. More water were coming out, I will test with a cooler flame temperature from Acetone later, I will try to find whats the most amount of water you can get out of this trihydrate before decomp.



If you want to try to do this without decomposing it, you don't want to use a flame. The heat will be too uneven. Use a boiling water bath for 100 oC, or a carefully heated oil bath for higher temperatures.

Did the copper(II) nitrate show any sign of reacting with the steel?

KonkreteRocketry - 14-2-2013 at 13:56

Quote: Originally posted by DraconicAcid  
Quote: Originally posted by KonkreteRocketry  
Ok so i did a test, i put copper nitrate on a alcohol flame, i diluted the alcohol to 70% so the flame is less hot.

Flame temp is around 180, sugar was decomped in the flame after around 10-15 sec.

Then, i put 3.3 gram of Cu(NO3)2(H2O)3 and heated in a steel container.

after around 10 seconds, steam started coming out..

after around 10 seconds, A small sign of decomp on the side of the container and i stopped the flame. I restarted it after a minute.

Again, steam coming out again, then i tried to stir it, after around 15 sec, another sign of decomp, and i stopped.

End : 3.1 grams.
Start : 3.3 grams.

I shall end up with 2.4 if all of it were to turn Anhydrous, so yeah, you can take some water out.

So In the end, resulted in 6% Anhydrous, or just took 6% water from the whole thing. More water were coming out, I will test with a cooler flame temperature from Acetone later, I will try to find whats the most amount of water you can get out of this trihydrate before decomp.



If you want to try to do this without decomposing it, you don't want to use a flame. The heat will be too uneven. Use a boiling water bath for 100 oC, or a carefully heated oil bath for higher temperatures.

Did the copper(II) nitrate show any sign of reacting with the steel?


Nah used it for other oxidizers like sodium nitrate, kno3, and this time copper nitrate, its fine, its stainless steel, they might coat something on it im not sure, but i know that Iron powder can burn with KNO3.

AndersHoveland - 14-2-2013 at 15:06

Same thing with the hydrates of Mg(NO3)2 and Al(NO3)2, if you heat them to try to drive out the water, they will just decompose.

some of these types of salts are actually more covalent in character than ionic. Magnesium ions and aluminum ions are fairly strong lewis acids. The hydrates of many of these salts are actually more analogous to addition compounds, where the water molecules of hydration are equivalent to a Lewis base.

KonkreteRocketry - 15-2-2013 at 04:18

I will change copper nitrate to sodium nitrate now, by mixing sodium carbonate and copper nitrate, but my copper nitrate is trihydrate, will it work ?

jamit - 15-2-2013 at 04:58

It is impossible to convert hydrated copper II nitrate into anhydrous copper II nitrate by heating it. It can only be done as noted by hissingnoise (also Woelen) by reacting copper in dry NO2. The best you can do is to put your hydrated copper ii nitrate in a desiccator and make it as dry as possible before putting it to a bottle for storage. Believe me, I tried doing what you did for months before discovering that it cannot be done!

Even making really pure copper II nitrate trihydrate is hard as it is extremely hydroscopic bordering on deliquescent.

DraconicAcid - 15-2-2013 at 09:12

Quote: Originally posted by KonkreteRocketry  
I will change copper nitrate to sodium nitrate now, by mixing sodium carbonate and copper nitrate, but my copper nitrate is trihydrate, will it work ?


Yes, that should work. You'll be doing that in solution, I'm sure, and once it's dissolved in water, it doesn't matter what hydrate it used to be.

KonkreteRocketry - 15-2-2013 at 09:28

Quote: Originally posted by DraconicAcid  
Quote: Originally posted by KonkreteRocketry  
I will change copper nitrate to sodium nitrate now, by mixing sodium carbonate and copper nitrate, but my copper nitrate is trihydrate, will it work ?


Yes, that should work. You'll be doing that in solution, I'm sure, and once it's dissolved in water, it doesn't matter what hydrate it used to be.


Yeah of course, and ok thank you, And, the result would be ANHYDROUS Sodium nitrate right ? where did the trihydrate go ? it became water ? :D

Umm seems legit :D I will try it now.

Boffis - 15-2-2013 at 10:07

When I was at university the preparation of anhydrous copper nitrate was a standard 2nd year practical and was described in "Practical inorganic chemistry" by G. Marr & B. Rockett. I t is in the section on the preparation and handling of liquid nitrogen dioxide.

AndersHoveland - 16-2-2013 at 22:31

Anhydrous Cu(NO3)2 is apparently very reactive:
Quote:

"at temperatures below –5°, the reaction of diethyl ether with copper(II) nitrate yields exclusively gaseous ethyl nitrite and complexed acetaldehyde. At higher temperatures, the acetaldehyde is further oxidised to acetate, with liberation of gaseous NO and NO2"

Anhydrous copper(II) nitrate as an oxidising agent.
L. C. Coard and R. E. Powell. Journal American Chemical Society, 296-297 (1967)


Apparently anhydrous Al(NO3)3 can be prepared by reacting dry NO2 with AlCl3, dissolved in cold liquid SO2 as the solvent.
(discussed in the "Lithium Perchlorate LiClO4" thread)

This might work on CuCl2 also, but the reaction might have to be modified. Either the initial CuCl2 reactant or the expected product (or both) must be soluble in the solvent, otherwise the reaction would not take place. AlCl3 is a strong lewis acid, and can form a transient adduct with the SO2, explaining its solubility.

Sometimes coppric nitrate is used for selective nitrations in organic chemistry. The hydrate is simply dissolved in acetic anhydride, and the nitration takes place under what is known as Menke conditions.

[Edited on 17-2-2013 by AndersHoveland]

virgilius1979 - 18-2-2013 at 11:22

What would be the best way to dry NO2 ? Pass it through MgSO4, or P4O10, or pure H2SO4 ? Would CaCl2 react with it ?

AndersHoveland - 18-2-2013 at 23:19

Quote: Originally posted by hissingnoise  

The anhydrous form is obtained by reaction of Cu with dry NO2!

I am not completely sure about this...

I know it works in the presence of water. But can the anhydrous form be produced from this reaction? Cu(NO3)2 is an oxidizer in its anhydrous form, and nitric oxide is a reducing agent, even reacting with oxygen at ambient conditions.

It might be better to try CuCl2 instead. At least that way there is chlorine to bind with the NO which is left over.

The reaction of nitrogen dioxide with aqueous base first produces a mixture of nitrites and nitrate, but additional nitrogen dioxide will oxidize the nitrite to more nitrate.

2 NaOH + 2 NO2 --> NaNO2 + NaNO3 + H2O

NaNO2 + NO2 --> NaNO3 + NO
Quote: Originally posted by virgilius1979  
What would be the best way to dry NO2 ? Pass it through MgSO4, or P4O10,

Seems good to me. ... Actually, I would not completely dismiss the possibility that P2O5 might be able to absorb NO2. I think the safest option would be anhydrous MgSO4.

You know, NO2 should not really be that difficult to dry. You could probably just use finely crushed Ca(NO3)2, it is hygroscopic.
Quote: Originally posted by virgilius1979  
or pure H2SO4 ?

I am not entirely sure. Concentrated sulfuric acid can actually absorb a mixture of nitric oxide and nitrogen dioxide gases to form nitrosylsulfuric acid.

Quote: Originally posted by virgilius1979  
Would CaCl2 react with it ?

Probably.

CaCl2 + 4 NO2 --> Ca(NO3)2 + 2 NOCl

[Edited on 19-2-2013 by AndersHoveland]

AJKOER - 22-2-2013 at 18:53

OK, if you have Thionyl chloride on hand, you may wish to test a small amount of your hydrated nitrate in the hope of making it anhydrous. To quote Wikipedia:

"Anhydrous metal chlorides may be obtained from hydrated metal chlorides by refluxing in freshly distilled thionyl chloride:[14]

MCln·xH2O + x SOCl2 → MCln + x SO2 + 2x HCl "

which is one expensive path with, in the current case, no guarantee.

[Edited on 23-2-2013 by AJKOER]

Motherload - 22-2-2013 at 19:37

I can't try this myself as I can't generate NO2 in my basement suite.
Landlords won't appreciate it too much..
I would suggest reacting CuO with N2O4 don't know how exothermic or violent the reaction might be ..... So a 100 mgs CuO to test.
And/Or
CuCl2 with AgNO3 with acetone as solvent.
CuCl2 and AgNO3 dissolve for sure and Cu(NO3)2 (Anhydrous) being more covalent than ionic should also dissolve in acetone as well. AgCl isn't soluble for sure.
Do not use ethanol as it presents a possibility of AgCNO to form.

KonkreteRocketry - 22-2-2013 at 23:10

Quote: Originally posted by Motherload  
I can't try this myself as I can't generate NO2 in my basement suite.
Landlords won't appreciate it too much..
I would suggest reacting CuO with N2O4 don't know how exothermic or violent the reaction might be ..... So a 100 mgs CuO to test.
And/Or
CuCl2 with AgNO3 with acetone as solvent.
CuCl2 and AgNO3 dissolve for sure and Cu(NO3)2 (Anhydrous) being more covalent than ionic should also dissolve in acetone as well. AgCl isn't soluble for sure.
Do not use ethanol as it presents a possibility of AgCNO to form.


I got an idea, get the NO2 gas from the decomp, then cool it down to -11 degree to get N2O4, then put the dry CuO to get Copper nitrate :)

Motherload - 23-2-2013 at 07:54

No matter how you generate NO2 .... It will need to be dried.
CaCl2 or MgSO4 column maybe.
You can easily condense it in a salt/ice bath cooled flask.
At around -5*C ... There should be very little fumes.

woelen - 23-2-2013 at 12:15

Making anhydrous Cu(NO3)2 is not easy at all. Before appr. 1960 people believed that this compound did not exist at all. But in the 1960's the anhydrous salt was prepared by reaction of metallic copper with a solution of thoroughly dried NO2 in ethyl acetate at low temperature. This does not sound like impossible for a home lab, but it will be hard. The experiment is dangerous. NO2 is very reactive and easily gives a violent reaction with ethyl acetate, hence the low temperature and the perfect drying of the reagents (traces of water may initiate a very violent reaction).

You can safely assume that if a compound only could be made somewhere in the 1960's and not before, that it is not an easy thing to do and that it certainly is not something for the inexperienced hobbyist.

AndersHoveland - 23-2-2013 at 14:17

Quote: Originally posted by woelen  
Making anhydrous Cu(NO3)2 is not easy at all. Before appr. 1960 people believed that this compound did not exist at all. But in the 1960's the anhydrous salt was prepared by reaction of metallic copper with a solution of thoroughly dried NO2 in ethyl acetate at low temperature.

Woelen, I am probably completely wrong here, but I question whether they actually prepared pure copper nitrate by this method. After all, such a reaction would have to result in nitric oxide as a by-product. And is not anhydrous copper(II) nitrate somewhat of a reactive oxidizer? How then can nitric oxide be formed? I wonder if perhaps what these researchers actually prepared was the nitrosyl complex of copper nitrate, NO·Cu(NO3)2. The formation of such an adduct would make the copper nitrate less oxidizing. Such nitrosyl complexes are quite common. I wonder whether the researchers bothered to quantitatively measure the nitrogen oxides resulting from decomposition when their product was heated. They may have just confirmed that no water was emitted, and jumped to an erroneous conclusion that they had prepared the pure copper nitrate.

If you have a link to the paper, we can check to see whether they quantitatively measured the nitrogen oxides given off per gram of their copper nitrate when heated.

[Edited on 23-2-2013 by AndersHoveland]

woelen - 24-2-2013 at 04:53

The information I have is not from a paper, but from a book. IIRC it is the book "General Chemistry" from Pauling. The material is described as deep blue and somewhat volatile.

I think that NO.Cu(NO3)2 would have copper in oxidation state +1, similar to the [Fe(NO)](+) complex, which has iron in oxidation state +1. But copper in oxidation state +1, combined with NO2 in the reaction vessel and with the NO3(-) ions in the final compound? Is that possible? So, I'm inclined to believe that they really made Cu(NO3)2 and not some nitroso-complex.

AndersHoveland - 24-2-2013 at 23:04

It would not really be in the +1 oxidation state. In fact, if you want to think of it this way, it is probably somewhere between +1 and +2. Do not forget that anhydrous copper(II) nitrate is an oxidizer.

The nitrosyl complex is really more of an adduct. And the common hydrate of copper(II) can be seen as an adduct also. Yet you would not say that the copper in copper(II) nitrate hydrate is in the +1 oxidation state.

Think of it this way: Cu(NO3)2 is like a lewis acid and NO behaves as a lewis base in this adduct. Because nitric oxide exists as a radical, the distinction between it acting as a base and acting as a reducing agent is a blurry one.

The definition of a lewis base is essentially the donation of an electron, with the formation of a covalent bond. The definition of a reducing agent in this case would be if the nitric oxide singularly formed a polar bond.

It is no coincidence that oxidizers are more likely to be acids, and reducing agents are more likely to be bases.

If the nitric oxide merely formed a polar bond, that would require the copper atom to be trivalent, since it is also bonded to two nitrate groups. Not that trivalent copper does not exist, but that is not the case here. The nitric oxide not only forms a bond with the copper atom, it also transfers charge into the nitrate groups. This explains how the copper can bond to 3 different groups - the bond order with each is less than 1. Alternatively this can be understood through resonance structures. Because the nitric oxide both bonds and transfers charge, it is acting as a base, not merely just a reducer.

There is nothing in NO·Cu(NO3)2 for NO2 to oxidize.

NO·CuCl2 is a more well known nitrosyl complex of copper. It is not stable in the presence of water, releasing its nitric oxide. Consider this: If the nitric oxide really had reduced the Cu+2, then a molecule of H2O would not be able to so easily displace it. Thus it is clear the NO is acting as a base in the complex.

[Edited on 26-2-2013 by AndersHoveland]