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KonkreteRocketry
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About to anhydrate Copper Nitrate, any suggestions ?
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]
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hissingnoise
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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!
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KonkreteRocketry
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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!
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What if i heat it under the decoposition temperature ? like 101 degree celsius just above the boiling point of water
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DraconicAcid
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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]
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KonkreteRocketry
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| Quote: Originally posted by DraconicAcid |
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.
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DraconicAcid
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| 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.
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KonkreteRocketry
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| 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 ?
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Mercedesbenzene
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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.
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KonkreteRocketry
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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.
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DraconicAcid
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| 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?
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KonkreteRocketry
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| 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.
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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.
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AndersHoveland
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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.
I'm not saying let's go kill all the stupid people...I'm just saying lets remove all the warning labels and let the problem sort itself out.
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KonkreteRocketry
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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 ?
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jamit
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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.
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DraconicAcid
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| 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 ?
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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.
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
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KonkreteRocketry
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| 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 ?
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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.
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Yeah of course, and ok thank you, And, the result would be ANHYDROUS Sodium nitrate right ? where did the trihydrate go ? it became water ? 
Umm seems legit I will try it now.
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Boffis
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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.
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AndersHoveland
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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)
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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]
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virgilius1979
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What would be the best way to dry NO2 ? Pass it through MgSO4, or P4O10, or pure H2SO4 ? Would CaCl2 react with it ?
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AndersHoveland
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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
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.
I am not entirely sure. Concentrated sulfuric acid can actually absorb a mixture of nitric oxide and nitrogen dioxide gases to form nitrosylsulfuric
acid.
Probably.
CaCl2 + 4 NO2 --> Ca(NO3)2 + 2 NOCl
[Edited on 19-2-2013 by AndersHoveland]
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AJKOER
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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]
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Motherload
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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.
"Chance favours the prepared mind"
"Fuck It !! We'll do it live !!"
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KonkreteRocketry
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| 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 
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Motherload
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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.
"Chance favours the prepared mind"
"Fuck It !! We'll do it live !!"
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woelen
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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.
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