kt5000
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Copper(II) nitrate crystals melted
I did the Nurdrage HNO3 process recently:
2KNO3 + Cu + 4HCl --> 2KCl + CuCl2 + 2H2O + 2NO2
2NO2 + 2H2O --> 2HNO3 + H2 (I think)
I don't yet have warm fuzzies about storing HNO3, so I reacted it with a bit of copper to make Cu(NO3)2. After the
reaction, I removed the excess copper and let the solution dry. I got some awesome looking Cu(NO3)2 crystals, then something
odd happened. I left the crystals set out in the dish for 3-4 days, covered with a coffee filter. When I went back to the dish, I found they had
"melted" into a slurry that was the same color as the crystals.
No color change and no added reactants makes me think it's still Cu(NO3)2. Is that hygroscopic? It's all I can think at the
moment. But the whole thing evaporated to dryness previously. I've been meaning to put that dish under some low heat and see what happens. Anyone
know what may have happened? Can a material evaporate to dryness then become hygroscopic?
[Edited on 6-10-2014 by kt5000]
[Edited on 6-10-2014 by kt5000]
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kt5000
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Another theory: It was cold outside when I did the reaction between HNO3 and Cu. Wikipedia lists the boiling point of
Cu(NO3)2 hexahydrate as 26.4 C and seems to indicate it decomposes to the trihydrate above that temp. Is it possible the
initial reaction produced non-hygroscopic Cu(NO3)2 hexahydrate crystals, then when the weather warmed, they decomposed into a
hygroscopic Cu(NO3)2 trihydrate form? (which would pull water from the air and create the blue slurry)
If so, can I simply add cold distilled water to the dish to get the hexahydrate and place in the fridge to slowly dry? (I love these problems)
[Edited on 6-10-2014 by kt5000]
[Edited on 6-10-2014 by kt5000]
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DraconicAcid
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26.4 oC isn't the boiling point- it's the temp at which it loses water of hydration, and then dissolves in the water it's just given off. You could
argue that it's really a melting point.
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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kt5000
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Quote: Originally posted by DraconicAcid  | | 26.4 oC isn't the boiling point- it's the temp at which it loses water of hydration, and then dissolves in the water it's just given off. You could
argue that it's really a melting point. |
Makes sense, I think.. So it loses its water of hydration and ends up as the trihydrate dissolved in that water just given off. I don't understand
why that water would stay there and not evaporate. If I cool the dish below 26 C, will the hexahydrate form and crystallize?
[Edited on 7-10-2014 by kt5000]
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Texium
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| Quote: Originally posted by kt5000 | | Quote: Originally posted by DraconicAcid | | 26.4 oC isn't the boiling point- it's the temp at which it loses water of hydration, and then dissolves in the water it's just given off. You could
argue that it's really a melting point. |
Makes sense, I think.. So it loses its water of hydration and ends up as the trihydrate dissolved in that water just given off? I don't understand
why that water would stay there and not evaporate. |
It should evaporate, but not immediately, as long as it
stays above the temperature at which the hexahydrate forms anyway. It seems that letting it evaporate under a heat lamp, or outside on a hot sunny day
if weather permits would convert it all to the trihydrate. The only copper(II) nitrate that I currently have is sitting in my shed that currently has
a giant yellow-jacket nest in it, so I'd do some tests on it too if I could, but right now it's inaccessible.
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kt5000
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I'll try to add a bit of distilled water, put the dish on my hot plate at around 35-40 C, and see it I can crystallize the trihydrate. I don't want
the hexahydrate badly enough to refrigerate it.
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Metacelsus
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The actual reaction is NO2 + H2O -> HNO3 + HNO2
The nitrous acid, depending on concentration, may be oxidized by atmospheric oxygen to nitric acid. In any case, no hydrogen is formed.
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kt5000
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Disappointed at the trihydrate crystal attempt.. I added a bit of distilled water, stuck it on the hotplate above 30 C, and let it dry. The result
looks way more like CuSO4 or CuCl2 with some white junk around the edges. It definitely doesn't have the deep blue of
Cu(NO3)2
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jamit
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Copper nitrate is extremely hygroscopic and deliquescence. What probably happened is that the solution dried up during a very dry whether and then
over a period of 4 days the whether changed and was very humid and the copper nitrate picked up water and dissolved in its own water of hydration.
This has happened to me with my copper nitrate. Once you dried copper nitrate you need to store it in a tightly sealed bottle to prevent moisture
from getting in and even then it picks up some moisture and get wet. Next to sodium hydroxide, it's one of the most water loving solid chemical I
have in my lab.
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