dawsonsuen
Harmless
Posts: 29
Registered: 26-7-2015
Member Is Offline
Mood: Complexed
|
|
Colour and synthesis of Copper Nitrate?
I have reacted some copper metal with some nitric acid earlier today and I have ended up with a green solution. (Stayed green upon adding water)
But last week, I have done the exact same experiment and my result was a dark blue solution.
Is this because copper reacts differently with different concentrations of nitric acid to produce different coloured copper nitrates?
Or is the green colour due to the formation of Copper(I) Nitrate? (If this exists)
Thank you in advance.
|
|
|
woelen
Super Administrator
Posts: 7977
Registered: 20-8-2005
Location: Netherlands
Member Is Offline
Mood: interested
|
|
The green color most likely is due to formation of nitrogen oxides, which remain dissolved in the liquid and on dilution lead to formation of nitrite
ion, which gives green complexes as well.
If you use excess nitric acid and heat the liquid after all copper has dissolved, then you get a much purer product. After heating, dilute with water
and then heat to boiling again, to be sure that any nitrous acid is destroyed and converted to nitrate and nitrogen monoxide, which is driven out of
the liquid. The final result will be a purely blue liquid, which contains copper nitrate and excess nitric acid.
Under these conditions you certainly don't get copper(I). That would immediately be oxidized to copper(II).
|
|
|
aga
Forum Drunkard
Posts: 7030
Registered: 25-3-2014
Member Is Offline
|
|
It was a while back when i made it and i still have a jar of royal Blue crystals.
|
|
|
ave369
Eastern European Lady of Mad Science
Posts: 596
Registered: 8-7-2015
Location: No Location
Member Is Offline
Mood: No Mood
|
|
I've made some copper nitrate recently by reacting copper with 70% nitric acid. The solution was blue. But when I evaporated it dry in a bowl, some
(not all) of the crystals came out green.
That was a headscratcher...
Smells like ammonia....
|
|
|
Texium
Administrator
Posts: 4508
Registered: 11-1-2014
Location: Salt Lake City
Member Is Offline
Mood: PhD candidate!
|
|
You were able to dry your copper nitrate by evaporation? I always have to dry it in a desiccator inside of the jar I want to keep it in and then put
the lid on before removing it just to keep it from liquefying... and it still gets damp.
|
|
|
aga
Forum Drunkard
Posts: 7030
Registered: 25-3-2014
Member Is Offline
|
|
Likely that mine are still saturated with HNO3 now i think about it.
|
|
|
ave369
Eastern European Lady of Mad Science
Posts: 596
Registered: 8-7-2015
Location: No Location
Member Is Offline
Mood: No Mood
|
|
Quote: Originally posted by zts16  | | You were able to dry your copper nitrate by evaporation? I always have to dry it in a desiccator inside of the jar I want to keep it in and then put
the lid on before removing it just to keep it from liquefying... and it still gets damp. |
Well, it was still somewhat damp, but definitely not liquid. The consistence of clay, or something like that.
Smells like ammonia....
|
|
|
aga
Forum Drunkard
Posts: 7030
Registered: 25-3-2014
Member Is Offline
|
|
Mine is like that : like clay (wet) yet granular too.
|
|
|
ave369
Eastern European Lady of Mad Science
Posts: 596
Registered: 8-7-2015
Location: No Location
Member Is Offline
Mood: No Mood
|
|
I think the green color of the salt is caused by a small amount of unreacted nitric acid that permeates and wettens the "clay".
Smells like ammonia....
|
|
|
Volanschemia
Hazard to Others
Posts: 340
Registered: 16-1-2015
Location: Victoria, Australia
Member Is Offline
Mood: Pretty much all of them!
|
|
Excess Nitric Acid won't cause a green colour. It is most likely, as woelen said, due to dissolved NOx forming nitrites,
which in turn form complexes.
"The chemists are a strange class of mortals, impelled by an almost insane impulse to seek their pleasures amid smoke and
vapor, soot and flame, poisons and poverty; yet among all these evils I seem to live so sweetly that may I die if I were to change places with the
Persian king" - Johann Joachim Becher, 1635 to 1682.
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
I would appreciate if someone would try to resolve the issue of unwanted NO, NO2, nitrite or free HNO3 in the aqueous Cu(NO3)2 by the following
procedure:
1. Add a small amount of CuCO3 not to exceed the total expected moles of the impurites mentioned above or the capacity of the vessel to address the
corresponding moles of CO2 released. Quickly close the vessel (to retain the CO2).
2. Apply pulse radiation (short microwave bursts) or uv treatment (vessel or its cover must permit penetration of the uv rays) being mindful of the
volume of expected gas generated in the closed vessel.
3. Periodically shake the solution mix (not stir).
4. Stop after numerous iterations or when solution color suggests only Cu(NO3)2 is present.
Optionally, upon release of CO2 gas at conclusion of treatment, measure volume of gas released as an indicator of the level of impurities.
I will reveal my understanding of the underlying chemistry, if there is any interest, upon the procedure performing at all satisfactory.
[Edited on 1-9-2015 by AJKOER]
|
|
|
Marvin
National Hazard
Posts: 995
Registered: 13-10-2002
Member Is Offline
Mood: No Mood
|
|
I ended up with a mixture of copper and sodium nitrate after electrolysis which I tried to separate by fractional crystallisation. What was odd was
the solution was blue when cold and green when hot.
The water of crystallisation seems important and pH may matter. You may also find that the reduction produces ammonium salts, I'm unsure if this is
favourable with copper, aside from colour this might make the solution dangerous to crystallise.
Colour changes are not too unusual with copper. Anhydrous copper (II) sulphate is white. The form we know and love is blue and some of the
intermediate hydrates I think are green or yellow. I am surprised the solution stayed green when water was added.
|
|
|
Metacelsus
International Hazard
Posts: 2531
Registered: 26-12-2012
Location: Boston, MA
Member Is Offline
Mood: Double, double, toil and trouble
|
|
Blue when cold, green when hot: it sounds like the formation of a complex is being affected by temperature.
In fact, in my high school chemistry class, the tetrachlorocuprate complex was used as an example of how temperature affects reaction equilibrium.
|
|
|
DraconicAcid
International Hazard
Posts: 4278
Registered: 1-2-2013
Location: The tiniest college campus ever....
Member Is Offline
Mood: Semi-victorious.
|
|
| Quote: Originally posted by Marvin | | Colour changes are not too unusual with copper. Anhydrous copper (II) sulphate is white. The form we know and love is blue and some of the
intermediate hydrates I think are green or yellow. |
No- the monohydrate is white, and the trihydrate is a light blue. I made them for the carnival.
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.
|
|
|
![[*]](images/xpblue/default_icon.gif){kind=icon}
