| Pages:
1
2 |
jamit
Hazard to Others
Posts: 375
Registered: 18-6-2010
Location: Midwest USA
Member Is Offline
Mood: No Mood
|
|
Wow, that is beautiful. What did you do ' chemcam' to make copper nitrate like that?
|
|
|
chemcam
Hazard to Others
Posts: 423
Registered: 18-2-2013
Location: Atlantis
Member Is Offline
Mood: I will be gone until mid-september, on a work contract.
|
|
Well thank you, what I did was use steam to evaporate the excess water from the solution and when I saw crystals forming I took the heat away and let
it further crystallize over night. It was a solid-like mush the next morning so I stirred it to get the shapes seen in the photo then just let it air
dry for a couple more days.
|
|
|
zirconiumiodide
Harmless
Posts: 11
Registered: 6-6-2014
Location: UK
Member Is Offline
Mood: No Mood
|
|
Impurity formed upon Drying Copper Nitrate
So over the course of several months i have managed to dry some Copper Nitrate solution i produced by the reaction of Copper Carbonate with Conc.
Nitric Acid to form some nice crystals but upon crushing they were still quite moist. To try and rectify this i placed back in my dessicator box and
left to dry further for a couple of weeks. But DISASTER! Some of the Copper Nitrate, although a nice powder now has decomposed to a red Cupric Oxide
which adds a rather unwanted impurity to my product. There is only a bit of this impurity as far as i can see and i don't want to ruin the otherwise
fine product further so as far as i can see i have three options...
1) Ignore the annoying impurity and store anyway
2) Add a very small amount of HNO3 via pipette or store for a few days under Nitric Acid fumes
3) Filter of the impurity
2 will probably work but don't want to damage the product further and 3 is undesirable as will have to go back through the drying procedure which took
months so might just have to ignore the impurity 
The reason for this decomposition is probably because i dry my chemicals in a desicator box on the window sil and the direct sunlight probably
catalysed the breakdown.
Any suggestions would be greatly appreciated, cheers
ZnI4 
|
|
|
Texium
Administrator
Posts: 4724
Registered: 11-1-2014
Location: Salt Lake City
Member Is Offline
Mood: Seeking gainful employment
|
|
That's certainly odd... the red copper oxide would be cuprous rather than cupric, so if that's indeed what it is then you would have had to have some
other impurity in there with it capable of reducing the copper(II) to copper(I). Where did your reactants come from? That could potentially help
determine what the problem might be.
|
|
|
zirconiumiodide
Harmless
Posts: 11
Registered: 6-6-2014
Location: UK
Member Is Offline
Mood: No Mood
|
|
Copper Nitrate was produced from 69% HNO3 with Copper Carbonate.
Copper Carbonate was produced from Copper Sulphate by reaction with Sodium Carbonate. Very little Sodium Sulphate should be in the product as Copper
Carbonate precipitate was rinsed with water, so product should be reasonably pure.
I decided to store the Copper Nitrate anyway as the decomposition appears to have only occured on one side of the evaporating dish containing the
compound and is limited to the surface. Probably <1% of the Copper Nitrate has decomposed. Upon grinding with pestle and mortar the resultant fine
powder is only slightly tinted by the impurity.
Thanks for the response and correction zts16. As the impurity is red it is as you suggested more likely to be cuprous oxide as appose to cupric oxide
which is black.
I think the sunlight (or heat from it) catalysed the breakdown as the side decomposed was that on the window side of the dish. I can't think of any
impurity in the product that would catalyse the breakdown. This decomposition was completely unexpected - i know high temperatures decompose Cu(NO3)2,
but sunlight? 
As not as much damage as previously thought has been done to the compound i will put up with the impurity. But in the future i will have to make sure
that i carry out the final drying of the crystals out of direct sunlight, if i make some more of this compound 
ZnI4
|
|
|
woelen
Super Administrator
Posts: 8143
Registered: 20-8-2005
Location: Netherlands
Member Is Offline
Mood: interested
|
|
I hardly can imagine that copper(I) is formed in your situation. The red stuff really must be some other impurity. I can imagine it is some iron
compound. Cheap copper sulfate can contain some ferrous sulfate (I had a sample from a pottry supplier which gave a red color with thiocyanate,
besides the black precipitate of cupric thiocyanate).
 Btw, zirconiumiodide = ZrI4, not ZnI4
|
|
|
Texium
Administrator
Posts: 4724
Registered: 11-1-2014
Location: Salt Lake City
Member Is Offline
Mood: Seeking gainful employment
|
|
Quote: Originally posted by zirconiumiodide  |
Thanks for the response and correction zts16. As the impurity is red it is as you suggested more likely to be cuprous oxide as appose to cupric oxide
which is black.
|
Yeah, I'm sorry, I was not actually suggesting that I thought that you had copper(I) oxide, I was just
trying to show how it would be very unlikely. Now that I have returned to re-read my previous post, I see how that could have been misleading.
I agree with woelen about the iron impurities. You never really know what's going to be in the store-bought stuff. Most of the time when I use mine,
there are little bits of mystery compounds left over in the product that I can't explain, unless I use the stuff I've recrystallized.
Do the thiocyanate test on a solution of your copper sulfate if you have the means to. In fact, what might be better would be to just test your
copper(II) nitrate solution to see if the iron impurities made it all the way through.
|
|
|
zirconiumiodide
Harmless
Posts: 11
Registered: 6-6-2014
Location: UK
Member Is Offline
Mood: No Mood
|
|
Don't know why i've been putting Zn instead of Zr! Must be writing these posts
late at night. Either that or Zinc impurity!
I was very shocked by what seemed like decomposition. But why else was it only seen on the sunlit side and surface of the product?
Also there have been no signs of impurity up until this point. Two weeks ago they were lovely (all beit damp) blue crystals ..So can only believe that under very dry conditions and under direct sunlight although
slow, decomposition occured over the course of the two weeks it was drying.
Copper Sulphate used is meant to be high purity - bought of ebay. But saying that seems to have some insoluble (and probably soluble) impurities!
Haven't currently got thiocyanate as still working in a rudimentary home lab set up and building up my stock of chemicals so will have to wait a few
months before performing this test when i get some Potassium Thiocyanate.
ZrI4
[Edited on 16-7-2014 by zirconiumiodide]
[Edited on 16-7-2014 by zirconiumiodide]
|
|
|
Texium
Administrator
Posts: 4724
Registered: 11-1-2014
Location: Salt Lake City
Member Is Offline
Mood: Seeking gainful employment
|
|
Well, the impurity could be more soluble than the copper nitrate, thus being hidden when the crystals were still wet and only coming out of solution
when they were more dry. It's really very unlikely that your copper nitrate is actually decomposing. You'd get a definite odor of NO2
coming off of there if it was.
|
|
|
zirconiumiodide
Harmless
Posts: 11
Registered: 6-6-2014
Location: UK
Member Is Offline
Mood: No Mood
|
|
Impurities would have been seen throughout the compound as the Copper Nitrate was powder dry, and not just restricted to the surface as this was. And
on one side, answer that question?
Decomposition can't keep being dismissed in the light of the evidence!
If i had blow torched it i would expect the same thing, probably more extreme, but there would be decomposition on the surface and on the side i
blowtorched.
There was some tinting of the NaOH drying agent yellow suggesting gas may have eminated from the compound but any NO2 gas present in the drying box
would have been absorbed by the NaOH meaning there could have been very little present upon opening of the box. It may also have occured days before
the box was opened giving ample time for the gas to be absorbed.
The only way to be certain, even though from the evidence i'm quite sure it was decomposition, would be to repeat what occured by placing a small
amount of pure Cu(NO3)2 in direct sunlight under the same conditions.
But generally i decided to raise this issue not only for suggestions on how to cure my now tainted Copper Nitrate but also to warn people against
making the same mistake i did as it was totally unexpected. 
ZrI4
|
|
|
hyfalcon
International Hazard
Posts: 1003
Registered: 29-3-2012
Member Is Offline
Mood: No Mood
|
|
| Quote: Originally posted by zirconiumiodide | Impurities would have been seen throughout the compound as the Copper Nitrate was powder dry, and not just restricted to the surface as this was. And
on one side, answer that question?
Fractional crystalization
ZrI4 |
|
|
|
chloric1
International Hazard
Posts: 1222
Registered: 8-10-2003
Location: GroupVII of the periodic table
Member Is Offline
Mood: Stoichiometrically Balanced
|
|
Concentrating solutions with a desk fan
I bought a 8 inch desk fan from Walmart for $14 and a cheap one gallon flat plastic container. I made copper nitrate from 166 grams of copper sulfate
and 157 grams of calcium nitrate tetrahydrate. After washing the precipitated calcium sulfate I end up with 850 ml of solution. I use this fan set up
and even in 60 to 87% relative humidity in the middle of July I was able to reduce the volume significantly but now I got a much more concentrated
solution that I don’t think I can get dry without a dedicator so that’s where the Tupperware and sodium hydroxide will come in more to follow up
with later
Fellow molecular manipulator
|
|
|
chloric1
International Hazard
Posts: 1222
Registered: 8-10-2003
Location: GroupVII of the periodic table
Member Is Offline
Mood: Stoichiometrically Balanced
|
|
I managed to concentrate the volume to less than a third even in hot humid conditions. I’m considering not crystallizing it and keeping it as a
concentrated solution.
 
Fellow molecular manipulator
|
|
|
Radiums Lab
Hazard to Others
Posts: 236
Registered: 18-3-2025
Location: India
Member Is Offline
Mood: Experiencing the elegance of science.
|
|
If the pure crystals need to be dried we can use di ethyl ether and store it in a dessicator.
[Edited on 11-7-2025 by Radiums Lab]
Water is dangerous if you don't know how to handle it, elemental fluorine (F₂) on the other hand is pretty tame if you know what you are doing.
|
|
|
woelen
Super Administrator
Posts: 8143
Registered: 20-8-2005
Location: Netherlands
Member Is Offline
Mood: interested
|
|
Copper nitrate is really bad stuff on storage. It absorbs water from the air and even storage in a tightly closed container can give problems.
A few years ago, I purchased copper(II) nitrate x-hydrate,. 1 kg, for cheap (spare material from an old lab). It was a little clumpy and the container
was old and tattered. I transferred all of it to a big glass jar with a metal cap and a plastic liner inside. I stored it in the attic in our house,
where in summer it can become quite hot (35 C or so) and in winter it can become as cold as 5 C.
One year later, I found the jar with the copper(II) nitrate and all of it was one big crystalline lump, hard like rock, sealed into the glass jar. I
think it has liquefied somehow (maybe in hot summer, melting in its own water of crystallization) and later it crystallized again, but now in one big
lump. It is impossible to get it out of the jar, without hitting that with a hammer. I did not do that, it still is there as one big lump, sealed in
its jar.
On storage, it does not really degrade chemically (it remains copper(II) nitrate, no decomposition to other copper compounds), but its physical form
is not stable. All samples of copper nitrate I had in the past turn into hard to handle lumps. I think that it is best to make a highly concentrated
solution of this, and store that instead of the solid form. Especially if you intend to do experiments with it in aqueous solution, then there is no
need to keep it in solid form.
[Edited on 11-7-25 by woelen]
|
|
|
Precipitates
Hazard to Others
Posts: 183
Registered: 4-12-2023
Location: SE Asia
Member Is Offline
Mood: Acid hungry
|
|
Those were the good old days, when we could buy concentrated nitric acid in the UK.
Nostalgic reading these old posts from me (that I had mostly forgotten about).
| Quote: Originally posted by zirconiumiodide | | But DISASTER! Some of the Copper Nitrate, although a nice powder now has decomposed to a red Cupric Oxide which adds a rather unwanted impurity to my
product. |
That was weird, although I think I stirred the resulting mixture and ended up with a mostly blue powder.
11 years of storage and counting...
|
|
|
chloric1
International Hazard
Posts: 1222
Registered: 8-10-2003
Location: GroupVII of the periodic table
Member Is Offline
Mood: Stoichiometrically Balanced
|
|
| Quote: Originally posted by woelen | Copper nitrate is really bad stuff on storage. It absorbs water from the air and even storage in a tightly closed container can give problems.
A few years ago, I purchased copper(II) nitrate x-hydrate,. 1 kg, for cheap (spare material from an old lab). It was a little clumpy and the container
was old and tattered. I transferred all of it to a big glass jar with a metal cap and a plastic liner inside. I stored it in the attic in our house,
where in summer it can become quite hot (35 C or so) and in winter it can become as cold as 5 C.
One year later, I found the jar with the copper(II) nitrate and all of it was one big crystalline lump, hard like rock, sealed into the glass jar. I
think it has liquefied somehow (maybe in hot summer, melting in its own water of crystallization) and later it crystallized again, but now in one big
lump. It is impossible to get it out of the jar, without hitting that with a hammer. I did not do that, it still is there as one big lump, sealed in
its jar.
On storage, it does not really degrade chemically (it remains copper(II) nitrate, no decomposition to other copper compounds), but its physical form
is not stable. All samples of copper nitrate I had in the past turn into hard to handle lumps. I think that it is best to make a highly concentrated
solution of this, and store that instead of the solid form. Especially if you intend to do experiments with it in aqueous solution, then there is no
need to keep it in solid form.
[Edited on 11-7-25 by woelen] |
I think your spot on I just really wanted the concentrate solution to put patina on metals and have something besides copper sulfate for copper
expiriments. I went ahead and set up my DIY Tupperware dessicator with sodium hydroxide just to see if it can be done with rudimentary means. I will
probably just redissolve in solution with known copper content. The alkali does an admirable job in pulling water. When I free the brittle solid
caustic brick from the clutches of my dessicator, I weigh it then combine it with same weight of distilled water. That way I get 50% caustic soda of
commerce and any carbonate or other impurities are precipitated. Filter through cotton at 50 ° C to lower viscosity and I have refractive crystal
clear lye solution that I use to make alkaline earth hydroxides and process heavy metals.
Fellow molecular manipulator
|
|
|
CuriousOnlooker
Harmless
Posts: 35
Registered: 2-1-2025
Member Is Offline
|
|
~12 months later
Condemnant quo non intellegunt.
Never fire a warning shot. It is a waste of ammunition. ~ Hunter S. Thompson
|
|
|
chloric1
International Hazard
Posts: 1222
Registered: 8-10-2003
Location: GroupVII of the periodic table
Member Is Offline
Mood: Stoichiometrically Balanced
|
|
After 8 weeks over NaOH
Fellow molecular manipulator
|
|
|
chloric1
International Hazard
Posts: 1222
Registered: 8-10-2003
Location: GroupVII of the periodic table
Member Is Offline
Mood: Stoichiometrically Balanced
|
|
Sodium hydroxide is good desiccant but one needs to keep and eye on it as it is one of those inorganic solids that like to “creep” vertically vis
capillary action as moisture is absorbed!
Fellow molecular manipulator
|
|
|
| Pages:
1
2 |
![[*]](images/xpblue/default_icon.gif){kind=icon}
