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Pommie
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[*] posted on 29-3-2006 at 06:01
Drying copper nitrate?


A while ago I made some copper nitrate by mixing calcium nitrate and copper sulphate. The ice cream that I ended up with got filtered. I think I now have some fairly pure copper nitrate. My intention is to (eventually) make some lead dioxide electrodes. Being a bit of a pedant, I'd like to know how much copper nitrate I have. How on earth do I get this stuff dry? Can I bake it in the oven and then quickly squirrel it into an air tight jar? At what temp would I have to bake it to get it anhydrous? Is it white when anhydrous? This stuff turns into a liquid so fast. At what temp would I have to bake it to get the trihydate?

When (if) I get to the point where I have some lead nitrate, can it be used in a similar way to silver nitrate to show presence of chloride?

Mike.
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[*] posted on 29-3-2006 at 06:27


Making anhydrous Cu(NO3)2 cannot be done from an aqueous solution. The anhydrous compound has eluded chemists for more than a century or so. Only in the 1960's the anhydrous form was made by reacting copper with liquid NO2 (N2O4), crystallizing the adduct Cu(NO3)2.N2O4 and then carefully heating at around 85C to drive off the N2O4 solvent.

If you have Cu(NO3)2.xH2O, then heating results in loss of water vapour, oxygen, NO2. What remains behind is plain CuO.
The hydrated solid can be made from the solution with some difficulty. Heat the liquid, until it becomes dark blue and syruppy, then stop heating. Then let the stuff cool down. It gives a blue crystalline mess, which is fairly pure Cu(NO3)2.xH2O (3 <= x <= 6).
If you heat for too long, then you'll see brown fumes of NO2. If that happens, then you screwed up your stuff and basic copper salts will be formed.




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Pommie
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[*] posted on 29-3-2006 at 06:41


I heated it for a couple of days on my lab power supply. I'd guess it was around 40C and got a crystalline mass that looked like long needles. When I broke it up I had wet crystals. I left it to dry for a few more days and it just didn't get any dryer.

My question really boils down to, how can I use this in any quantative reaction when I can't weigh it.

That just made me think of specific gravity. Is that the only way?

Mike.

P.S There must be a temp where the trihydrate would be produced. Any idea?
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12AX7
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[*] posted on 29-3-2006 at 10:28


Make a solution and titrate it?

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chromium
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[*] posted on 29-3-2006 at 11:34


You can make solution of copper nitrate and this can be titrated with Na2CO3 or NaOH to get idea how much copper nitrate known volume of your solution contains. This solution can be used for reactions or evaporated before use. In any case amount of copper nitrate is not unknown any more.

Another way is to use tables of specific gravity for copper nitrate solutions (if you can find such tables). Boiling point of solution and optical activity can also be used to determine concentration if you have proper tables.

[Edited on 29-3-2006 by chromium]
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[*] posted on 12-4-2006 at 07:46


Iodometry-- What about using standardized potassium iodide solution? The cupric ion oxidizes iodide to elemental iodine which, in turn, can be titrated with standard thiosulfate. Starch should be mixed so the color transistion is more stark. I have not actually tried this myself but quantitive testing literature often states this method for estimating copper technique. Be mindfull of any free nitric acid.



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[*] posted on 12-4-2006 at 18:55
Cu(NO3)2


Another way would be to put a known quantity into solution and perform an electrolysis.
The weight of the copper on the cathode will indirectly indicate the amount of copper
nitrate available. You should get 33.88 grams of copper for every 100 grams of nitrate.




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[*] posted on 12-4-2006 at 20:18


- You could add some ammonia to that solution so you get a better visual indicator of how much copper is left. :)

Tim




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BromicAcid
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[*] posted on 13-4-2006 at 12:15


All of these methods sound like a pain in the butt. How about assuming your crystals are homogenous and quickly weighing out a known quantity then heating until they are throughly decomposed to CuO and weighing that product, they have used a similar procedure in industry for many applications where the pure compound cannot be isolated or is difficult to isolate. Then you'll know the precent copper by weight and you might assume that all of the copper was present in the form of the nitrate hence you know what weight of each gram of your hydrated salt you weigh is actually the copper nitrate.



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[*] posted on 13-4-2006 at 12:42
Determination methods


Bromic, that's the simplest method I've seen yet. That would work out to 42.41 grams of CuO
for 100 grams of the nitrate. Every time I made the nitrate I always got the wet crystalline
mass. I used the electrolysis because I already had all the gear on hand. I guess it's really
a matter of preference and whatever you have on hand. Your method requires the most
basic of equipment. One thing I'll try in the future is to remove the excess moisture under
warm conditions and vacuum. I'll be happy if I can at least get it into a dry hydrate.




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[*] posted on 27-3-2013 at 07:00


I Have to renew this thread because now I have an ice cream slurry myself. I reacted about 660 ml of nitric acid with about 200g of copper I removed the excess copper at what I hope was the end of the reaction. Now i have a small cake pan of copper nitrate but actually don't have any need of it. I want to just store it for the day I do need it.
Can I just put it in a jar. or will there be some unwanted degradation of the product and if so what might that be?
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[*] posted on 27-3-2013 at 07:20


As far as I know it will 'degrade' unless you put it in an ampule. Though I'm sure it has some sort of a shelf life in an container. I would think that it reacts with air exactly the same way as it does in solid copper form.

[Edited on 27-3-2013 by Hockeydemon]
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[*] posted on 27-3-2013 at 11:14


I tried this and i think i got a 2.5 hydrate instead of trihydrate, because the mass decreased, without decomposition.

also, i think a vacuum flash might do the work ;)
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[*] posted on 27-3-2013 at 11:45


Quote: Originally posted by KonkreteRocketry  
I tried this and i think i got a 2.5 hydrate instead of trihydrate, because the mass decreased, without decomposition.

also, i think a vacuum flash might do the work ;)


How hard is it to seal an ampoule under vacuum? Would sealing a glass tube at one end, attaching a vacuum to the other end, and sealing somewhere in the middle? What kind of torr do I need to be able to achieve?
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watson.fawkes
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[*] posted on 27-3-2013 at 15:23


Quote: Originally posted by Hockeydemon  
How hard is it to seal an ampoule under vacuum? Would sealing a glass tube at one end, attaching a vacuum to the other end, and sealing somewhere in the middle? What kind of torr do I need to be able to achieve?
http://www.ilpi.com/glassblowing/tutorial_ampule.html
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[*] posted on 27-3-2013 at 19:03


Quote: Originally posted by Pommie  
A while ago I made some copper nitrate by mixing calcium nitrate and copper sulphate. The ice cream that I ended up with got filtered. I think I now have some fairly pure copper nitrate. My intention is to (eventually) make some lead dioxide electrodes. Being a bit of a pedant, I'd like to know how much copper nitrate I have. How on earth do I get this stuff dry? Can I bake it in the oven and then quickly squirrel it into an air tight jar? At what temp would I have to bake it to get it anhydrous? Is it white when anhydrous? This stuff turns into a liquid so fast. At what temp would I have to bake it to get the trihydate?

When (if) I get to the point where I have some lead nitrate, can it be used in a similar way to silver nitrate to show presence of chloride?

Mike.




As noted by Woelen, anhydrous copper II nitrate cannot be made using aqueous solution.

Making the hydrated version of copper II nitrate is not easy either. Once you know that you have removed all of calcium sulfate and have a pure solution of copper nitrate, you need to carefully, at low heat, reduce the solution down to a concentrated solution of copper II nitrate. Then place the solution in an evaporating dish and put it in a desiccator and put the desiccator into the refrigerator. Allow it to dry for about one or two weeks.

I've been experimenting with copper II nitrate and its really difficult to crystallize the solid hydrate, especially during summer. In the winter where the garage is freezing and very dry, its alot easier. Once its dry, I bring the desiccator back into the house from the garage and allow it to dry further at room temperature until it is completely dry. I then quickly place it into a bottle for storage. In fact, that's what I'm doing right now. Good luck. Here's some pictures.


IMG_2835.JPG - 29kB IMG_2838.JPG - 42kB
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Steve_hi
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[*] posted on 31-3-2013 at 05:30
Air dried CuNO3?


From Beginning to end

It seems just being patient and letting it dry thinly spread on a glass pan does the trick

[img]C:\Users\Steve\Pictures\2013-03-31 2013\1up.jpg[/img]

[img]C:\Users\Steve\Pictures\2013-03-31 2013\2up.jpg[/img]


[img]C:\Users\Steve\Pictures\2013-03-31 2013\3up.jpg[/img]

1up.jpg - 206kB 2up.jpg - 208kB 3up.jpg - 212kB
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[*] posted on 5-4-2013 at 09:53


You have it in contact with a metal spoon? It will react with it and make copper + FeNO3. I would use glass and plastic only.
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Steve_hi
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[*] posted on 5-4-2013 at 15:51



You have it in contact with a metal spoon? It will react with it and make copper + FeNO3. I would use glass and plastic only.


Now I need to make FeNO3 are there other precautions advice or recommendations that differ from Copper nitrate? Instead of Iron Powder as I read in wicki. Could I use steel wool Or Iron Filings? I only want to make about an ounce Im starting a collection of compounds like Woelen has on his web site the first are Copper nitrate and copper carbonate I want to have as many colours as possible. Thanks




















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[*] posted on 8-4-2013 at 08:17


I have used both steel wool and iron filings for this. Both will give you quite some impurities from carbon but this can be filtered easily after the reaction has completed. Also a little soak in acetone helps to clean the steel wool, just make sure that after a soak you let it dry thoroughly in a warm spot, at least an hour. Forgot to add the highly dangerous nitrous oxide gas you will get, so outside is a must, and preferably down wind away from neighbours if you have any.

[Edited on 8-4-2013 by CHRIS25]




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[*] posted on 8-4-2013 at 10:16


Steel wool can contain a noticeable amount of chromium as well.



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[*] posted on 21-4-2013 at 09:33


If there was an anhydrous solvent that could dissolve both KNO3 and CuSO4, it would be easy, but I'm afraid that such a solvent doesn't exist. Yet. :P



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[*] posted on 21-4-2013 at 20:46


Methanol may be a possibility (pure speculation, no references). Dissolves CuSO4 to some extent (Wiki), no data listed for KNO3.



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[*] posted on 22-4-2013 at 11:58
Possible nitric acid fumes


I was wondering if its possible that after the copper nitrate is put into glass jars if fumes from the nitric acid could still be leaching out into the air. I notice after being in my lab for a while that my lips and nose become dry hard to describe but I think some chemical or another is giving off irritating vapours and I'd like to find the offending chemical. Ive bought and stored a number of chemicals in the last month or two, today I removed the copper nitrate and put them outside to see if it would make a difference. I also noticed that the bottle which I stored the remaining mother liquer from the copper nitrate was forming crystals along the outside bottom of the bottle. They are amber pet bottles which I bought from the Pharmacy.
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[*] posted on 23-4-2013 at 13:17
Copper Nitrate


The best results I have obtained by trying to make copper nitrate dry is just by letting it sit in dry, open air for a few days. Usually takes about 3-4 days for no more mass to be lost.





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