Sciencemadness Discussion Board - Drying copper nitrate?

Drying copper nitrate?

Pommie - 29-3-2006 at 06:01

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.

woelen - 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.

Pommie - 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?

12AX7 - 29-3-2006 at 10:28

Make a solution and titrate it?

Tim

chromium - 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]

chloric1 - 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.

Cu(NO3)2

MadHatter - 12-4-2006 at 18:55

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.

12AX7 - 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

BromicAcid - 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.

Determination methods

MadHatter - 13-4-2006 at 12:42

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.

Steve_hi - 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?

Hockeydemon - 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]

KonkreteRocketry - 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

Hockeydemon - 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?

watson.fawkes - 27-3-2013 at 15:23

Quote: Originally posted by Hockeydemon

http://www.ilpi.com/glassblowing/tutorial_ampule.html

jamit - 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.

Air dried CuNO3?

Steve_hi - 31-3-2013 at 05:30

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]

vmelkon - 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.

Steve_hi - 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

CHRIS25 - 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]

elementcollector1 - 8-4-2013 at 10:16

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

Adas - 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.

elementcollector1 - 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.

Possible nitric acid fumes

Steve_hi - 22-4-2013 at 11:58

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.

Copper Nitrate

chemcam - 23-4-2013 at 13:17

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.

jamit - 24-4-2013 at 00:25

Wow, that is beautiful. What did you do ' chemcam' to make copper nitrate like that?

chemcam - 24-4-2013 at 17:36

Quote: Originally posted by jamit

Wow, that is beautiful. What did you do ' chemcam' to make copper nitrate like that?

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.

Impurity formed upon Drying Copper Nitrate

zirconiumiodide - 15-7-2014 at 05:20

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 - 15-7-2014 at 09:04

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 - 15-7-2014 at 14:51

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 - 15-7-2014 at 23:08

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 - 16-7-2014 at 05:21

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 - 16-7-2014 at 06:39

Quote: Originally posted by woelen

Btw, zirconiumiodide = ZrI4, not ZnI4

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 - 16-7-2014 at 10:02

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 - 16-7-2014 at 10:57

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 - 16-7-2014 at 11:18

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