potassium nitrate and copper sulphate

Quote: Originally posted by Vargouille

The solubility for copper sulfate is the lowest,

No. Potassium sulfate is the least soluble and would precipitate first, especially at high temperatures (e.g, as you boil off water) as its solubility does not rise as rapidly as that of copper nitrate (or copper sulfate, for that matter) with increasing temperature.
You could proceed as follows:
Dissolve four moles (404g) of KNO3 in a liter of warm water
Dissolve two moles (499g) of CuSO4 (pentahydrate) in a liter of hot water
Mix these, boil off 1500g of water (you would likely need to filter or decant at times to get rid of the precipitating K2SO4, or deal with a lot of bumping...). Cool to freezing, filter to remove more K2SO4. The remaining liquid should still contain all the copper and nitrate, but less than 40g of K2SO4 (the remaining 310g of it having precipitated). You could then drive off the remaining water (carefully, so as not to decompose the copper nitrate) and have ~90% pure Cu(NO₃₂ with the remainder as potassium sulfate. Alternatively, adding 200g of absolute ethanol would precipitate 90% of the remaining K2SO4 and you could then obtain 98-99% pure copper nitrate.

Quote: Originally posted by Vargouille

Most interestingly, you might be able to make a double salt, but

Quote: Originally posted by Hexavalent

I have always wondered about the decomposition of copper nitrate - I once had a solution of copper nitrate that I evaporated....it absolutely reeked. The smell hit the back of your throat and was unbearable. At first, I thought it was excess nitric acid but now, looking back, it could be the decomp. of the nitrate.

Quote: Originally posted by Vargouille

...I should think that, considering the nature of ionic compounds, copper nitrate would decompose at boiling temperatures. Have you tried this procedure before?

Quote: Originally posted by bbartlog

No, it is entirely speculative. Given the existence of a double salt, it seems unlikely to work, and I can certainly believe that copper nitrate undergoes some slight decomposition at boiling temperatures, giving off NOx gases. But my CRC Handbook gives a decomposition temperature of 170C.

Just curious, if you heated the CuSO4 to 150 Degrees
Celcium, until it melt, and then put the KNO3 inside the
liquid CuSO4 would the reaction work? Of course its just
an imaginary situation, since I guess the Cu(NO3)2 has
a quite low decomposition temperature, so controlling
the reaction temperature would be really hard. Maybe if
that was possible it would be a way to get anhydrous
Cu(NO3)2, right?
I might be tripping here

Quote: Originally posted by tetrahedron

chris, you're interested in the anhydrous salt right? http://toxnet.nlm.nih.gov/cgi-bin/sis/search/a?dbs+hsdb:@ter... Quote: Methods of Manufacturing: ... Sublimation of copper(II) nitrate under vacuum from ... mixt of copper(II) bromide and silver nitrate at about 200 deg C. ...[Kirk-Othmer Encyclopedia of Chemical Technology. 3rd ed., Volumes 1-26. New York, NY: John Wiley and Sons, 1978-1984., p. V7 (1982) 105] **PEER REVIEWED**

Quote: Originally posted by AJKOER

...supposedly one can convert any metal hydrate

Quote: Originally posted by bbartlog

Quote: Originally posted by Vargouille   The solubility for copper sulfate is the lowest, No. Potassium sulfate is the least soluble and would precipitate first, especially at high temperatures (e.g, as you boil off water) as its solubility does not rise as rapidly as that of copper nitrate (or copper sulfate, for that matter) with increasing temperature. You could proceed as follows: Dissolve four moles (404g) of KNO3 in a liter of warm water Dissolve two moles (499g) of CuSO4 (pentahydrate) in a liter of hot water Mix these, boil off 1500g of water (you would likely need to filter or decant at times to get rid of the precipitating K2SO4, or deal with a lot of bumping...). Cool to freezing, filter to remove more K2SO4. The remaining liquid should still contain all the copper and nitrate, but less than 40g of K2SO4 (the remaining 310g of it having precipitated). You could then drive off the remaining water (carefully, so as not to decompose the copper nitrate) and have ~90% pure Cu(NO<sub>3</sub><sub>2</sub> with the remainder as potassium sulfate. Alternatively, adding 200g of absolute ethanol would precipitate 90% of the remaining K2SO4 and you could then obtain 98-99% pure copper nitrate.

Tried this at 1/100 scale, 4g KNO3 (stump burner) 5g CuSO4.5H2O (root killer);
20mL of water.
Result of first attempt ... reddish brown sludge on bottom of flask as soon as water got to boiling, lots of fizzing and splashing with choking smell that attacked back of throat.
I capped flask immediately, and reduced heat. Under vaccum condition, boiling is now proceeding in a reflux situation stably. Red brown sludge is not dissolving or turning white (color of potassium sulfate.)

This is consistent with later posts claiming that decomposition of copper nitrate produces that smell; and wikipedia stating that decompostiion occurs at 80C and above.

When I allow boiling to stop, color of solution is blue (not green as in photo).

I Will retry experiment with 99% Isopropyl alcohol to see if reaction will carry out slowly with only a little water present and desired product being dissolved in alcohol at low heat.

Edit: In first photo, I have added an extra 20ml of water when trying to reduce loss of nitrate.
Second photo is after boiling over-night. The brown sludge has all turned to pretty copper metal. The flask has been boiling under a partial vaccuum as the heat was turned down after stoppering the flask; so little to no air / oxygen from outside got in. It's difficult to see copper color through the solution; the color in second picture is now the true color of the solution.



[Edited on 26-1-2018 by semiconductive]

Quote: Originally posted by Hexavalent

I have always wondered about the decomposition of copper nitrate - I once had a solution of copper nitrate that I evaporated....it absolutely reeked. The smell hit the back of your throat and was unbearable. At first, I thought it was excess nitric acid but now, looking back, it could be the decomp. of the nitrate.

Quote: Originally posted by semiconductive

This is consistent with later posts claiming that decomposition of copper nitrate produces that smell; and wikipedia stating that decompostiion occurs at 80C and above.

Well there’s a reason why we shouldn’t trust Wikipedia, I’ve got a completely different set of values. When I redissolve my salts, including copper nitrate, I always heat to boiling and have had no issue with decomposition below 100C. Maybe there are impurities affecting the stability?



[Edited on 26-1-2018 by LearnedAmateur]

Quote: Originally posted by Vargouille

Barium has the least soluble sulfate of the three, but is quite toxic.

Quote: Originally posted by AJKOER

A word of caution with respect to avoiding iron impurities and air. [Edited on 26-1-2018 by AJKOER]

The second attempt:
4g Stump burner (KNO3)
5g root killer (CuSO4.5H2O)
+5g of H2O, +2g of Isopropyl alcohol ( 500mL water left vs. 200g alcohol in original post. )

I am attempting to crudely mimmic original formulation's final ratio of water to alcohol to see if reaction would proceed and both dissolve and precipitate slowly. eg: If I warm and cool the solution, I expect solubilities would change and allow reaction to proceed step-wise. The object was to avoid boiling or raising temperatures significantly above 40C.

I did NOT pre-mix chemicals with water or alcohol, but instead purposely added clumps of crystals so I could ascertain which chemicals had been used up, and hopefully this would slow down the reaction.

Upon warming the mixture, a surprise happened. The solution immediately began to fizz slightly and separated into two layers. Presumably, the top layer is isopropyl alcohol and the bottom layer (bright yellow) is an unknown hydration or perhaps double salt ?

I immediately added vinyl electrical tape to attempt to slow down the loss of gas. It held the stopper.


Once I realized heating would cause problems, then I tried shaking in a vibrator instead. That also caused it to fizz and stretch the vinyl tape (bad odor); so I stopped and let it set on a counter at about 8:30PM. By 9:30PM the solution had separated into three layers: A bottom layer of yellow, a middle layer of green (probably Cu[NO3]2.??H2O), and a clear layer of liquid on top. After an hour, the green layer had grown in thickness ... so it appeared the reaction was proceeeding toward Cu-Nitrate. Noticably, the blue CuSO4 crystals were on the bottom of the flask, while the while KNO3 crystals floated to the top of even the greenish layer of fluid. The KNO3 crystals were discolored extra orange on their surfaces. Small bubbles were also slowly being released from the yellow solution below.

So, the NO3 gas appears to form in the posassium nitrate crystals and make them buoyant; and also in the yellow sludge.

In any event, the final outcome this morning was surprising again. The green layer had completely disappeared and only pure yellow was on bottom with a clear layer of ?alcohol? on top.

I moved the solution back and forth, and could still see bluish-green copper sulfate crystals. Some very small potassium nitrate crystals are still attempting to float on the yellow solution.

After shaking the soltuion, some bluish-green fluid formed -- presumably dissolving of more copper sulfate crystals; but it's hard to tell because even the copper sulfate crystals look more greenish now -- as if they lost water of crystalization.

I'm wondering if the water I added simply isn't enough because the Cu(NO3)2 requires more water of hydration than I allowed for and it is dehydrating as the reaction proceeds.



[Edited on 26-1-2018 by semiconductive]
Note: After removing the tape, the flask decrompressed a little. It's been in the dark overnight, but under light and with loss of presure the yellow sludge is rapidly getting a brown hue on top, under the greenish copper solution I noted before.




[Edited on 26-1-2018 by semiconductive]

Quote: Originally posted by PirateDocBrown

Quote: Originally posted by Vargouille   Barium has the least soluble sulfate of the three, but is quite toxic. Uh, no. People eat BaSO4 every day, as an X-ray contrast for the GI tract.

Quote: Originally posted by clearly_not_atara

Quote: Originally posted by PirateDocBrown   Quote: Originally posted by Vargouille   Barium has the least soluble sulfate of the three, but is quite toxic. Uh, no. People eat BaSO4 every day, as an X-ray contrast for the GI tract. More importantly, barium nitrate is not sold as a fertilizer for less than ten bucks a kilo. Using potassium nitrate for this is amazingly wasteful, in money, time, and nitrate.

Quote: Originally posted by clearly_not_atara

You oxidized the isopropanol, reducing some nitrate present to nitrite. Copper nitrite forms a yellow-orange complex: http://www.sciencemadness.org/talk/viewthread.php?tid=29805 The "sludge" could be explained by some copper-nitrite complex migrating into a phase which also contains polymers of acetone. [Edited on 26-1-2018 by clearly_not_atara]

Please excuse the first sentence of the last post, I meant "nitrite" not nitrate; they both cause green colored solutions.

EDIT:::: Part of this post was originally an error, and I don't know how to delete it. The stump burner I bought is likely sodium pyrosulfate, which is a first for me. It wasn't marked on the label, anywhere. So, take this post as an explanation of how I got suspicious because of the solubility error in what I thought was potassium nitrate.The math is correct, the chemical likely wrong. I have since purchased genuine stump burner with potassium nitrate in it "spectracide Stump Remover" brand, from Fred Meyers, and will try again. The other stump burner, "stump out" from Ace Hardware is likely NOT potassium nitrate; though I wasn't able to get an MSDS from the manufacturer several people have complained about it, online and ace's computer listing says it's pyrosulfate.

---------------------------------------------------------------------------------------------------------
I have done part of the next expermient. I did not expect to get yellow material in solution with ammonia before dessication, but I did get it. That yellow stuff is a colloidal suspension that turns into a precipitate of micro-crystalline material. Yellow is not a normal copper ligand-ion interaction color, I'm pretty sure, because it's opaque. I am not sure if it's some kind of double salt, or just the effect of having two different crystals intermixed at a microscopic level.

However there was another unexepected surprise in the ligand coloration before it became yellow colloidal suspension and then crystals.

During the reaction with ammonia present, a clear red liquid formed. It's dark red, but it looked very transparent and not cloudy for a short period of time. Therefore, I assume it's a new color caused by some combination of sulfate and ammonia acting as ligands in solution.

I haven't desiccated the solution yet, as I discovered something else on the way. Let me first describe the experiment, the surprises, and show the photos of the steps.

This time, I used ammonia water from Ace Hardware which is suppsed to be a 10% solution; nothing but ammonia is listed on the bottle or the MSDS that I saw when I bought the bottle a year ago. It's lost some of it's smelling strength, which means that a certain amount of ammonia worked it's way out of the cap over the year. It's still pretty strong, though.

In any event, I decided to try and dissolve 4g of -----(stump burner) from ace into ammonia water and then chill it in a freezer before adding 5g of CuSO4.5H2O crystals. I hoped chilling the dissolved stump burner in ammonia would stop the release of gas that I noticed when I mixed the CuSO4 in the previous experiment and boiled. Again, the point of this attempt is to mix reagents with insufficient water to completely dissolve either of them, and see if the reaction can be carried out without having to boil off large amounts of solute afterward.

These are the constants I have looked up from the original recepie for the substances used:
4.04g of KNO3 / 101.1032 g/mol = 0.040 mol (Sol is 316mg/mL @ 20C --> 12.7mL water/4 gram KNO3, 2.46g/mL @ 100C --> 1.6mL water/4 gram KNO3)
4.99g of CuSO4.5H2O / 249.685 g/mol = 0.020 mol (Sol is 320mg/mL @ 20C --> 15.6mL water/5 gram CuSO4.5H2O )

So, the approach I wanted to take was to dissolve KNO3 in boiling ammonia water to bring it to super-saturation, then chill it below room temperature (10C or so). Ammonia is not a common ion with nitrate, so I decided to measure out enough water to dissolve all the KNO3 just assuming the excess concentration of ammonia would not affect the nitrate solubility. (But it didn't go as planned, which made me suspicious. )

I measured 2.46mL of water times 110% since ammonia is less dense than water; I hoped that would give a rough starting point for the solubility of KNO3 in 10% ammonia water. ( 2.7g of ammonia water. ) and the maximum water I wanted would be the amount that could dissolve all the copper nitrate formed, but not enough to dissolve the potassium sulfate that would be formed.

K2SO4 * 0.02mol * 174.259g/mol = 3.485g ( sol. W 0.11g/mL 20C --> 31.682g H2O )
Cu(NO3)2 * 0.02mol * 241.60g/mol= 4.832g ( sol. W 1.25g/mL 20C --> 3.866g H2O )

I brought 4g of stump burner it to a boil in a capped earlenmeyer flask with 2.7g of ammonia water. If this had been pure water for re-crystallization, it would have easily dissolved the whole amount at boil. But it did not. So, I kept adding ammonia water and watching the dissolution until I had enough water to dissolve all the copper nitrate in the final solution (which was not my original goal.) I ought to have stopped at 3.866g of H2O, but went all the way to 10.63g before quitting. Even then, the KNO3 had not all dissolved. Since this was almost 4x the ammonia water compared to the necessary distilled water, I figured the remaining stump burner had to be impurities; and just removed them. The wet impurities only ammounted to 100mg of the 4g put in the jar; so it's not critical.

In any event, even though the Erlenmeyer flask was capped with a glass stopper; none the less, I'm sure I lost most of the ammonia. The ability of strong ammonia in suppressing the solubility of KNO3 was not expected (Edit: because it wasn't KNO3)

I then chilled the solution in a freezer, down to about 10C. Absolutely NO KNO3 precipitate fell out of solution after 15 minutes of chilling.
I measured out 5g of CuSO4.H2O, and added a few crystals to the flask. There was no bubbling, the surface of the blue crystal immediately turned green, and a greenish diffusion streams could be seen dissolving in the solute. After about a minute, I could see white precipitate start falling out of solution. I don't know if it's KNO3 or K2SO4. I was satisfied that the solution wouldn't out-gas, so I added all 5 grams of CuSO4. No gassing happened. after a few seconds, I smelled the solution and there was no odor at all; neither ammonia, nor foul NO2 gas smell. So, this is much better than the prior experiment; but it makes me think I lost almost all the ammonia.

I capped the flask and clamped it tight, and let the crystals slowly dissolve while I gently swirled the solution once every five minute or so. It was now warming up toward room temperature. After an hour I smelled the solution to verify that nothing had changed. There were no strong smells all the way to the completion of the experiment.

The water continued to precipitate white material, and the water itself had a colloidal suspension until approximately half the copper sulfate had dissolved. Then suddenly, an organgish color could be seen coming off the crystals. Soon this color turned a transparent and very dark red. I did not monitor the solution for about a half hour, and when I came back; the whole flask was filled with a yellow suspension and the red coloring was gone. I still couldn't smell ammonia, or nitrate; and since the color was yellow, I decided to add ammonia water (1 gram) to see if the blue color of copper-ammonia complex would show up before mixing. It did, very dark blue (see photo). Once the solution mixed, though, the blue turned green (the color of Cu(NO3)2. ) and the yellow suspension fell to the bottom as a micro-crystalline precipitate. Since there was no alcohol, I assume there is nothing to oxidize to make nitrite and this ought to still be a nitrate compound. Ammonia complex is typically blue, and once dissolved I got green liquid. I have purchased sodium nitrite in order to repeat the experiment with a pure sodium nitrite for preserving meat; then I can check nitrite color in an analogous experiment.

I went to bed and when I got up in the morning I noticed that the top surface of the yellow precipitate had turned a brilliant white. The water above the precipitate was now a darker green. When I turned the light on, after a few seconds, a small amount of black precipitate formed on top of the white and then stopped. I could see it darken with time over a period of a minute. So, the solution appears to have something which is photo-sensitive (at least to UV light) and not RED LED night lights in it. After shaking the solution, a few times, no more dark precipitate formed.



Edit: The precipitate has since turned grey, slowly; while the solution remains green. The number of correspondences between the expected behavior of the potassium nitrate with copper sulfate, compared to the probable sodium pyrosulfate are astonishing.

[Edited on 30-1-2018 by semiconductive]

Quote: Originally posted by LearnedAmateur

Quote: Originally posted by semiconductive   This is consistent with later posts claiming that decomposition of copper nitrate produces that smell; and wikipedia stating that decompostiion occurs at 80C and above. Well there’s a reason why we shouldn’t trust Wikipedia, I’ve got a completely different set of values. When I redissolve my salts, including copper nitrate, I always heat to boiling and have had no issue with decomposition below 100C. Maybe there are impurities affecting the stability? [Edited on 26-1-2018 by LearnedAmateur]

Quote:

I've completed the second round of experiments. So far, the first two replies in the thread appear to correctly predict the difficulty in getting CuNO3, even though the second reply is based on incorrect solubility data which is pointed out by a third poster. Quote: Originally posted by Hexavalent   Yes, this could be problematic as both copper (II) nitrate and potassium sulfate are soluble in water, quite well. ... The overall equation is; 2KNO3 + CuSO4 &rarr; K2SO4 + Cu(NO3)2, but you will, of course, not just get copper nitrate as the product. What actually happens is; 2[K+] + 2[NO3-] + [Cu2+] + [SO42-) &rarr; 2[K+] + [SO42-) + [Cu2+] + 2[NO3-] Quote: Originally posted by Vargouille   I doubt it would work. The solubility for copper sulfate is the lowest, and that of copper nitrate is the highest. Most interestingly, you might be able to make a double salt, but it won't be simple to purify into either nitrate. Versus this post: Quote: Originally posted by bbartlog   Quote: Originally posted by Vargouille   The solubility for copper sulfate is the lowest, No. Potassium sulfate is the least soluble and would precipitate first, especially at high temperatures (e.g, as you boil off water) as its solubility does not rise as rapidly as that of copper nitrate (or copper sulfate, for that matter) with increasing temperature. You could proceed as follows: Dissolve four moles (404g) of KNO3 in a liter of warm water Dissolve two moles (499g) of CuSO4 (pentahydrate) in a liter of hot water Mix these, boil off 1500g of water (you would likely need to filter or decant at times to get rid of the precipitating K2SO4, or deal with a lot of bumping...). Cool to freezing, filter to remove more K2SO4. The remaining liquid should still contain all the copper and nitrate, but less than 40g of K2SO4 (the remaining 310g of it having precipitated). You could then drive off the remaining water (carefully, so as not to decompose the copper nitrate) and have ~90% pure Cu(NO<sub>3</sub><sub>2</sub> with the remainder as potassium sulfate. Alternatively, adding 200g of absolute ethanol would precipitate 90% of the remaining K2SO4 and you could then obtain 98-99% pure copper nitrate. The results I explain in this post are definitely done with the correct chemicals. eg: 4g & 5g of K2NO3 and CuSO4. However, the results are still complicated because of the alcohol. I need to re-test, to make sure the (disputed) breakdown CuNO3 data from Wikipedia + an unknown issue with CuNO3/K2SO4 and alcohol aren't ruining results. I also want try adding ammonium sulfate to see if Le chatlier's principle can be used to change solubility of K2SO4. My tentative conclusion, is that there is one of the following: decomposition of the nitrate, creation of a double salt with solubility LOWER than K2SO4 (!!!??), or else interaction (adduct?) of ions with alcohol to change the solubility of a double salt. I've tried the experiment under isopropyl, ethanol, and methanol; with water limited to the theoretical amount needed to dissolve the CuNO3 product. Eg: The reaction is carried out with that amount of water, plus the same or more in alcohol stirred vigorously for 24 hours under heat of 75C. I then transfer the solute to a second container and either distill off the alchohol (75C) or chil the solute to precipitate some of the dissolved salt; I then repeat the washing of the original precipitate by returning only the distilled or chilled alcohol solution to the original container. (Without more water in the case of distilled alcohol.) The distill/chill/rewash cycle was repeated ten to twenty times each. In chill cycles, a bright white precipitate formed in the second flask ... but was indistinguishable from the precipitate in the first flask; as both turn faintly green when exposed to sufficient water or dried in air. In distill cycles, the remaining water solution that is kept in the second flask turns increasingly blue. In all reactions, the alcohol solution is either aqua-marine or clear. Under strong alcohol the precipitate is always white. When reacting the checmicals under IPA, a very curious effect happened. As temperature is raised from 70 to 80C all the colored solution precipitates. The precipitation is so rapid (1-2 seconds) that I believe it's very sensitive to an exact temperature. I don't trust my thermometer to measure the actual temp within 1C, so am not reporting the threshold. But adding either ethanol or methanol will cause the hot precipitate to re-dissolve and the sensitivity to temperature is canceled. Re-running the experiment under ethanol-water or methanol-water does not have the negative solubility issue. The precipitate formed under alcohol solutions is always white. However, almost all the copper is in the white precipitate. Letting water get to the precipitate turns it faintly green, and ammonia water turns dark purple, no matter how many washings are done. Repeatedly washing the first precipitate in re-distilled methanol is the ONLY way that I was able to produce a concentrated and beautiful blue solution in the second flask. But the solubility of copper nitrate is much smaller than expected and re-distilling is impractical for producing CuNO3; even after 10 cycles only a tiny amount of CuNO3 is isolated. Returning our attention to the initial precipitate in the first flask. As a final check ... I re-dissolved ALL of the initial precipitate in 500mL of water. It turned light blue. I then evaporated off around 100mL of water at 90C using air to speed up the evaporation. Surprisingly the vast majority of precipitate to fell out of hot 90C solution a second time with 400mL!! of water still present. The color of the water solution is blue, but the precipitate is light green. This extremely low solubility and color of precipitate is a big surprise. I wonder if there is some kind of alcohol ?adduct? that can survive re-dissolution and precipitation under water. Wikipedia's dubious/ambiguous alcohol vs. water solubility data is a severe complicating factor. I am not certain which alcohol they are referring to, and if alcohol can chemically react with the nitrate or sulfate ions sufficiently to disrupt the experiment. [Edited on 11-2-2018 by semiconductive]

Summary of quantitative results of Cu2NO3 preperation via alcohol-water route and multi-distillation.
Using 4g K2NO3, 5g CuSO4, and 3.87g of water with a small amount of ammonia, and 4g of alcohol; eg: making a ~50% solution; I evaporated the remaining water at 90C. There are obviously some impurites left, but the crystal still looks decent. A small odor came off while drying at 98C which was not ammonia smell; I don't know if I lost a small amount of nitrate, ammonia, alcohol, or a combination. It wasn't very strong, so I don't think I lost much mass. I end up with 0.888 grams of dark blue crystal.



The theoretical yield should be:

K2SO4 * 0.02mol * 174.259g/mol = 3.485g ( sol. W 0.11g/mL 20C --> 31.682g H2O )
Cu(NO3)2 * 0.02mol * 241.60g/mol= 4.832g ( sol. W 1.25g/mL 20C --> 3.866g H2O )

So I ended up with: 0.888/4.832 = 18% yield even after many distillations to increase effective solvent used (like a soxhelet extractor, but done manually ... sigh.)

After re-dissolving the bulk of the precipitate in 500mL water (excluding the blue stuff in the picture above), a large amount of the re-dissolved crystal re-precipitated by the time I got to 400mL of water. I removed that precipitate and set it aside. The remaining 400mL of solution was evaporated to dryness at 90C.

Three layers of crystal formed. First a greenish color on bottom, then a green-whitish color, then a deep blue layer.

Because of the large amount of water, the alcohol should have been long boiled off before evaporation completed. The fact that I got blue precipitate on top made me suspicious that there is a water-alcohol double salt/adduct that is more soluble than copper nitrate in alcohol.

Since there should be absolutely no water in the remaining precipitate, I tried adding straight methanol to it as a test. Sure enough, only the blue crystals dissolved into methanol. Photo, below. The color of the alcohol is blue and not aqua-marine. This bolsters my suspicion that a large part of this last process's problems were because of using both water and alcohol at the same time. Apparently the alcohol can somehow react/adduct with postassium sulfate and copper nitrate when water / water of crystallization is present.




[Edited on 14-2-2018 by semiconductive]

[Edited on 14-2-2018 by semiconductive]

Oh, to the guy above who posted the pics of the pretty copper salts. If you want to purify and find out what it is, see if you can pull some mono crystals off. They are either a rhomboid or something not rhomboid. I believe nitrates form needle-like crystals. Tie a string to the little monocrystal and submerge this crystal in a saturated solution of your mystery salt(s). Nice pics. Here's my cuso4*5h20 monocrystal.

I forgot, that's one of the full size bic lighters, so three inches was modesty.

[Edited on 16-2-2018 by Omniquist]

[Edited on 16-2-2018 by Omniquist]

Hanging out in the library and I saw this in a book on organic chem.. Looks like there's a possibility that you could do the nitrate salt and sulphate salt swap, but it would seem adding a current of so2 gas catalyzes the reaction to completion. However, this is for the chloride, so this could be irrelevant. I just took a screen shot so....here, hope this helps.

Partial Quantitative success: Water only method produces Cu(NO3)2 From KNO3 and CuSO4.5H2O

I've been running the experiment in several variations. I've had some successes, now. The third poster who said Cu(NO3)2 could be produced is correct. Impure Copper Nitrate CAN be produced from solutions in water with evaporation at 65C.

I will report the results from crude (non-recrystallized) starting reagents first. This is most likely what other amateurs would try for a quick test; and this post is dedicated to that set-up.



I dissolved 4g KNO3 ( Spectracide brand stump burner) in 25mL hot R.O. water.
I dissolved 5g CuSO4.5H20 ( Roebic (tm) K-77 brand ROOT KILLER ) in 25mL R.O. water.

There is up to 1% impurity in root killer, and spectracide stump remover is expected to be contaminated with excess potassium oxide, hydroxide, and carbonate from air. I tested root killer sample. Dehydrating 5g on dull red burner yields: 3.180g. Some black carbon was formed. Theoretically dehydration should yield ~3.2g ; therefore 20mg error/5g ~= 0.4% impurity + excess water. This should not affect final yields by much.

I combined the two solutions and placed on a magnetic stirrer at 65C in a 50mL Erlenmeyer flask. I placed silicone tubing above the liquid and attached tubing to an aquarium pump. Gentle air blowing evaporated the warm water fairly quickly (hours vs. days unaided).

After evaporating 25mL of solution, a greenish-opaque crystalline mass formed. I transferred all liquid to another flask, and weighed the crystalline mass with tare weight of flask. I dried the crystalline mass for an hour before weighing.

--------------------------
Result I: Evaporating 25mL solution produces 0.520g of light green crystalline mass.
~0mg of first crystalline mass is soluble in methanol.
--------------------------

I used methanol to test alcohol solubility (HEET, Yellow bottle with blue writing for car gas dryer is methanol, Don't use ISO-HEET, that's Iso-propyl.).

CuSO4 and K2SO4 are insoluble in methanol, but CuNO3 with water of crystalization is very soluble. No mass loss was detected. Greenish precipitate must be some kind of double salt of copper/copper hydroxide and potassium?
I then re-dried the crystalline mass for 20 minutes. (Edit: Fixed a typo in earlier post).

I re-placed the remaining solution back on stirrer in new erlenmeyer flask. I evaporated off more liquid. Again a greenish crystalline mass formed ... but this time it was far more translucent and lighter green.

-------------
Result II: evaporating additional ~12mL of water produced 0.956g of crystalline mass.
26mg of crystal was soluble in methanol.
-------------

Therefore, some Cu(NO3)2 was likely in the second crystalline mass.

I again evaporated more liquid in a clean, tared flask. As distinct blue crystals began forming when around ~6mL of solution was left, I decided to add just enough R.O. water to redissolve them; and then poured off the remaining solution into a clean flask. Then I weighed the remaining greenish crystalline mass.

--------------
Result III -- 1.504g of greenish crystalline mass from ~6mL of evaporation.
methanol solubility was not tested.
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I collected all the crystalline green mass from all the evaporation cycles. I attempted to redissolve the crystalline mass in 50mL of R.O. water with vigorous stirring. The crystalline mass only partially re-dissolved. A concentrated powdery green precipitate remained on bottom of the flask. Heating and prolonged stirring did not change the result. The Greenish copper salt that precipitates is almost entirely insoluble in water. The translucent part does redissolve. I assume that the translucent part must be K2SO4, and that because there is excess potassium base in spectracide stump burner that the greenish salt is a basic copper salt; maybe Cu(I)NO3+OH, or CuOH, etc.

I evaporated the remaining 6mL of solution, and it yielded 4.73g of bluish salt with some layers of green mixed in.

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Result IV -- 4.73g of blue crystalline mass, small green layer on bottom and top of crystal.
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The final CuSO4 salt is still quite impure. The experiment was unfortunately ruined at the next step; I added a small amount of water to see if blue crystal would dissolve and leave green crystal. But they both dissolved appreciably, with green being slower. So, the nitrate salt apparently increases the solubility of the green salt. I set the impure salt aside on the stirrer to re-dry. I then took the solution made only from green precipitate and tried adding vinegar (which would evaporate or react only with bases) and a little more K2NO3. I wanted to see if I could cause the hard to dissolve copper salt to go back into solution and react with excess nitrate . eg: the starting stoichometry is a little off for this experiment, and eI thought excess KNO3 should be easy to separate from sulfate salts so I tried to drive the reaction to completionby acid+KNO3 to drive the copper into solution.

I put it back on the hot plate to evaporate again, when an unfortunate side effect occurred.
I happened to be drying one flask of Cu(NO3)2 on the same magnetic stirrer while evaporating the liquid off another flask. Although the liquid was kept at 65C, the hot plate was actually significantly warmer than that. I don't know an exact temperature, but in order to keep a solution at 65C with fish tank air pump cooling it .... generally the hot plate stays 10C to 15C hotter. It didn't burn my fingers, so it wasn't above boiling point of water. However, I estimate the dry crystals of CuNO3 were at around 75C to 85C. A choking gas suddenly started coming off the dry crystals after they had been sitting at that temperature for about a half hour. No such smell came from the liquid.

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Summary of results:
Total mass of results: 7.71g, total starting mass: 9.00g.
Original CuSO4 mass due to .5H2O = 1.8g, Expected loss to become .3H2O = 0.72g.
Actual mass loss, 9-7.71=1.29g ; and 0.72g expected < 1.29g actual < 1.8g max

Unexpected loss due to drips, outgassing, or green crystalls unknown water content: 1.29-0.72=0.570gram.
Raw Yield: 4.73g of blue crystal, out of maximum of 4.832g of tri-hydrate; = ~98%
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Therefore, I can't prove that nitrate based mass loss didn't happen from the final product, nor separate the issue from water of crystallization of whatever the green precipitate is. This reminds me that wikipeda said 80C can cause breakdown of Copper Nitrate into basic copper nitrate.

Quote: Originally posted by LearnedAmateur

Well there’s a reason why we shouldn’t trust Wikipedia, I’ve got a completely different set of values. When I redissolve my salts, including copper nitrate, I always heat to boiling and have had no issue with decomposition below 100C. Maybe there are impurities affecting the stability?