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Author: Subject: Interesting Resultant Solution from NO2 Generator Reaction

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[*] posted on 13-4-2016 at 10:40
Interesting Resultant Solution from NO2 Generator Reaction

Buckle up, this is a long post. Sorry in advance but I'm just trying to be thorough in describing my observations.

A while ago I was fiddling around with a NO2 generating reaction that I performed in order to bubble the gas through chilled distilled water or even some weak hydrogen peroxide in order to make some weak nitric acid. I'm sure many of you have heard of the reaction I performed, and I know Nurd Rage did a video on it, which is where I got the idea. So my reagents were copper metal, HCl, and KNO3.

I decided to do the stoich and estimate reactant quantities myself, but I went at it a little too haphazardly and didn't quite think through it enough. So my wrong stoich was as follows: $$ 1\space Cu{(s)} \space +\space 4\space HCl{(aq)}\space +\space 2\space KNO_3{(aq)} \rightarrow \space 2\space NO_2 {(g)}\space +\space 2 \space KCl{(aq)}\space +\space 1\space CuCl_2{(aq)}\space+ \space2\space H_2O{(l)}$$
After throwing this bit of science at the wall, I got some interesting results. An incredibly deep green solution that often results from an extremely concentrated copper nitrate solution, and not copper chloride. That was my first indication that it would be silly for CuCl2 to form. Then once I actually thought about the chemistry occuring instead of slinging things into a flask, I realized how this reaction is similar to that of nitration when using a nitrate salt, H2SO4, and whatever you want nitrated. And the solution looked identical to the product of reacting copper and nitric acid (who'd a thunk). This made me realize that indeed Cu(NO3)2 is produced, not CuCl2.
So the non-moronic stoichiometric reaction I formulated was:
$$ 1\space Cu{(s)} \space +\space 4\space HCl{(aq)}\space +\space 4\space KNO_3{(aq)} \rightarrow \space 2\space NO_2 {(g)}\space +\space 4 \space KCl{(aq)}\space +\space 1\space Cu(NO_3)_2{(aq)}\space+ \space2\space H_2O{(l)}$$
This time I doubled the amount of KNO3 in order to account for the fact that it is Cu(NO3)2 forming. But I still got a very similar result, and likely one reason for this was I did not add enough water to fully dissolve the greater amount of KNO3 in the colder temperatures in which I was doing this reaction. So after the reaction stopped evolving gas, which ended up being a rather little amount, the solution still smelled very strong of HCl and had some undissolved KNO3 left in the bottom of the flask.

Given the very cold temperatures in which I often did this reaction, is it possible that some of the NO2 gas remained dissolved in solution? Once I heated up the solution in an attempt to dissolve remaining KNO3, a noticeable amount of characteristically brown NO2 gas seemed to form inside the flask. I would attribute this to simply enough heat being supplied to drive the reaction forward again, but this would happen whenever I even brought the flask from the cold garage to indoors, and once I took it back out and stoppered the flask the brown color receded. Another theory I have is that if there was plenty of moisture in the air the NO2 could have reacted with it, causing the color to dissipate and the rather acrid odor.

After a while I began repeating the reaction not to generate the gas but rather for the solution. And furthermore, the crystals that result from it. So far I've gotten a wide range of crystals. But considering the conditions of the solution (relative degree of acidity & presence of a myriad of cations and anions) it's no wonder that I've gotten various results. I'm not very good at identifying specific types of crystal structures but I'll do my best. In general the crystals I have obtained from the solution have been small and more cubic in nature similar to KCl, but some have been longer in a manner similar to KNO3. Nearly all of them have been a very wet, limey green though. I believe that they are likely the result of various complex ions forming in solution. Considering that most of all the solution seems to remain acidic after the reaction reaches completion, it would be a very good breeding ground for complexes to form.

Whenever I attempt to heat these crystals in an attempt to dehydrate them, they usually tend to change color and the green tint begins to darken and it also begins to gain a yellow tinge. Usually as you dehydrate a compound it begins to turn lighter or to black in my experience. However, at one point the color stops darkening and then the crystals begin to decompose in a manner akin to copper nitrate. So much of the behavior of copper nitrate is there, just not the color or structure. On occasion as some blue crystals form that are likely just copper nitrate, but they are much more infrequent. I believe that in this solution a copper hexahydrate ion could initially form under the acidic conditions, and it is not out of the realm of the possible for potassium ions to act as a ligand as ammonium might and thus form a complex with copper. Copper would be acting as the central metal ion, giving the crystal it's major semblance to copper nitrate. But this is pure conjecture on my, rather inexperienced and not very knowledgeable part, so if some of the much more experienced and intelligent chemists that frequent this board could weigh in on that especially I would greatly appreciate it. I know I could and probably should look into the lattice energies of what I believe might be forming to see if it is in the realm of the possible. But since I don't have an exact enough idea yet that becomes rather difficult.

What is also quite odd about this solution in my opinion is what happens when I react it with aluminum. In a very highly concentrated solution, when aluminum foil was added great heat resulted, and a layer of faintly blue and light precipitate formed and rose to the top. Also, a the solution turned blue. No copper sponge resulted as you would expect from a replacement reaction between copper and aluminum. But I think this could probably just be attributed to the fact that the concentration of copper nitrate in solution would decrease.

This solution has perplexed me for a while now, that's why I collected so many observations about it. I do apologize for my lack of quantitative data in advance. I suppose that I should obtain some densities of these crystals. At some point I might bring it into school to see what spectra it absorbs best. Oh well, I might obtain some densities today if I'm able to get some crystals to reform as I've redissolved about 90% of what I had. Any ideas about what crystals could be forming from this solution would be indeed greatly appreciated. And thanks for reading my novella of a post.

[Edited on 13-4-2016 by Chemist_Cup_Noodles]

I'll be honest-- We're throwing science at the wall here to see what sticks. No idea what it'll do.
-Cave Johnson, Portal 2
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