@ j_sum1 I just finished reading your posts detailing your efforts to synthesize anhydrous NO2 with urea and calcium hypochlorite. For someone who
was experimenting "in the blind" as I was tonight, your work proceeded in a very cautious and logical manner and you managed to resist the urge to
scale up or rush your experiments. All of us have, at one time or another, given in to that "more, bigger, faster" urge to save time. Most of us
have gotten away with it. Some of us didn't. Anyway, I just wanted to point out what you were doing right. Much respect sir.
Now, a few tidbits I'd like to add regarding the synthesis you attempted.
In the first phase you were reacting calcium hypochlorite and urea in the presence of slaked lime. You mentioned that a popping began as you were
mixing the dry reagents. That sounds almost exactly like what happened to me tonight. I believe, in your case, you forgot that your oxidizing agent
was a powerful chlorinating agent as well. Exposing Ca(OCl)2 to either harsh acidic or basic conditions will result in the creation of free chlorine
and/or monoatomic chlorine radcals which would have made short work of the two amine groups on the urea molecule. As soon as each nitrogen added a
third chlorine, its covalent bond with the carbonyl carbon would vanish. Almost immediately two things would occur. The newly formed molecule of
nitrogen trichloride would drift away from the carbon and a chlorine radical would take its place. The same radical halogenation would occur at the
other nitrogen, completely replacing both amine groups with chlorine atoms. Instead of solid urea, you now have gaseous phosgene, an insidious poison
which leaves the reaction mixture, driving the reaction to the right, increasing the rate of phosgene production. Meanwhile, the NCl3 that is
building up in the reaction mixture suddenly notices that it's really hot and the pH is unbearably high. It isn't happy. Any time the local
concentration of NCl3 exceeds the critical threshold, it makes its displeasure known to anyone within earshot.
I'm afraid that urea is just about the worst possible reactant you can use if your goal is make NO2. NO2 reacts vigorously with urea. All the
products of this reaction are colorless gases, nitrogen, water vapor, CO and CO2 which, as in the previous example, leave the reaction.
LeChattelier's principal kicks in and the reaction rate kicks into high gear in an attempt to make more product and reestablish equilibrium between
the forward and reverse reactions. If that weren't bad enough, the water produced by the reaction acts as a catalyst, speeding up the destruction of
NO2. Dropping urea prills into red fuming nitric acid is an easy way to convert it into white fuming nitric acid as urea destroys dissolved NO2 but
does not react with HNO3. When the fizzing stops, you're done.
This reaction made world news last year when it was exploited by Volkswagen in a recent US emissions testing scandal. The level of oxides of nitrogen
(NOX) at the tail pipe had to meet US EPA standards before the company's diesel cars could be imported. Knowing that they couldn't pass the tests, VW
did what anybody faced with the same situation would do. They cheated. Canisters of dry prilled urea were hidden in the exhaust system of the test
vehicles. NOX entering the canisters exited as water vapor, nitrogen and CO2. Since the tests were specific for NOX, the kapooter ignored the
abnormal CO2 emissions. It wasn't until a human being correlated the low NOX/high CO2 pattern that anybody suspected that a respected multi billion
dollar world corporation was trying to slip its pecker into the EPA's pocket. The issue that really pushed the bizarre meter up past eleven was the
fact that the guilty scumbag corporation wasn't even Chinese, it was German.
Anyway, I can think of maybe a dozen ways to make dry NO2 from OTC ingredients other than urea. I assume that air is OTC enough for you? When I get
a chance, I'll post some of them to the other thread.
Snap, Crackle, BANG!
Y71 |
