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AJKOER
Radically Dubious
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With respect to scrubbing NO2, why not pass the target gases over heated Fe2O3, creating Iron nitrate in the process? This is a general technique for
creating anhydrous nitrate to quote Wikipedia on NO2 (https://en.m.wikipedia.org/wiki/Nitrogen_dioxide):
"NO2 is used to generate anhydrous metal nitrates from the oxides:[9]
MO + 3 NO2 → M(NO3)2 + NO "
where the harvesting of the formed anhydrous nitrate(s) may paid for the scrubbing operation.
Here are my prior comments on SM upon my successful small scale scrubbing of the problematic nitric oxide using NaClO/HOCl. Apparently, the best pH is
from 4 to 7, and not more alkaline. Please see "Oxidation of Nitric Oxide in Two Stage Chemical Scrubbers Using DC Corona Discharge" available at https://www.google.com/url?sa=t&source=web&rct=j&... which presents some comparative views on many existing routes.
However, some recent research per Atomistry.com on nitric oxide (link: http://nitrogen.atomistry.com/nitric_oxide.html ), suggests yet another possible simple route to me. To quote:
"The absorption of nitric oxide by aqueous solutions of various salts has been the subject of much investigation. In the case of ferrous salt
solutions the solubility of the gas increases with the concentration of the solution. The limit is reached when the proportions of iron to nitric
oxide are in the ratio 1:1, both in aqueous and alcoholic solutions. It is assumed that unstable chemical compounds are formed of the type FeSO4.NO,
FeCl2.NO, etc., but the ready dissociation of such compounds under the influence of heat indicates only a feeble combination. Usher has investigated
the freezing-point of such solutions, and finds that neither the freezing-point nor the pressure of the nitric oxide remained constant, and hence no
conclusion can be drawn as to the nature of the compound FeSO4.NO. The absorption of nitric oxide by bivalent salt solutions of nickel, cobalt, and
manganese is of a similar nature. Ferric salts also absorb nitric oxide readily, as also do many metallic and non-metallic halides. Nitric oxide
dissolves in solutions of copper sulphate, producing a violet unstable compound, CuSO4.NO. "
Namely, direct absorption into an aqueous ferrous or ferric salt soution.
I also came across some interesting chemistry surrounding the production of peroxynitrite as yet another possible method which arises from the
reaction of nitric oxide radical (NO.) with a superoxide radical, per the reactions below:
NO. + 02.- -------) ONOO-
ONOO- + H+ -----) ONOOH
ONOOH ------) OH.+ NO2.
See https://www.google.com/url?url=http://scholar.google.com/sch... which I find interesting based on a possible ongoing photocatalytic source of the
superoxide radical via a suspension of sayTiO2 as to quote:
"A luminol chemiluminescence (CL) probe method was successfully applied to the investigation of the superoxide radical (O2−˙) formed on
photoirradiated TiO2 powders. "
Reference link: http://pubs.rsc.org/en/Content/ArticleLanding/2002/CP/B10844...
which, unlike prior methods, does not require replenishing of the oxidizer and thus may be able to better address a continuous scrubbing issue.
[Edited on 15-7-2015 by AJKOER]
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annaandherdad
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Quote: Originally posted by Boffis  |
Molecular Manipulations: did you dry yours prior to condensing it? How does the presence of wáter effect the condensed product, do you get ice
crystals in the liquid NO2 or is ice soluble in liquid NO2? |
The book on nitric gases in the forum library gives some information about this. Water floats on liquid NO2, and of course the NO2 reacts with it to
make HNO3 and NO. But unlike the case of passing NO2 gas through water (or over it in a tower) apparently in the presence of oxygen this combination
of liquids is capable of reaching any concentration of HNO3, up to 95% or more.
I'm in the process of trying this out. I made NO+NO2 by HCl+NaNO2 (the latter very cheap on Amazon), and fed the combination of gases into a RBF with
another tube leading in from a balloon filled with oxygen. I got the idea for doing this from Woelen's description of the NO+O2 reaction. I put
the flask in an ice bath, no salt, just at 0 degrees. As the NO reacts with the oxygen and the NO2 condenses to a liquid, the vacuum sucks more
oxygen from the balloon as needed. The balloon is mylar, so it holds the oxygen at almost zero pressure (above atmospheric). It's a fairly big
balloon, I put maybe 50 liters of oxygen in it.
Here are the photos: http://bohr.physics.berkeley.edu/lab/lab1.html
The blue thing in the first photo is the mylar oxygen balloon. The second photo shows the liquid NO2. Tomorrow I'm going to add some water (again
with the oxygen from the balloon) and see what I get.
Any other SF Bay chemists?
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battoussai114
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It seems the reaction of NO2 with H2O2 proceeds rapidly if booth are in gas phase: http://ntrs.nasa.gov/search.jsp?R=19730009428
By quickly glossing over the article it seems it might be an option for removing the gas in your case... passing the NO2 stream by a vessel with
heated H2O2 and then condensing the nitric acid in another flask. The nitric acid formed is a useful reagent. But before trying you'd better check the
vapor pressure of H2O2 in different temps for designing the equipment.
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annaandherdad
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Continuing from above....
Added some cold water to my flask of NO2/N2O4, with a connection to the oxygen balloon which the flask was free to draw on. The book, "Absorption of
Nitrous Gases" in the library speaks of bubbling oxygen through the mixture, but I didn't want to do that in order to keep down the fumes, so I just
sloshed the mixture around. It definitely started drawing oxygen from the balloon, and I figure I was making nitric acid because it started to attack
the rubber stopper pretty badly.
Kept this up for a while then let the thing sit for 4 hours in an ice bath. Then I poured off the liquid into a graduated cylinder. It showed two
layers, both red brown but different shades. I believe the upper layer was RFNA, that is, a higher concentration than can be achieved by bubbling NO2
through water. I say this, not because I measured the percentage of HNO3, but because I tested it on copper. There was no or very little reaction at
first, but when I added some water it instantly started a vigorous reaction with the liquid turning blue.
Here is what the book says:
"By adding an excess of liquid nitrogen tetroxide to water,
the concentration of nitric acid present may be made to reach
any desired figure, even at ordinary temperatures (less than
20° C.)- Furthermore, by bubbling oxygen into the mixture,
the nitrogen trioxide present is slowly oxidized to the tetroxide
so that the ultimate product may consist of nitric acid containing
dissolved nitrogen tetroxide." (Webb, Absorption of Nitrous Gases, p. 122).
I don't think this method of making HNO3 is competitive with the saltpeter route, but I was curious to try it. However, dealing with NO2 is a pain,
and I don't have a fume hood (just an open garage, a fan, and the great outdoors).
Any other SF Bay chemists?
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franklyn
International Hazard
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I thought it was common knowledge Urea and NO2 mutually annihilate in solution.
One mol of Urea destroys two mols of nitrous acid , the reason this is used to purify nitric acid.
(NH2)2CO + 2 HNO2 => CO2 + 3 H2O + 2 N2
2 NO2 + H2O => HNO3 + HNO2
On balance you'll also get some Urea Nitrate as a product
(NH2)2CO + HNO3 => (NH2)2CO•HNO3
.
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an Internet Forum In case you wonder W T F ! is going on here
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Boffis
International Hazard
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@annaandherdad. These are a really interesting series of experiments and thank you for sharing them with the community.
I notice in the pictures there is a drying tube, can I ask what is in it?
Another idea occurred to me, could you use 30% H2O2 to get rid of the NO2 from the nitric acid or how about condensing the NO2 directly into ice cold
H2O2 solution. With 30% H2O2 you should end up with >61% nitric acid.
The next question is could you use waste NO2 to up grade azeotropic nitric acid by mixing liquid NO2 with Ice cold nitric acid and then exposing it to
oxygen? Or is this just wishful thinking.
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annaandherdad
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| Quote: Originally posted by Boffis | | @annaandherdad. These are a really interesting series of experiments and thank you for sharing them with the community. |
Thanks!
It has sand in it. I plugged the hole with a little glass fibers. The purpose is to remove the little droplets that are produced in the reaction
making the NO2. I can't bubble it through water for this purpose, sand works very well. I've used the same technique with HCl gas.
| Quote: Originally posted by Boffis | | idea occurred to me, could you use 30% H2O2 to get rid of the NO2 from the nitric acid or how about condensing the NO2 directly into ice cold H2O2
solution. With 30% H2O2 you should end up with >61% nitric acid. |
I haven't tried H2O2 but the point of these experiments was to get a higher concentration of nitric acid than is available by just bubbling NO2
through water (or even using an absorption column). The maximum concentration obtainable this way is less than the 68% azeotrope (I'm recalling this
from the book on nitrous gases, cited above). But I hadn't tried to calculate the percentage one would get from concentrated H2O2. It's an
interesting idea. If need be, one could obtain more concentrated H2O2 (for example, by freezing). I got my oxygen from H2O2, one might have said
that it would have been easier to react the NO2 directly with the H2O2, except as I explained I was interested in getting more concentrated HNO3.
| Quote: Originally posted by Boffis | | The next question is you use waste NO2 to up grade azeotropic nitric acid by mixing liquid NO2 with Ice cold nitric acid and then exposing it to
oxygen? Or is this just wishful thinking. |
Yes, it sounds to me like that would work perfectly well, and would be a good economy on the oxygen, which is some trouble to make.
I have another comment to make about this experiment. I explained previously that my nitric acid behaves as concentrated in terms of its action on
copper (passivating until water is added).
But today I tried another test, namely, pouring some on nitrile gloves. I only had a few ml left, but it was enough. The gloves smoked and fumed
quite a bit, and then burst into flames. Impressive for the neighbors. Ordinary concentration HNO3 (68%) won't do this.
Any other SF Bay chemists?
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annaandherdad
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Hi, Boffis, I repeated your calculation, and got 61% HNO3 from passing NO2 through 30% H2O2, same as you. What interested me is that in this process
overall the water doesn't change, that is, the reaction is
2NO2 + H2O2 --> 2HNO3
(with no water on either side). Once you have converted all the H2O2 to HNO3 then you might think that more NO2 would react with the water to make
more HNO3 (plus NO). If you're using NO2 gas, then Webb says you won't, not above 68.4%, anyway.
Here are the passages from Webb that I was thinking about :
"When nitrogen tetroxide is led into water, nitric acids
containing (60-65 per cent. HNO3 can readily be obtained.2
It was shown by Koerster and Koch3 that nitric acid containing
68.4 per cent. HNO3 can be realized experimentally
under these conditions, and is the maximum concentration
obtainable in this way.
When liquid nitrogen tetroxide is added to a small quantity
of water a separation into two layers occurs. The upper
layer consists of nitric acid containing dissolved nitrogen
tetroxide, while the lower bluish-green layer consists mainly
of nitrogen trioxide
2N2O4 + H2O <--> N2O3 + 2HNO3
If, however, oxygen is bubbled through the mixture, the
latter assumes an orange-yellow colour and the upper layer
is found to contain nitric acid of 95-99 per cent. HNO3—
while the lower layer contains nitrogen tetroxide with 5-10
per cent. HNO3" (Webb, op cit, p. 38)
A comment on the balloon method: You have to have quite pure oxygen in the balloon (in particular, air won't work). That's because if you used air,
after a while the reaction chamber would have only nitrogen and nitrogen oxides in it, and pushing more NO into it would just force those gases back
into the balloon. I was pretty careful to make my oxygen dry and as pure as possible.
Earlier in Webb's book he talks about either using all glass equipment (which I don't have) or using paraffin to protect cork and rubber. At least he
says this when things are exposed to NO2; as for HNO3, I'm not sure if concentrated HNO3 will attack paraffin, perhaps violently.
Any other SF Bay chemists?
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