Booze
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Nitric acid diffuculties
I am having serious trouble making nitric acid. I used Nile Red's method using ammonium nitrate, sulfuric acid drain cleaner, and water. Here is the
video: https://www.youtube.com/watch?v=KBeo8nww21g
I did exactly as he did, except with ammonium nitrate. No nitrogen dioxide gas was produced, and I just gave up after the distillation was done. So
what did I do wrong, and is there another way to produce nitric acid that you recommend?
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Texium
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Thread Moved
26-2-2017 at 12:11 |
UC235
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I don't understand what the problem is. NOx should ideally not be produced though some is unavoidable.
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Booze
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See, I want to produce nitrogen dioxide so it reacts with water to produce HNO3.
If I did use potassium nitrate, this is what I think would happen:
KNO3 + H2SO4 -> HNO3 + KHSO4
I used ammonium nitrate, and according to the video I was following it should have still produced nitric acid.
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UC235
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NO2 is a byproduct from decomposition of nitric acid during heating. Simply mixing the nitrate salt and sulfuric acid produces nitric acid, it's just
a matter of removing it from the bisulfate side product by distillation.
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Booze
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Right. I did do that, but it appeared that no nitric acid was being produced, which is why I am asking for help. Because ammonium nitrate did not
work, I was thinking about using potassium nitrate.
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ficolas
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Quote: Originally posted by Booze  | | Right. I did do that, but it appeared that no nitric acid was being produced, which is why I am asking for help. Because ammonium nitrate did not
work, I was thinking about using potassium nitrate. |
Why no nitric acid was being produced? How did you notice?
Amonium bitrate works. Unless what you have isnt amonium nitrate, you are doing something wrong or miss interpreting something
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UC235
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Well, then one of your reagents is not what you think it is. Ammonium nitrate works very well for nitric acid production. In fact, I prefer it to any
other salt because of it's high solubility in water and sulfuric acid, and the low melting point of ammonium bisulfate.
Are you using instant cold packs for your "ammonium nitrate"? If so, are you sure it wasn't urea based (in the majority these days, or a mixture with
ammonium nitrate).
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Booze
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It says ammonium nitrate right on the box of the cold packs. I know there is no nitric because I add bits of copper to both the distillate and the
ammonia bisulphate and nothing. I also tested my drain cleaner and it is definitely sulfuric.
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UC235
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Your crude distillate will not necessarily attack copper at room temperature. It's too dilute to be a potent oxidizer. Try heating it.
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Booze
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Oh. This was a while ago and I have gotten rid of all the products of my experiment. I will try again later today and see if it works. I have also
gotten proper distilling equipment.
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Amos
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Just keep on heating. It's not the happiest compound to distill, especially not when there's excess sulfuric acid is present. Just be patient.
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Booze
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Okay, that is it. I am not going to try to make nitric again. I mixed concentrated sulfuric acid and ammonium nitrate, and 2 minutes in my
distillation a white gas filled the entire apparatus and then it got so violent the joints fell apart and filled my garage with the gas. There is a
bunch of white solid in my apparatus and almost ruined it.
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Corrosive Joeseph
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'nitric acid from calcium nitrate' thread.....................
http://www.sciencemadness.org/talk/viewthread.php?tid=9848
in which blogfast25 links to what l consider to be one of the best Nitric acid vids on youTube................
[Easy Nitric Acid]
https://www.youtube.com/watch?v=7akk5ppJjEw&feature=yout...
/CJ
Being well adjusted to a sick society is no measure of one's mental health
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AJKOER
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If you can obtain stump remover (nearly pure KNO3) and Epsom salt (pure MgSO4), one can prepare Mg(NO3)2 as I did once. Basically, mix and freeze out
the K2SO4:
2 KNO3 + MgSO4 = K2SO4 + Mg(NO3)2
The magnesium nitrate is, I suspect, a good starting point to prepare many other nitrates by adding the appropriate hydroxide, for example, with
ammonia water, leading to aqueous NH4NO3, after letting the Mg(OH)2 precipitate settle:
2 NH3 + H2O + Mg(NO3)2 = Mg(OH)2 (s) + 2 NH4NO3
In addition, per this comment on another forum, one Kumaraswami Sathiavasan, whose holds a MSc in Chemistry, claims that the nitrates of divalent and
trivalent metals, including Mg, Ca, Ba, Pb, Cu, Zn, Hg, Al, Fe,..., can be thermally decomposed to release NOx. Examples provided include:
2Pb(NO3)2 = 2PbO + 4NO2 + O2
4Al(NO3)2 = 2Al2O3 + 6NO2 + 3O2
2Hg(NO3)2 = 2Hg + 4NO2 + 2O2
2AgNO3= 2Ag + 2NO2 + O2
but not the monovalent alkali metals like sodium and potassium. Also, careful heating (and not possibly explosive rapid heating) of Ammonium nitrate
yields Laughing gas (N2O).
Link: https://www.quora.com/Which-nitrate-will-decompose-to-give-N...
With respect to Magnesium nitrate, per another source to quote:
"For example, a typical Group 2 nitrate like magnesium nitrate decomposes like this:
2 Mg(NO3)2 --) 2 MgO + 4 NO2 (g) + O2 (g) "
Link: http://www.chemguide.co.uk/inorganic/group1/compounds.html
So, the thermal decomposition of, say, Mg(NO3)2 (even if impure with residual KNO3 or MgSO4) liberates a little O2 and more NO2, which can be scrubbed
with, for example, HOCl (preparation described previously on SM by myself, see http://www.sciencemadness.org/talk/viewthread.php?tid=71477#... ). This quickly leads to HCl and HNO3, with the nitric acid preparation being a
point of interest of this thread.
Note, no strong starting acid is required or expensive reagents (like 30% H2O2). The hypochlorous acid can be prepared with CO2 (from vinegar and
baking soda) acting on an available hypochlorite (like chlorine bleach, which is a mix of water, NaOCl and NaCl) with added CaCl2 (sold as a drying
agent for basements). Unfortunately, the final product with this scrubbing routine is likely dilute HNO3 with a chloride impurity.
[Edit] An interesting idea I had was to alternately scrub the NO2 in water containing N2O (sold in stores in cartridges for beverage injection) in the
presence of pure O2 and sunlight in a closed vessel (the latter allowing the action of solar light).
Logic: the photolysis of aqueous N2O produces hydroxyl radicals, which can further interact with HNO2, NO2- and O2, leading to NO2, NO3-, nitrate
radicals, and possibly other reactive oxygen species, which on net may favorably contribute to rapid HNO3 formation. One could replace the use of
solar light with the action of pulses from a microwave oven.
Note, a related hydroxyl radical approach (that also creates in situ H2O2, but which I view less favorably leading to a transition metal impurity in
the Nitric acid itself and possibly reduced stability with time) would be to air treat a solution containing dissolved NO2 in the presence of say a
very small amount of Fe2O3 (or, better a much smaller amount of easily prepared nano-Fe2O3 from the combustion of FeCl3/Ethyleneglycol/Ethanol)
preferably also in the presence of sunlight.
[Edited on 27-2-2017 by AJKOER]
[Edited on 27-2-2017 by AJKOER]
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Amos
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| Quote: Originally posted by Booze | | Okay, that is it. I am not going to try to make nitric again. I mixed concentrated sulfuric acid and ammonium nitrate, and 2 minutes in my
distillation a white gas filled the entire apparatus and then it got so violent the joints fell apart and filled my garage with the gas. There is a
bunch of white solid in my apparatus and almost ruined it. |
sounds like the conversion to nitric acid was not very complete, and your excessive heating caused the ammonium nitrate to decompose, forming nitrous
oxide and water:
NH4NO3 → 2 H2O + N2O
The water thus produced only causes even more heating when it contacts sulfuric acid; sounds like a chain reaction. Nitric acid is not at all
difficult to make if you approach it correctly. If you don't absolutely need fuming nitric acid, start by carefully diluting your sulfuric acid down
to 70 or 80% in water; start with a beaker of water and add acid to the same beaker the whole time, in portions so it doesn't boil over. When it has
cooled down, put it in your distillation flask and add the ammonium nitrate. Mix the whole thing thoroughly and then carefully begin increasing the
heat on your hot plate. Don't aim to start forcing the acid to distill until you can see that everything is dissolved and only liquid is in the flask.
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Booze
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Thanks for all the feedback from you experts. From what I saw on YouTube, they made a nitrogen dioxide generator to condense it with water to yield
nitric acid. I will take whatever nitric I can get and maybe fractional distill it if I need to. I will probably try to not use ammonium nitrate
anymore and instead use KNO3 or heat a metal nitrate as AJOKER suggested.
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Amos
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| Quote: Originally posted by Booze | | Thanks for all the feedback from you experts. From what I saw on YouTube, they made a nitrogen dioxide generator to condense it with water to yield
nitric acid. I will take whatever nitric I can get and maybe fractional distill it if I need to. I will probably try to not use ammonium nitrate
anymore and instead use KNO3 or heat a metal nitrate as AJOKER suggested. |
If you have any use for ammonia, there's a route to nitric acid I used to perform before I had cheap potassium nitrate.
Heating sodium hydroxide and ammonium nitrate together with a little water produces ammonia gas and sodium nitrate; so I used to do this and bubble
all of my ammonia into cold water and prepare solution that way. The stuff leftover in the reaction vessel can be dissolved in water, filtered, and
slowly evaporated to give crystalline sodium nitrate, which I used for making nitric acid. I do highly recommend a suckback trap for the ammonia
generator, though.
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highpower48
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Maybe wrong but potassium nitrate is very cheap, lots cheaper than cutting open cold packs. I'm not sure but isn't most of the cold packs urea or I
believe calcium ammonium nitrate. I think I remember members saying they were having issues with the cold packs.
To make nitric acid: just mix correct amounts of potassium nitrate, water and sulfuric acid. Then distill. I'm a new comer to the forum, so I may be
incorrect on some of this info.
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Amos
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The cold packs at Walgreen's, a very common pharmacy chain in the US, are prilled ammonium nitrate, nearly pure save some prilling agent.
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AJKOER
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| Quote: Originally posted by Booze | | Thanks for all the feedback from you experts. From what I saw on YouTube, they made a nitrogen dioxide generator to condense it with water to yield
nitric acid....... |
First, I start with some educational references beginning with an old source from Mining and Engineering World, Volume 37, p. 996, link: https://books.google.com/books?pg=PA996&lpg=PA996&dq... , to quote:
“ In oxidation chambers, 500° to 50° C., 2NO + O = 2NO2. In absorption chambers, where it comes in contact with H2O, 2NO2 + H2O = HNO3 + HNO2.
Nitrous acid breaks down in the presence of water, giving HNO3 and NO2 thus: 3HNO2 + H2O = HNO3 + 2NO + 2H2O. The NO again reacts with O and the above
series of reactions is repeated. ”
Another source is from Atomistry.com on HNO2 (link: http://nitrogen.atomistry.com/nitrous_acid.html ), to quote:
“The aqueous solution of nitrous acid is blue in colour, which quickly fades with evolution of brown fumes, leaving a solution containing only
nitric acid:
3HNO2 ⇔ HNO3 + 2NO + H2O.
The decomposition of nitrous acid is unimolecular, and the strongest solution which can be obtained at 0° C. is 0.185N, prepared from the
decomposition of barium nitrite with dilute sulphuric acid.
The aqueous solution is more stable when kept at low temperatures and small concentration, and also when under pressure of nitric oxide. The
decomposition of a cold dilute solution follows the reaction
3HNO2 ⇔ HNO3 + 2NO + H2O,
whereas stronger solutions at higher temperatures decompose according to the equations
2HNO2 ⇔ N2O3 + H2O ⇔ NO + NO2 + H2O.
Other factors also influence the decomposition, such as agitation, surface area, and presence of nitric acid. “
So the industrial process, which involves the use of ‘towers’ (which I suspect are constructed of steel, which could introduce small amounts of
the transition metals Iron and Chromium to the reaction mix) to process the action of air/O2 on NO2 dissolved in water, appears to rely on the
reaction of NO2 with water, Nitrous acid's subsequent decomposition, and the gas phase reaction of NO with O2, via the reactions:
3 [ 2 NO2 + H2O = HNO3 + HNO2 ]
3 HNO2 ⇔ HNO3 + 2 NO + H2O
2 NO + O2 = 2 NO2
For an implied net reaction of:
4 NO2 + 2 H2O + O2 = 4 HNO3
The industrial method is actually, in my opinion, a hard path to follow, as first, the air/NOx gases are under pressure, and one needs a gas phase
reaction between NO and O2 to form NO2. Since the decomposition of HNO2 to HNO3 + NO is an equilibrium reaction, with increasing Nitric acid
concentration, the equilibrium shifts unfavorably towards HNO2 (all other things being equal). Next, one must concentrate the amount of NO2 in water
as only the more concentrated Nitrous acid is unstable, otherwise the so called net reaction cited above is inaccurate. In fact, the cool starting
solution with little NO2 produces dilute HNO2 (and HNO3), and under these initial conditions, Nitrous acid is more stable. However, in reality,
surface area contact in the towers could accelerate decomposition as would transition metal impurities via, for example, where I start with the so
called metal auto-oxidation reaction of molecular oxygen by ferrous:
Fe2+ + O2 = Fe3+ + •O2- (Source: See Table 1. p. 10 at https://pdfs.semanticscholar.org/b017/68c3b9a544d3357ab27198... )
HO2• + Fe2+ + H+ = Fe3+ + H2O2 (at pH < 4.8)
2HO2• → H2O2 + O2
Or: Fe2+ + 2 H+ + •O2- = Fe3+ + H2O2 (at pH > 4.8)
Fe2+ + H2O2 → Fe3+ + OH− + •OH (best at pH < 4)
Fe3+ + H2O2 = Fe2+ + H+ + HO2•
Fe3+ + HO2• = Fe2+ + O2 + H+
•OH + H2O2 → HO2• + H2O
•OH + Fe2+ → Fe3+ + OH−
Where the in situ formation of H2O2 could react with HNO2 creating Nitric acid:
HNO2 + H2O2 → HNO3 + H2O
Or, possible radical interactions favoring the direct or eventual formation of Nitric acid, including for example:
•OH + NO2 → HONO2
•OH + NO → HONO
•OH + NO2- → OH- + NO2
•OH + •OH → H2O2
There is also an important reaction to bio-chemists of superoxide and NO forming peroxynitrite (ONOO−, as a reference, please see, for example, http://www.sciencedirect.com/science/article/pii/S1054358908... ):
•O2- + NO --> ONOO-
Any peroxynitrous acid (ONOOH) formed is expected to release •OH and the NO2 radical as oxidizing intermediates (see, for example, http://pubs.acs.org/doi/abs/10.1021/ja982887t ). Also, in more acidic conditions:
•HO2 + NO --> HNO3 (see http://www.atmos-chem-phys.net/8/4061/2008/acp-8-4061-2008.p... )
There are also negative reactions detracting from HNO3 yield (and I would argue long term stability of the Nitric acid) including, for example:
Fe2+ + HONO2 → Fe3+ + NO2− + •OH
Fe2+ + NO2 + 2 H+ → Fe3+ + NO + H2O (Source: See Table 1 at https://pdfs.semanticscholar.org/b017/68c3b9a544d3357ab27198... )
Fe2+ + NO + H+ → Fe3+ + HNO
Finally, all of the above system reactions have parallel surface based reactions with iron oxides at a much wider pH ranges (reference, please see,
for example, Eq. 10 at https://www.hindawi.com/journals/ijp/2012/801694/#B73 ), where the transition metal oxides may be on the surface of the metal processing towers.
Also, the presence of solar light would help reduce ferric to ferrous and therein promote the yield of HNO3.
My supposition from all of the above is that, for example, simply passing air into dilute aqueous NO2 in an open vessel may not significantly increase
the Nitric acid content absence transition metal impurities (which may not be otherwise desirable).
[Edited on 28-2-2017 by AJKOER]
[Edited on 1-3-2017 by AJKOER]
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