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[*] posted on 1-6-2014 at 07:44
Making Nitrates


Hello, where I live Fertiliser is extremely hard to come by unless you are a farmer and carry a license due to fertiliser bombs.
I can buy 1kg of KNO3 for 21 USD, that is really expensive.

I saw Plante's thread on the Ostwald process and it looks to be dangerous and difficult.
I do not like the idea of pissing in a bucket to retrieve 10grams of nitrates.
Any ideas are welcome.
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[*] posted on 1-6-2014 at 08:15


Arc process if electricity cost you nothing, since it is likely to cost more then 20 per kg of nitrate made this way.



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[*] posted on 1-6-2014 at 09:14


Quote:
I can buy 1kg of KNO3 for 21 USD, that is really expensive.

Ebay is always worth a look . . . ?


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[*] posted on 1-6-2014 at 09:51


As far as I know KNO3 doesn't react fast enough to explode. KNO3 is also not a very good fertilizer so I'm wondering if maybe you mean NH4NO3, but I would doubt that's what you mean. Could you tell us where you live?



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[*] posted on 1-6-2014 at 11:28


I bet Canada as i'm in the same situation and they banned ALL nitrates, like you can't even get them in the instant cold packs :p I guess KNO3 can still explode if used in a very solid and very airtight container so the pressure can build up. Anyways, personnaly I found no simple way to make nitrates during all this time.
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[*] posted on 1-6-2014 at 11:59


Ammonium sulfate, until that is banned as well.
http://www.sciencemadness.org/talk/viewthread.php?tid=11388




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[*] posted on 1-6-2014 at 12:13


I am from Northern Ireland, wish I was in Canada :(
Yes, fertilisers are incredibly difficult to come by, and they are urea and in low concentrations.
You need a farmers license to buy them, these heavy restrictions are due to the IRA making fertiliser bombs.
Instant Cold Packs contain Urea in one pharmacy I checked.

Hissingnoise, anytime you link to ebay it is a .ie site, where are you from? Maybe you can tell me where to buy fertiliser down south.
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[*] posted on 2-6-2014 at 11:04


Quote:
Hissingnoise, anytime you link to ebay it is a .ie site, where are you from?

That was the main site, but yes, I'm just down the road from you in Connacht . . .
I bought my KNO3(25kg) yonks ago from a British supplier!
When I do get through it ─ it'll certainly be ebay for me.

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[*] posted on 2-6-2014 at 18:26


I have witnessed the action of dilute H2O2 on aqueous ammonia (cheap household brand including surfactants) in the presence of Copper and a touch of sea salt. The latter serves as an electrolyte as this is centrally an electrochemical reaction with a sprinkle of more standard chemical side reactions.

In accordance with a reference source (see http://www.academia.edu/292096/Kinetics_and_Mechanism_of_Cop... ) apparently one of those side reactions includes the formation of HNO2 (and/or NH4NO2). In my opinion, the latter is visually created within 5 to 10 minutes as one of many side reactions evidenced by its subsequent rapid decomposition via a formation of a column of suds. Actually, I mistakenly left this reaction run unattended once in a plastic vessel to return to a complete mess. Some caution is advised per Nitrous acid's history of sudden, rapid and significant gas release.

I suspect with the timely addition of Na2CO3, some NaNO2 could be formed.

A step closer to nitrate, with nitrite for all.

[Edited on 3-6-2014 by AJKOER]
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[*] posted on 3-6-2014 at 02:44




This method might fill the void:

http://www.sciencemadness.org/talk/viewthread.php?tid=25829
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[*] posted on 3-6-2014 at 03:46


Jock88:

Your reference reminded me of a source than stated PbO2 can oxidize ammonia to nitrate as well. A good reference on PbO2, please see http://lead.atomistry.com/lead_dioxide.html to quote:

"The dioxide [PbO2] likewise combines quantitatively with nitrogen peroxide to form lead nitrate, and also oxidises ammonia with production of the same salt. "

PbO2 can conveniently be prepared by the action of NaClO (Chlorine bleach, or one can use chlorine water or HOCl), NaOH (or MgO in chlorine water) on a soluble Lead salt (for example, Lead acetate, but here is another reference using Lead nitrate http://oxfordchemserve.com/lead-nitrate-pbno32-99-200-g/ ). There is first the formation of Pb(ClO)2 as a precipitate, that, in short order, decomposes leaving PbO2.

Now, one would think that Pb was cheap and available, but I have to resort to a Pb solder alloy available at Radio Shack.

Of course, one must always be mindful and take appropriate safeguards given Lead's toxicity (even fumes/dust from working with Pb salts). Also, perhaps lesser known to some, is the toxic nature of Mn, to quote (link: http://digitalfire.com/4sight/hazards/ceramic_hazard_mangane... ):

"Even though I was aware that manganese is considered a highly toxic material, it took me a long time to realize that my becoming ill might have anything to do with my raku pottery work."
-------------------------------------------------

One might also consider in place of H2O2, KMnO4 or PbO2, using CaO2 that I prepared easily recently as I discussed on SM. Per Atomistry (link: http://calcium.atomistry.com/calcium_peroxide.html ) to quote:

Freyssing and Roche proposed the use of " Bicalzit " to sterilise drinking water. The action is not very rapid, but is complete after two or three hours. According to Hetscli, under certain conditions it [CaO2] is more energetically than hydrogen peroxide, and 0.5 grm. per litre is fatal to typhoid bacilli."

Actually, this not an entirely new preparation as the action of water slowly, or with an acid more rapidly, forms H2O2:

CaO2 + 2 H2O ----) Ca(OH)2 + H2O2

[Edited on 3-6-2014 by AJKOER]
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[*] posted on 11-9-2016 at 08:32


Here is a link to a thread of mine concerned with nitrate and chlorate formation complete with pictures: http://www.sciencemadness.org/talk/viewthread.php?tid=34429#... .
Here is a partial quote concerning nitrate formation:

"I have performed multiple runs of a new path to nitrate based on microwave treatment of an aqueous mixture of HOCl (preparation discussed below), ..........., that was infused with N2O gas (source is an 8 gram of N2O pressurized gas cartridge designed for use in a cream whipper which was employed here) and oxygen enriched air. A small amount of chlorate is also formed which can subsequently be greatly increased via photolysis using the reported photocatalytic properties of the created nitrate. Actual portions employed in a particular run are detailed below along with precautions and photos.

First, some theoretical background to provide understanding, promote safety and good yield.

One interesting paper is "Treatment of N2O impulsed microwave torch discharge " by M.Jasinski, et. al., April 2004 . To quote: "Results of using a moderate–power (several hundred Watts) pulsed microwave torch plasma (MTP) to the conversion of atmospheric-pressure nitrous oxide (N2O) into nitrogen oxides( NO, NO2 and N2O4) are presented." with a claimed efficiency between 70% to 90% depending on the power of the microwave. Link: https://www.google.com/url?sa=t&source=web&rct=j&...

This possible conversion of N2O is also cited in this source, "Focus on Hazardous Materials Research", edited by Leonora G. Mason, to quote : "In our previous study using microwave discharge, N2O could be efficiently decomposed into N2 and O2 at atomospheric pressure [8]. However, when N2O is diluted in N2/O2 mixtures was decomposed by microwave discharge, a large amount of NO was emitted due to fast reactions as N(2D,2P) + O2 --> NO + O link: https://books.google.com/books?id=8MXX01Qw_G0C&pg=PA144&...

So, for the purpose of constructing a net reaction for the gas phase in the radiolysis of N2O diluted in N2 and a O2 enriched mixture in a microwave, I will assume, as an undoubted simplication, a reaction sequence that has one of the larger associated oxygen (which is supplied pre-run) and nitrogen demands:

1/2 N2 + pulse radiation ---> N(2D,2P)
N(2D,2P) + O2 --> NO + O
O + N2O ---> 2 NO

Simplied net reaction ignoring possible benefical formation of any NO2 and several other intermediate nitrogen and oxygen species:

1/2 N2 + O2 + N2O --pulse radiation--> 3 NO

With shaking of the pulsed radiated vessel, a reaction with Hypochlorous acid:

2 NO + 3 HOCl + H2O --> 3 HCl + 2 HNO3

where the scrubbing of the NO is best accomplished between a pH of 4 to 7, and not more alkaline (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&...

Next, apparently N2O in an aqueous setting has somewhat different chemistry on irradiation. Per this source, "Chemistry of Ozone in Water and Wastewater Treatment: From Basic Principles to Applications", on page 229, the author Clemens Sonntag cites the following reaction involving a solvated electron (actually, rescaled by 4, link: https://books.google.com/books?id=Om_TKidEjToC&pg=PA229&...

4 e-(aq) + 4 N2O + 4 H2O ---> 4 ·OH + 4 N2 + 4 OH-

where there is no nitrate formation, and one source notes at elevated pH an associated increase in hydroxl radical activity ..."

The gas phase synthesis proceeded by heating an expandable container equiped with an Aluminum filament (created by twisting Al foil and securing it to the mouth of the vessel) containing air with added oxygen. Upon heating, it is assumed that the N2O/HOCl infused liquid largely liberates its N2O gas permitting the gas phase creation of nitric oxide. As most people have accidentally observed, Aluminum placed in a microwave produces sparking, which apparently can foster NO creation. With the HOCl upon shaking the full expanded vessel after 10 seconds burst in the microwave, results in a quick gas contraction as the NO is scrubbed. This 10 second heating, expansion and shaking contraction (which is followed by cooling in a cold water bath) exercise is repeated until there is little observed further contraction of the vessel after cooling indicating no more NO formation.

Note, per the gas phase reaction above, each mole of N2O employed (with an added 1/2 mole of N2 from air) is expected to form at most (100% yield) 3 moles of corresponding nitrate. As such, each 8 gram canister of compressed N2O (which should produce around 4 liters of laughing gas) is expected to form at most 3*8/(44) = .545 moles of nitrate, not a bad small scale yield.

[Edited on 11-9-2016 by AJKOER]
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[*] posted on 12-9-2016 at 00:27


Nitrites can be purchased as food additive (curing salt). Oxidizing nitrite to nitrate may be an option. Adding H2O2 to a solution of a nitrite and allowing this to stand for a long time? You could give it a try.

Another option is to add an acid to nitrite and lead the resulting NOx gas mix, together with a lot of air, through water, or even better, through 12% H2O2. This yields nitric acid, which can be neutralized to make nitrates.

But be prepared, nitrates made in this way will be nearly as expensive as the $21 price tag for 1 kg of KNO3, mentioned at the top of this thread. Nitrite as starting material is more expensive, especially, because nitrites which can be purchased by the general public must be food grade and that makes them more expensive (I purchased 5 kg of food grade NaNO2 for appr. EUR 30, which already is very cheap).




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[*] posted on 12-9-2016 at 11:40


Quote: Originally posted by woelen  
Nitrites can be purchased as food additive (curing salt). Oxidizing nitrite to nitrate may be an option. Adding H2O2 to a solution of a nitrite and allowing this to stand for a long time? You could give it a try.

Another option is to add an acid to nitrite and lead the resulting NOx gas mix, together with a lot of air, through water, or even better, through 12% H2O2. This yields nitric acid, which can be neutralized to make nitrates.

But be prepared, nitrates made in this way will be nearly as expensive as the $21 price tag for 1 kg of KNO3, mentioned at the top of this thread. Nitrite as starting material is more expensive, especially, because nitrites which can be purchased by the general public must be food grade and that makes them more expensive (I purchased 5 kg of food grade NaNO2 for appr. EUR 30, which already is very cheap).

You could also just add a small amount of nitric acid to a solution of your nitrite salt in peroxide. Then just neutralize with the appropriate base, ie K2CO3.
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[*] posted on 12-9-2016 at 23:04


But the problem is that the OP wants to make nitrates. Nitric acid almost certainly will not be available. If someone has nitric acid, then someone also has nitrates.



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[*] posted on 13-9-2016 at 07:23
Nitrate by Fermentation


A neat idea I was toying with for Aga's last challenge is to artificially mimic the dung and urine heaps used originally to produce potassium nitrate, sort of like poo hydroponics. In the original method, poo was piled into a heap and doused repeatedly with urine. Urea in urine would break down and release ammonia. Over time, nitrosobacter in the heap would oxidize the ammonia to nitrite, which would then be further oxidized by nitrobacter into nitrate. The dung heap was turned over periodically to supply the oxygen needed, and leachate from the bottom of the pile was collected and treated with wood ashes (KOH/K2CO3) and then crystallized to end up with potassium nitrate, mainly for gunpowder.

But, playing with poo and boiling piss is neither fun nor neighbor-friendly. However, it seems that the solid matter in the turds is just acting like a matrix for the nitro/nitrosobacter while providing trace nutrients and porosity for aeration.

All you might need is a large, temperature controlled fermentation tank filled with plain water, with a strong aerator to keep it oxygenated. Cheap ammonium sulfate would supply a culture of nitro/nitrosobacter with food, and over time the tank would essentially become a solution of ammonium nitrate and ammonium sulfate. The solution must be carefully buffered with phosphates and carbonates, according to my research, and nutrients such as iron, calcium, and chloride must also be present.

There is an aquarium product called "Nitromax" which is essentially a buffered starter culture of these bacteria. It is used to keep ammonia levels down in fish tanks by converting it to nitrate. I do not think it is out of our reach to set something like this up and try it. It is much less expensive than the Birkeland–Eyde process, although obtaining product of useful concentration will take a lot of boiling. A 10-gallon stock pot over a bonfire should do the trick though.

http://www.ncbi.nlm.nih.gov/pmc/articles/PMC238513/pdf/aem00...

http://www.drsfostersmith.com/product/prod_display.cfm?pcati...




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[*] posted on 13-9-2016 at 10:50


Quote: Originally posted by Praxichys  
There is an aquarium product called "Nitromax" which is essentially a buffered starter culture of these bacteria. It is used to keep ammonia levels down in fish tanks by converting it to nitrate. I do not think it is out of our reach to set something like this up and try it.
I haven't seen any proof from this century that such a process could be made into a workable method of nitrate production, though by all means try it. A reliable, reproducible, and scalable procedure similar to the one in the quote below could be invaluable. My extremely limited research suggests that the aquarium bacteria are mainly active on the surface of the filter, and are ineffective when used in aquariums without the correct type of filter, so a 'bucket with a bubbler' might not cut it.

Based on the excerpt quoted below, a few years ago, I added some pottery-grade calcium carbonate to a colander full of vermiculite (peat of the right sort was unavailable, and anyway, carbonaceous materials are supposed to promote denitrification), hung it over a bucket, added a small pump to recirculate liquid dripping into the bucket back onto the vermiculite filled colander, and added to the bucket some diesel exhaust fluid (urea in water), diluted to about 0.75% urea, plus a few grams of monoammonium phosphate and a bottle of a similar aquarium product. This apparatus was kept in a dark, warm room, with good air circulation, and the solution tested daily. I did not detect more than a few PPM of nitrates (by colorimetric aquarium test strips) even after days and weeks. This was my first and only foray into microbiological biochemistry, so obviously I have no idea what I was doing or what may have gone wrong. The info in the paper you linked is new to me, though I'd expect the quantity of bacterial solution I used to contain enough micronutrients to support at least a small colony of the bacteria, just like how a pure sucrose solution can be fermented with enough yeast. That may very well have been one of the ways I went wrong.
Quote: Originally posted by The WiZard is In  
Thus Muntz and Laine (Compt. rend., 1905, 141, 861 ; 1906, 142,
430, 1239) impregnated peat with sufficient lime to combine with
the nitric acid formed, and then inoculated it with nitrifying
bacteria and passed through it a 0.75 per cent. solution of
ammonium sulphate (NH4)2SO4, at 30° C., thereby obtaining a I
per cent. solution of calcium nitrate, Ca(NO2)2. The bacteria
would oxidise quickly only dlilute solution of ammonium salts, but
even 22 per cent. nitrate in the solution did not interfere with the
process. Consequently, by sending the ammonium sulphate
solution five times through the peat beds there was finally
obtained a solution containing 41.7 e. of Ca(NO2)2 per litre.

Yield.—6.5 kg. of Ca(NOs)2 in twenty-four hours per cubic metre
of peat. The old saltpetre plantations yielded 5 kg. KXO2 in two
years per cubic metre.


[Edited on 13-9-2016 by zwt]
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[*] posted on 13-9-2016 at 11:57


It sort of makes sense that "artificial" fermentation to produce nitrates was never industrially popular. The Ostwald process was patented in 1902 and is incredibly efficient, and this was right in time for Fritz Haber's 1910 massive improvement on the production of the ammonia raw material.

This manuscript from 1863 shows that the last advancements in KNO3 production and refinement from manure heaps occured during this time. It was only during the last ten years of the 1800s that the germ theory of disease became widely accepted, so the process by which manure was converted to KNO3 was barely understood and certainly not well researched by the time it became economically obsolete. As such, no industrial-scale fermentation methods were ever developed.

Based on the quote from The WiZard, this method seems possible with the right matrix medium. The starting material is cheap, and the yields are apparently fantastic once you get it right. This seems to be the holy grail of home nitrate production, equivalent in utility to a well-designed chlorate cell in light of continuing restrictions on these salts. I will look into this.




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[*] posted on 15-9-2016 at 06:52


Some theory suggests, to create nitrite/nitrate from ammonia, start by employing a hydroxyl radical generating mechanism including photolysis (see http://pubs.acs.org/doi/abs/10.1021/jp0349132 ), fenton and fenton-type reactions involving transition metals (discussed below), radiation (microwave pulse), electrolysis, sonochemistry,...

Then, the action of the •OH on ammonia in the presence of dissolved oxygen (or employ H2O2) is claimed by one reference to proceed as follows:

NH3 + •OH → H2O + •NH2 (see Laszlo extract and link below)
•NH2 + O2 → NH2O2• (aminylperoxyl radical unstable) → NO + H2O (See , "On the aqueous reactions of the aminyl radical with molecular oxygen and the superoxide anion", Table 2.1 at https://www.google.com/url?sa=t&source=web&rct=j&... )

Having formed some NO, in the further presence of a hydroxyl radical source generator and hydrogen peroxide the following reactions, for example, could introduce the formation of NO2:

H2O2 + •OH → H2O + •HO2
•HO2 + NO → •OH + NO2 (see https://www.google.com/url?url=http://scholar.google.com/sch... )

or, as the net of last two reactions equals:

H2O2 + NO ---UV or Fe(++), Cu(+), Co(++),..→ H2O + NO2

And further:

NO2 + H2O → HNO2 + HNO3

which introduces aqueous nitrite that is a better light induced promoter of hydroxyl radicals then nitrate, which is superior to H2O2. In the presence of uv, strong solar light or select transition metals, this postulated reaction chain can likely accelerates the conversion of ammonia to nitrite/nitrate. Interestingly, there are researchers reporting tissue cell damage from the seeming direct action of NO on H2O2 in biological systems (see, for example, "Hydroxyl radical formation resulting from the interaction of nitric oxide and hydrogen peroxide.",by Nappi AJ, Vass E. ) although, in my opinion, such a direct radical reaction need not take place given the enabling effects of light or transition metals. Note, my simple reaction chain consumes and regenerates the very short lived hydroxyl radical which, in essence, could increase its apparent reactivity life span (as measured by random cell collisions leading to cellular damage) having been, in effect, resurrected.

Unfortunately, there are also possible Fenton-type REDOX reactions in the presence of, say, cuprous/ferrous ions (present in tap water) with the any newly created HNO3 (or HNO2) which may also consumes the targeted nitrites/nitrate formation: See, for example, "Fenton chemistry in biology and medicine*" by Josef Prousek, to quote reaction (15) on page 2330:

"For Fe(II) and Cu(I), this situation can be generally depicted as follows [20,39],

Fe2+/Cu+ + HOX → Fe3+/Cu2+ + .OH + X- (15)

where X = Cl, ONO, and SCN. "

I am assuming that the above applies not only to HOONO and HONOO but to HNO3, as well. In the current context, where one would then have:

(1) Cu(+) + HNO3 → Cu(++) + •OH + NO2-

My take on a possible reaction chain:

(2) HNO3 → H+ + NO3-
(3) •OH + HNO3 → H2O + •NO3 (see https://www3.nd.edu/~ndrlrcdc/Compilations/Hydroxyl/OH_169.H... )
(4) •NO3 + NO2- → NO3- + •NO2 (Reference, page 182, equations R24 in "Plasma Kinetics in Atmospheric Gases", by M. Capitelli, et al., link: https://books.google.com/books?id=2ofqCAAAQBAJ&pg=PA182&...

Net (1) to (4): Cu(+) + Cl- + 3 HNO3 → Cu(++) + 2 NO3- + NO2 + H+ + Cl- + H2O

There exists an example cited in a Russian educational site, page 18, #116, unbalanced, based on ferrous sulfide:

FeS + HNO3 → Fe(NO3)3 + H2SO4 + NO2 " (see https://www.google.com/url?sa=t&source=web&rct=j&... )

which agrees on the production of a nitrate, NO2 and an acid.

Those who believe I have presented an overly complex edition of the action of say a cuprous/ferrous with nitric acid should first read, "A kinetics study of the oxidation of iron(II) by nitric acid" by Irving R. Epstein, et al, J. Am. Chem. Soc., 1980, 102 (11), pp 3751–3758, link: http://pubs.acs.org/doi/abs/10.1021/ja00531a015?src=recsys&a... ), where it soon becomes clear my rendition likely errors more on the too simple as I have omitted, for example, speciation considerations.

A similarly destructive REDOX reaction and radical chain may similarly exist for any formed HNO2. So, pH control via addition of NaOH or buffering is likely needed to preserve any formed nitrate/nitrite.

Cited Laszlo reference: "Kinetics and Mechanism of the Reaction of •NH2 with O2 in Aqueous Solutions", by B. Laszlo , Z. B. Alfassi , P. Neta , and R. E. Huie, J. Phys. Chem. A, 1998, 102 (44), pp 8498–8504, DOI: 10.1021/jp981529+, July 15, 1998. To quote from the abstract:

"The reaction of NH3 with •OH or SO4•- radicals produces the aminyl radical, •NH2. Pulse radiolysis and laser flash photolysis techniques were utilized to study the formation of this radical, its absorption spectrum, its reaction with O2, and the mechanism of formation of subsequent intermediates and the mainfinal product, peroxynitrite. The rates of formation of •NH2 and its absorption spectrum are in agreement with previous reports. The reaction of •NH2 with O2, however, was observed to take place much more rapidly than reported before and to involve an equilibrium of these reactants with the aminylperoxyl radical, NH2O2•. The equilibrium is shifted toward completion of the reaction via catalyzed decomposition of this peroxyl radical, and this decomposition affects the observed rate of reaction of •NH2 with O2. The peroxyl radical deprotonates and isomerizes and finally forms NO. In the presence of O2•-, NO is converted rapidly to peroxynitrite, ONO2-. This product, which is stable at alkaline pH, was confirmed by γ-radiolysis of aerated ammonia solutions." Link: http://pubs.acs.org/doi/abs/10.1021/jp981529%2B?journalCode=...

Another more recent source: "Removal of ammonia by OH radical in aqueous phase." by Huang L, Li L, Dong W., Liu Y., Hou H., in Environ Sci Technol, 2008 Nov 1;42(21):8070-5.i .To quote select parts of the abstract:

"In this research work, H2O2 was selected as *OH precursor. The removal of ammonia under 253.7 nm irradiation from low-pressure mercury lamp in the presence of H2O2 was studied to investigate the ammonia removal efficiency by *OH. Results show that the *OH, generated by H2O2 photolysis, could oxidize NH3 to NO2- and further to NO3-. Removal efficiencies of ammonia were low and were affected by initial pH value and ammonia concentration. ....

*NH2, the main product of *OH with NH3, would further react with H2O2 to yield *NHOH. Since *NHOH could not stay stable in solution, it would rapidly convert to NH2O2- and consequently NO2- and NO3-..."

Link: http://www.ncbi.nlm.nih.gov/pubmed/19031904

Also, interesting mention of a target pH of 9.3 under photolysis in a prior thread of mine at https://www.sciencemadness.org/whisper/viewthread.php?tid=11... . In the reference provided there, an important observation is that the 9.3 pH cited is the pH associated with the dissociative equilibrium of ammonia in water:

NH4+ + H2O = NH3 + H3O+ pKa(NH4+)= 9.246

So, apparently free NH3 is much more readily attacked by •OH than NH4+. As such, a good candidate for exploration may be a closed NH3/N2O/H2O2/O2 system under pulse radiation.

[Edited on 16-9-2016 by AJKOER]

[Edited on 16-9-2016 by AJKOER]
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Melgar
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[*] posted on 15-9-2016 at 09:55


Quote: Originally posted by woelen  
But the problem is that the OP wants to make nitrates. Nitric acid almost certainly will not be available. If someone has nitric acid, then someone also has nitrates.

The idea is that once you have a small amount, you could continue making it via the route I described whenever you're running low. The first time you do it, you could also use something like oxalic acid, since oxalate salts would precipitate out first on evaporation. Filter, then titrate to neutralize any remaining acid. Although, I remember once mixing oxalic acid with a nitrate salt, and observing a very slow redox reaction, which resulted in the formation of what I assume were CO2 bubbles. My guess is that the oxalic acid caused a small amount of nitric acid to be present at any given time, and that the nitric acid oxidized the oxalate to CO2. That would generate NOx species, which would probably oxidize more oxalate. But because HNO3 is a much stronger acid than oxalic acid, that reaction would be much slower than the reaction with H2O2.

I wonder if you could just add oxalic acid to NaNO2 or KNO2 in H2O2, leave it sit for a few days, and eventually the whole thing would just neutralize itself to a nitrate salt solution containing some carbonate salts? If the goal is nitric acid, the carbonates wouldn't matter. If it wasn't, they'd fall out first on evaporation, and could be decanted then.

Incidentally, they make nitrate-selective ion exchange resins quite cheaply, and those work particularly well for removing nitrates from a dilute solution that also contains other salts. Those resins would probably work really well for isolating the nitrates generated using many of the ideas being discussed here, like ammonia fermentation, Fenton chemistry, etc.

[Edited on 9/15/16 by Melgar]
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