Sciencemadness Discussion Board

ammon-sulfate-nitrate to nitric acid

Kieron - 24-2-2019 at 15:32

Good morning,
please forgive me my poor english.

I have a double salt of ammonium nitrate and ammonium sulfate (2NH4NO3·(NH4)2SO4). What I want to achieve is to produce nitric acid out of it. Unfortunately I failed so far.

What I tried:
I tried to add sulfuric acid to the double salt, so that the ammonium nitrate part reacts to nitric acid. But no matter how the concentrations were that I used, the solution stayed inert.

So I tried to recrystalize it, but either I did it wrong, or it just doesn't work with this double salt to purify both components.

Then I came with the idea to add NaOH to react both salts NaSO4 and NaNO3 (and ammoniak, I know). NaSO4 would fall out in water were as NaNO3 is soluble. But this is pretty costly so I want to avoid that approach.

Can you give me a helping hand?

Best regards.

hissingnoise - 25-2-2019 at 06:58

Quote:
I have a double salt of ammonium nitrate and ammonium sulfate (2NH4NO3·(NH4)2SO4). What I want to achieve is to produce nitric acid out of it. Unfortunately I failed so far.

The most convenient way to produce NA from your salt is by adding sulphuric acid and distilling the nitric acid from the mixture...

Your English is fine, BTW.

Kieron - 25-2-2019 at 11:18

Thank you for your reply.
Quote:

The most convenient way to produce NA from your salt is by adding sulphuric acid and distilling the nitric acid from the mixture...

That's what I thought, too. But as I wrote, the mixture with 95% sulfuric acid stayed completely silent. Even after a day I could not destilate any nitric acid out of it.
I thought of a the posibility that my mixture has a lime component (it's fertilizer), but that's not the case, too.

RogueRose - 25-2-2019 at 19:58

How are you adding the compounds together? Are you dissolving the fertilizer into the H2SO4 (maybe diluted?)? Or are you dissolving the fertilizer in water and then adding the H2SO4?

I think I had similar results when trying to mix Ca(NO3)2 with H2SO4 (25-30%). I found that the Ca(NO3)2 needed to be dissolved into water (a saturated solution was fine, though more dilute works better) before the H2SO4 was added. Because, in this case CaSO4 is formed, which is SUPER fluffy in suspension, it was difficult to get everything to mix well when either the salt solution or the acid had a high concentration (g's/L).

I would make a saturated solution of your fertilizer. Then I'd try starting with 50-60% NaOH and
add some drops into the solution and see if any Na2SO4 forms. I'd use some small test tubes to test this out,maybe 10-20ml at a time. I'd also try 20-30% NaOH as well. The Na2SO4 should start to crash out once it reaches max solubility. I'd also make the fertilizer solution as cold as you can, as it really reduces the solubility of the Na2SO4. An alternative would be to heat the fertilizer solution up to BP or near BP, add the NaOH solution (equi-molar amounts to fertilizer's), mix well, then put it in the freezer and see how much Na2SO4 falls out.

You should get about 10x as much Na2SO4 in solution at BP than you will at freezing (0C), so it should be pretty noticeable when it crashes out.

Once you see everything is working properly, then you at least know this will work with NaOH and then I'd try doing the same thing with H2SO4.

You need to realize that both the ammonium sulfate and nitric acid are probably very soluble together and I'm wondering if it might not raise the BP of the nitric acid (especially if you have excess H2SO4 in the solution). What temp did you reach when trying the distillation?

Also, when you say "inert", what do you mean? Did you not see bubbled? You know that you will need to add enough H2SO4 to both the ammonia (forming either ammonium bisulfate or ammonium sulfate) from the ammonia nitrate, so you wouldn't see any bubbles or anything.


You might want to look at the solubilities of the two, the sulfate is insoluble in acetone and most alcohols, while nitrate is soluble in methanol (not greatly, but if you could distill off and reuse it that could work - IDK how much you need - I've done this before)
https://periodic-table-of-elements.org/SOLUBILITY/ammonium_n...
https://periodic-table-of-elements.org/SOLUBILITY/ammonium_s...

Herr Haber - 26-2-2019 at 04:13

What temperature did you reach during distillation ?

All other parameters are interesting too because there's no reason it shouldnt work.

hissingnoise - 26-2-2019 at 06:29

Quote:
That's what I thought, too. But as I wrote, the mixture with 95% sulfuric acid stayed completely silent. Even after a day I could not destilate any nitric acid out of it.

The only possible scenario I can think of is that because of the quantity of sulphate already present you may have added a large excess of H2O4 ─ that would dehydrate NH4NO3 to form nitroamide, an unstable compound that decomposes to N2O at 150°C.

Nitroamide can prevent the formation of NA in solution.

Careful addition of water can hydrate nitroamide (NH2NO2) so that NA then comes over at 90 ─ 120°C.

markx - 26-2-2019 at 23:12

Quote: Originally posted by Kieron  
Good morning,
please forgive me my poor english.

I have a double salt of ammonium nitrate and ammonium sulfate (2NH4NO3·(NH4)2SO4). What I want to achieve is to produce nitric acid out of it. Unfortunately I failed so far.

What I tried:
I tried to add sulfuric acid to the double salt, so that the ammonium nitrate part reacts to nitric acid. But no matter how the concentrations were that I used, the solution stayed inert.

So I tried to recrystalize it, but either I did it wrong, or it just doesn't work with this double salt to purify both components.

Then I came with the idea to add NaOH to react both salts NaSO4 and NaNO3 (and ammoniak, I know). NaSO4 would fall out in water were as NaNO3 is soluble. But this is pretty costly so I want to avoid that approach.

Can you give me a helping hand?

Best regards.


Are you sure your "double salt" even does contain a reasonable amount of the nitrate component? The apparent inertness upon treatment with highly concentrated sulfuric acid suggests it might be mostly ammonium sulfate.
You can do a small scale extraction with methanol, as mentioned, to see if it can be separated into components and in what approximate ratio the components are present in the mixture.

AJKOER - 3-3-2019 at 05:45

One can convert NH4NO3 to Mg(NO3)2 by warming with a suspension of Mg(OH)2 or possibly MgCO3.

Then, as I noted previously:

Quote: Originally posted by AJKOER  
No need for any acid to make dilute HNO3 with a few easy steps from cheap available chemicals.

First, mix up a concentrated solution of MgSO4 (Epsom salt) with KNO3 (EDIT: not a hydrate, that is my NaHSO4.H2O) from stump remover. Freeze till the K2SO4 formation is evident. Decant removing K2SO4, dilute and refreeze to obtain Mg(NO3)2•6H2O. Remove crystals, heat and collect vapors of dilute HNO3. Reactions:

MgSO4 + 2 KNO3 --> Mg(NO3)2 + K2SO4

Mg(NO3)2 + 6 H2O = Mg(NO3)2•6H2O

Mg(NO3)2•6H2O --heat to over 130 C--> Mg(OH)NO3 + HNO3 + 5H2O

Per a source (https://chemiday.com/en/reaction/3-1-0-7303 ), to quote:

“The thermal decomposition of hexahydrate nitrate magnesium to produce magnesium hydroxide-nitrate, nitric acid and water. This reaction takes place at a temperature of over 130°C.”

Interestingly, Wikipedia presents an alternate path on the thermal decomposition of magnesium nitrate, likely not referring to the hydrate:

2 Mg(NO3)2 → 2 MgO + 4 NO2 + O2

Reference: see https://en.wikipedia.org/wiki/Magnesium_nitrate

Note, my personal experience, as reported on SM, extends to aqueous Mg(NO3)2 preparation, and once on evaporation upon standing in air, the creation of a very hygroscopic salt. As such, I doubt if there is an easy path to anhydrous magnesium nitrate, other than discussed below involving NO2, so the Wikipedia path to NO2 appears circular in all likelihood.
[Edited on 10-12-2017 by AJKOER]

--------------------------------------------------------

For those wanting to avoid toxic fumes, an alternate but still potentially dangerous path which is theoretically interesting (especially for those OK with energetic experiments) would be to convert NH4NO3 to HNO3 still absence the use of acids or other not readily available or expensive reagents may be the following but leading to a transition metal contaminated nitric acid product:

> Apply electrodes of copper to NH4NO3 dissolved in ferrous rich tap water.

> Continuously pump (using an air pump, my choice is a leftover one employed for a fish tank) into the electrolysis cell above.

> As the classic warning of explosion hazard may be accurate for electrolysis of acidic ammonium salt solutions (see all my comments and links at http://www.sciencemadness.org/talk/viewthread.php?tid=18912 ) creating unstable NH4NO2, please perform this experiment in small amounts, monitor pH, use a plastic (as glass is not desirable in a possible explosion scenario) vessel not readily attacked by dilute HNO3 and apply other safety measures (as normally taken in energetic experiments).

I speculate that this variation of the classic electrolysis MAY not result in an explosion assuming the conversion from added oxygen/air to the superoxide radical anion (.O2-, or present as .HO2 at pH < 4.8) by solvated electrons generated in the electrolysis (O2 + e-(aq) --> .O2- (aq)). This could be followed by the reaction of the problematic NO2- with H2O2 (from .HO2 + .HO2 --> H2O2 + O2) forming nitrate (or, via other pathways, see, for example, reactions cited at https://www.tandfonline.com/doi/pdf/10.1271/bbb.64.1751 or, I would argue Electro-Fenton created hydroxyl radicals and H2O2 see https://www.hindawi.com/journals/ijp/2017/8528063/).

Basic logic of underlying process, attack of NH4+ which undergoes the equilibrium reaction:

NH4+ = H+ + NH3

Note, standard electrolysis of concentrations producing ammonia is recommended only to take place in alkaline media at low temperatures (like 25-60°C, see for example http://folk.ntnu.no/skoge/prost/proceedings/aiche-2005/topic... ).

Also, avoid a chloride presence due to the possible creation of highly explosive NCl3 (see, for example, http://www.freepatentsonline.com/2209681.html ).

[Edited on 3-3-2019 by AJKOER]

hissingnoise - 3-3-2019 at 07:21

Quote:

Mg(NO3)2•6H2O --heat to over 130 C--> Mg(OH)NO3 + HNO3 + 5H2O

“The thermal decomposition of hexahydrate nitrate magnesium to produce magnesium hydroxide-nitrate, nitric acid and water. This reaction takes place at a temperature of over 130°C.”

Interestingly, Wikipedia presents an alternate path on the thermal decomposition of magnesium nitrate, likely not referring to the hydrate:

2 Mg(NO3)2 → 2 MgO + 4 NO2 + O2

The Wiki decomp. reaction is correct ─ nitric acid cannot be made by the route you suggest...

Neither, unfortunately, can it be obtained by electrolysis of nitrate salts as NO3 is reduced at the cathode by hydrogen...

And efficient absorption of NO2 by water isn't really feasible in a home-lab...

AJKOER - 3-3-2019 at 08:08

Quote: Originally posted by hissingnoise  
nitric acid cannot be made by the route you suggest...

Neither, unfortunately, can it be obtained by electrolysis of nitrate salts as NO3 is reduced at the cathode by hydrogen...

And efficient absorption of NO2 by water isn't really feasible in a home-lab...



Actually, the radical reaction by the hydrogen atom radical (not H2) proceeds as:

.H + NO3- = OH- + .NO2

.NO2 + .NO2 = N2O4

N2O4 + H2O = HNO2 + HNO3

which does, indeed, introduces nitrite, but I argue oxygen (from air pump into the system) created H2O2 may convert nitrite back to nitrate.

Note, one of my references also cites the important of Cu(OH)2 presence (here from the corrosion of the copper electrode) in promoting a high nitrate product.

Bottom line, I suspect some tuning (like voltage, temperature, concentrations, running time,...) is likely required to increase yield

Vomaturge - 3-3-2019 at 12:00

Quote:
[Quote] assuming the conversion from added oxygen/air to the superoxide radical anion (.O2-, or present as .HO2 at pH < 4.8) by solvated electrons generated in the electrolysis (O2 + e-(aq) --> .O2- (aq)). This could be followed by the reaction of the problematic NO2- with H2O2 (from .HO2 + .HO2 --> H2O2 + O2) forming nitrate (or, via other pathways, see, for example, reactions cited at https://www.tandfonline.com/doi/pdf/10.1271/bbb.64.1751 or, I would argue Electro-Fenton created hydroxyl radicals and H2O2 see https://www.hindawi.com/journals/ijp/2017/8528063/).


And therein lies the problem. I won't argue against UV lamps, sunlight, or even electrolysis making these radicals, but I doubt you will have enough of them to make the process worthwhile. The reaction may well happen, but if it gives a trace of the desired product in a saturated solution of nitrate or nitrite salt, that's a hollow victory for someone who just wants some nitric for use in other experiments.

I'm not saying radical reactions aren't important; in some conditions they may dominate over other reaction mechanisms. But they don't always.

As a result, I'd suggest that the next time you propose a photolysis or other radical reaction as a means of creating a useful reagent, make an estimation (based on experiments, or references, which you are great at finding) of:

a) reaction conditions , including how much of each reactant, catalyst, solvent, time required to complete or reach equilibrium, etc... If light exposure or temperature are important variables, posting them will be good, too.

b) estimated quantity of the desired product expected to be formed.

c) expected side reactions, including non radical ones (like what happens to ammonia when it's released by electrolysis below the surface of a solution that's supposed to be turning into nitric acid). Even if the undesired reactions leave the desired products untouched, you still need to be aware of side products when trying to extract the desired ones.

This is not meant to be an attack on you, but more a suggestion. It might get people to read your posts and references when they otherwise would have glossed over it and dismissed it as 'another radical reaction post by Ajoker'. It might even help you decide which reaction pathways show promise, and which ones to discount.

As far as the OP's question, i would have thought that distilling that mixture would work. My guess would be that the fertilizer's composition was wrong (either mostly sulphate, or possibly containing some compounds besides sulphate and nitrate)

Heptylene - 4-3-2019 at 08:47

From my experience, adding conc. sulfuric acid to ammonium nitrate produces fumes of nitric acid and heats up so much I have to do the addition slowly (possibly from humidity in ammonium nitrate). if you don't get this reaction, your fertilizer might not contain significant nitrate.

Does it decompose upon heating? Heating a small amount of pure ammonium nitrate generates a lot of gas. If you try used only a few milligrams to be safe.

fusso - 4-3-2019 at 12:24

Quote: Originally posted by Heptylene  
Does it decompose upon heating? Heating a small amount of pure ammonium nitrate generates a lot of gas. If you try used only a few milligrams to be safe.
Spoiler: heating NH4NO3 just decomp it, no bang or explosions will occur:P

Heptylene - 4-3-2019 at 14:50

Quote: Originally posted by fusso  
Quote: Originally posted by Heptylene  
Does it decompose upon heating? Heating a small amount of pure ammonium nitrate generates a lot of gas. If you try used only a few milligrams to be safe.
Spoiler: heating NH4NO3 just decomp it, no bang or explosions will occur:P


That's what I think too. I've done it without problems in the past to determine ammonium nitrate content of fertilizer. But IF it ever explodes (maybe due to a sensitizing impurity...) I'd rather it be with a few milligrams that a few grams. I don't pretend to know for sure it won't explode and neither should you.

fusso - 4-3-2019 at 15:09

What if you use excess H2SO4? Did you try this?

Kieron - 21-3-2019 at 12:06

Hi folks, thanks for the input.

I'll try your solutions, take measurements and came back with the results.
Sorry to be so late, but I was in surgery for a week.