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Author: Subject: Preparation of ionic nitrites
Lionel Spanner
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[*] posted on 27-1-2023 at 21:52


Having made 7 attempts at it, 5 in tin cans and 2 in a stainless steel saucepan, I can only conclude the Grossmann method for making inorganic nitrites (US Patent 792,515, 1905) which is cited in the wiki, is not useful or reliable. The results vary wildly depending on the surface material of the reaction vessel, and when successful (in tin cans previously used for food) the nitrite was of very low purity, i.e. 40% or less. The two runs in stainless steel produced no nitrite at all; the only product recovered was unreacted sodium nitrate.



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[*] posted on 28-1-2023 at 01:37


I have been pursuing the old molten lead method over the last couple of weeks. 85 g Sodium nitrate and 207 g lead (I used lead fishing sinkers) were melted in a cast iron pot and stirred continuously. An orange oxide begins forming, and the mixture bubbles. After about 30 minutes, the mixture is a thick orange mud and the lead seems to have mostly reacted. Stirring is continued after the heat is stopped to prevent a solid mass forming. Once the mixture solidifies, about 200ml water is added and the mixture left to soak for 15 minutes or so. All chunks should disintegrate.

The mixture is filtered and CO2 is bubbled through the filtrate for a few minutes; a white precipitate forms, which is removed by filtration. The filtrate is then reduced by boiling. When the temperature reaches 125-130 degrees C, heat is removed and the mixture cooled to fridge temperature. The crystals are then filtered and the filtrate reduced again for a second crop. Adding HCl to the salt gives plumes of brown gas. Pretty straightforward, right?

Now it was clear to me from early on that not all the lead had reacted. There were small grains of metallic lead amongst the lead oxide. So I tested the purity of my product as follows. 1 g of my salt was dissolved in a couple of ml of water. 2.5 g silver nitrate was similarly dissolved separately in 5 ml or so. The two solutions were mixed; a white precipitate of silver nitrite formed instantaneously. This was filtered, dried and weighed. If the 1g was pure sodium nitrite, the yield of silver nitrite should be 2.23g.

After one lead/nitrate reduction, my salt mixture was 30% nitrite. I then repeated the procedure with this salt mixture and fresh lead and got it to 50%. A third cycle got me to 80%. For reference I also tested some sodium nitrite isolated from curing salt, which I measured at 95% (although, in this case, I think any sodium chloride contaminant would have formed the insoluble silver chloride which could throw my numbers off).

It's possible that a longer reaction could boost yield, however I'm also aware that sodium nitrite decomposes into the nitrate in the presence of oxygen at high temperatures, so a longer reaction time could possibly be counterproductive. So all in all, to me, the lead method is 'not great'.

While I haven't tried it, the use of silver nitrate could also be used as a purification method if you're happy to buy/make plenty of silver nitrate (you'll need 250g of silver nitrate to separate 100g of the sodium nitrite from nitrate contamination, but you can reclaim most of it!). After reacting as above, the silver nitrite would be mixed with an equimolar amount of sodium chloride in solution to give a precipitate of silver chloride, while sodium nitrite remains in solution. Filter the solid and evaporate the water for your sodium nitrite. The silver chloride can be made back into metallic silver with NaOH and sugar, then reacted with nitric acid to regenerate the silver nitrate.
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[*] posted on 18-2-2023 at 21:28


For what it's worth I've recently tried a couple of methods described in this thread, with no joy.

The reaction of sulphamic acid, calcium oxide and sodium nitrate (as described on the first page) resulted in a lot of water vapour, some nitrogen dioxide, and the precipitation of a non-reducing substance that isn't nitrite or nitrate, and is considerably less soluble in water than simple inorganic nitrites.

The reaction of sodium nitrate and sodium sulphide (as described in Morgan, 1908) produced sodium sulphate, and a hygyroscopic yellow product that was slightly less soluble in water than nitrate or nitrite, and was contaminated with unreacted sodium sulphide - possibly sodium sulphamate.

Given that inorganic homebrewed nitrite is the modern-day amateur chemists' equivalent of the philosopher's stone, I'm very glad that Poland exists, Polish vendors will freely sell sodium nitrite to private individuals, and although it's not cheap, the postage is not ridiculously expensive.

[Edited on 19-2-2023 by Lionel Spanner]




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[*] posted on 2-3-2023 at 16:29


Many years ago I used to work at a cosmetic/toiletry manufacturer that had once used bronopol (2-bromo-2-nitro-propane-1,3-diol) to preserve many of its products, and found some of them turned brown to black due to the reaction between nitrite ions released by bronopol degradation and cocamide DEA, turning the latter into an unstable N-nitrosamine due to an incomplete diazotisation reaction; when attempted with secondary amines, this reaction stops at the nitrosamine intermediate. (As nitrosamines are highly carcinogenic, this was extremely bad news, and the preservative system in those products was soon changed.)

As it turns out, bronopol is much more rapidly hydrolysed to nitrite in aqueous caustic soda at 100 °C. The initial products of the reaction, formaldehyde and 2-bromo-2-nitroethanol, are relatively volatile, boiling at -19 and 83 °C. This could potentially be turned into a useful preparation, though bronopol is hard to come by for amateurs.

Source: Sanyal, Basu, Banerjee. Rapid ultraviolet spectrophotometric determination of bronopol: application to raw material analysis and kinetic studies of bronopol degradation. , J. Pharm. Biomed Anal., 14 (1996), 1447–1453. doi:10.1016/0731-7085(96)01779-7




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[*] posted on 11-3-2023 at 09:29
A major breakthrough!


One of the biggest problems in producing nitrite by reduction of nitrate is its tendency to react with oxygen at the reaction temperature. The solution? Remove oxygen.

I recently got an argon cylinder over the counter from a local welding supply store (Machine Mart) and decided to retry the sodium sulphide reduction method described in Morgan, in the same manner as the molten lead method (reductant added to molten nitrate in portions), while sparging the flask with argon and keeping oxygen out.

It only went and bloody well worked!

I'll need to reproduce and refine the method before providing a full write-up, but this is definitely a viable way to produce nitrites at a small scale. Get in!




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[*] posted on 11-3-2023 at 20:20


Chalk up another victory to the inert atmosphere! First eugenol demethylation, now nitrite.

Lovely choice of reducing agents we have. Sodium sulfide or lead. Any word on nickel carbonyl? :D

But seriously, nice work!




[Edited on 04-20-1969 by clearly_not_atara]
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[*] posted on 12-3-2023 at 00:13


Quote: Originally posted by clearly_not_atara  
Lovely choice of reducing agents we have. Sodium sulfide or lead. Any word on nickel carbonyl? :D

To be fair, sodium sulphide is easily obtainable from photography suppliers (it's the traditional reagent used for sepia toning), and you only need 1 mole of sulphide per 4 moles of nitrate. Plus, technical grade sodium sulphide is hydrated and is a liquid at the temperature the reaction is carried out, so the reaction mixture is uniform, and you don't have the problem of uneven mixing that you'd get with a solid/liquid or solid/solid mixture.

The process described in Morgan was carried out in iron pans, with no mention of an inert atmosphere, which is a recipe for failure (and an awful lot of ammonia.) The chemistry itself was sound, but the process was pretty bad.




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[*] posted on 22-3-2023 at 10:24


I couldn't reproduce this method.

In my initial attempt, I stopped when addition of sulphide started producing small deflagrations at the surface of the molten mixture, likely due to formation of elemental sulphur - this started happening when 60% of the sulphide had been added. However, when I crystallised out nitrite, it was only 60% pure, suggesting an incomplete reaction had taken place. So the second time, I added all of the sulphide.

On dissolving the solidified reaction mixture, it appears the nitrite had been destroyed in the reaction, most likely due to sulphide reacting with nitrite, forming elemental sulphur and ammonia; as no ammonia vapours were evident during the reaction, it was most likely captured as ammonium sulphate. This would explain why although a highly soluble hygroscopic substance was produced on concentrating the mixture to near-dryness under vacuum, it did not look like either sodium nitrate or nitrite (smaller crystals), and the pH of a solution was too low for it to be either nitrite or nitrate (about 5).

Sod the shipping costs, going forward I'm just buying it from Poland.




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[*] posted on 19-6-2023 at 12:32


In the first post he do NaHSO3 + CaCl2 and then heat the resulting Ca(HSO3)2 to decompose to CaSO3 + H2SO3.

I was thinking.

Lets start with sodium metabisulfite.
Na2S2O5 + H2O = NaHSO3

Then
NaHSO3 + Ca(OH)2 = CaSO3 + NaOH + H2O
Can this be done?(i know weak base displace stronger base, not the opposite, but the low solubility of CaSO3 will drive it forward?)

or (this should definitely work)
NaHSO3 + NaOH = Na2SO3 + H2O
followed by
Na2SO3 + CaCl2 = CaSO3 + NaCl

Also about the thermal reduction of NaNO3 + CaSO3, i plan to do it in a crucible(porcelain) with bunsen burner.
Do i heat it till the nitrate/nitrite melt?, is it gonna be hard to tell when it reached completion? or theres some color change?.

[Edited on 19-6-2023 by fx-991ex]

[Edited on 19-6-2023 by fx-991ex]
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[*] posted on 20-6-2023 at 03:51


Quote: Originally posted by fx-991ex  
In the first post he do NaHSO3 + CaCl2 and then heat the resulting Ca(HSO3)2 to decompose to CaSO3 + H2SO3.

I was thinking.

Lets start with sodium metabisulfite.
Na2S2O5 + H2O = NaHSO3

Then
NaHSO3 + Ca(OH)2 = CaSO3 + NaOH + H2O
Can this be done?(i know weak base displace stronger base, not the opposite, but the low solubility of CaSO3 will drive it forward?)

or (this should definitely work)
NaHSO3 + NaOH = Na2SO3 + H2O
followed by
Na2SO3 + CaCl2 = CaSO3 + NaCl

Also about the thermal reduction of NaNO3 + CaSO3, i plan to do it in a crucible(porcelain) with bunsen burner.
Do i heat it till the nitrate/nitrite melt?, is it gonna be hard to tell when it reached completion? or theres some color change?.

[Edited on 19-6-2023 by fx-991ex]

[Edited on 19-6-2023 by fx-991ex]

My only comment is that you will need to work with an inert atmosphere, as nitrite salts are rapidly oxidised to nitrate by oxygen in their molten state.
Good luck!




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[*] posted on 20-6-2023 at 08:18


Quote: Originally posted by Lionel Spanner  
Quote: Originally posted by fx-991ex  
In the first post he do NaHSO3 + CaCl2 and then heat the resulting Ca(HSO3)2 to decompose to CaSO3 + H2SO3.

I was thinking.

Lets start with sodium metabisulfite.
Na2S2O5 + H2O = NaHSO3

Then
NaHSO3 + Ca(OH)2 = CaSO3 + NaOH + H2O
Can this be done?(i know weak base displace stronger base, not the opposite, but the low solubility of CaSO3 will drive it forward?)

or (this should definitely work)
NaHSO3 + NaOH = Na2SO3 + H2O
followed by
Na2SO3 + CaCl2 = CaSO3 + NaCl

Also about the thermal reduction of NaNO3 + CaSO3, i plan to do it in a crucible(porcelain) with bunsen burner.
Do i heat it till the nitrate/nitrite melt?, is it gonna be hard to tell when it reached completion? or theres some color change?.

[Edited on 19-6-2023 by fx-991ex]

[Edited on 19-6-2023 by fx-991ex]

My only comment is that you will need to work with an inert atmosphere, as nitrite salts are rapidly oxidised to nitrate by oxygen in their molten state.
Good luck!


Apparently it can be done without melting the mixture (230C-300C)
Heres a video where it seem to be working well: https://www.youtube.com/watch?v=5BLPoE6Y-ns
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[*] posted on 20-6-2023 at 13:07


If you can actually get it to work, that'd be one hell of an achievement.
As I said: good luck!




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[*] posted on 22-6-2023 at 09:52


Would it work better with potassium nitrate? KNO2 is 10 times more soluble than KNO3.



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[*] posted on 2-8-2023 at 14:54


Since reductive pathways to nitrite generally require high temperatures and produce poor purity products, or result in over-reduction to ammonia, it may be worth taking a cue from the industrial synthesis, namely the reaction between nitric oxide and caustic soda, which doesn't involve any oxidation or reduction at all.

Now nitric oxide itself is hard to come by for amateurs, but here's an idea for an indirect route: prepare nitrosyl sulphuric acid, dissolve it in sulphuric acid, cool the solution to near 0 °C, then add it (slowly and carefully!) to an equally cold hydroxide solution, under an inert atmosphere, to produce a mixture of nitrite and sulphate, that can easily be separated due to the large difference in solubility.

Brauer claims direct addition of water to nitrosyl sulphuric acid produces dinitrogen trioxide, a direct precursor to inorganic and organic nitrites, but this is probably only true at temperatures well below zero.

Nitrosyl sulphuric acid is not trivial to make, as it requires sulphur dioxide to be bubbled through cold fuming nitric acid, but as a non-volatile solid with a melting point of 70 °C it is easier to handle than other simple inorganic nitrosyl compounds, which are highly toxic gases or low-boiling liquids.

[Edited on 2-8-2023 by Lionel Spanner]




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[*] posted on 11-8-2023 at 09:34


Here's what may be a very interesting and relevant paper, describing selective production of nitrite or nitrate by electrochemical oxidation of ammonia with a copper electrode.
https://chemistry-europe.onlinelibrary.wiley.com/doi/epdf/10...

Unfortunately Sci-Hub hasn't indexed it, and I'm not in a position to drop the Britbongland equivalent of $49 just to download it, especially if it turns out to be hot garbage. Would anyone who has access to the Wiley online library be able to share it? Many thanks!




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[*] posted on 11-8-2023 at 23:09


Here it is.
Haven't read everything yet, but the procedure looks amateur friendly.

Attachment: ChemSusChem - 2021 - Johnston.pdf (714kB)
This file has been downloaded 152 times
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[*] posted on 12-8-2023 at 22:32


Quote: Originally posted by Parakeet  
Here it is.
Haven't read everything yet, but the procedure looks amateur friendly.


Fantastic, thank you very much!




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[*] posted on 13-8-2023 at 08:55


HNO3 --> isopropylnitrite!

i'd just like to report a pounding success. as im messing around with bulk dissolution of silver alloy in dilute nitric acid i showed off to one non-chemist what nitric and copper can do, because its so toxic and beautiful, he had some acetone standing around and i added a bit into it, for science- to see if anything interesting would happen. it died down and there was a vague organic nitrite scent, very mild vasodilatory effect, aha.
so i took a bit of copper wires, maybe 15% HNO3 and added about 1mL IPAlcohol to the.. 15mL mixture of acid
let it stand for maybe 3 hours roomtemperature, color changed very mildly due to copper metal dissolution
and, due to IPNitrite being very volatile i can indeed testify that theres a very strongly vasodilatory substance now in my 50mL erlenmeyer flask
so- simply scale this up, figure out how much % acid one can use, what temperature- preferably 50*C so you can immediatedly distill over the isopropyl nitrite and right away react with sodium hydroxide to form our dear sodium nitrite in quite pure form

alternatively ethanol may be used, nitric acid and alcohol may cause a runoff and cause explosive formation thus its ideal to not increase the temperature very much as for, distilling it out as its hard to tell what could set off organic explosives in a glass vessel, perhaps some solvent could be used to extract the nitrites, but that would also risk carrying over an organic nitrate, which would then contaminate the nitrite

its plausible we can go as low tech as HCl + KNO3 + Cu + IPAlcohol --> IPNitrite + NaOH --> NaNO2 + IPAlcohol

update: i mixed 30mL mL IPA with 30mL 62% HNO3, and 90mL H2O, giving 120mL 15% HNO3
at room temperature this caused a slight runoff but clearly giving off some vasodilatory fumes, more dilution or controlling temperature better might be the way to go, seems to work without copper, albeit more difficult to control

[Edited on 13-8-2023 by Fantasma4500]




~25 drops = 1mL @dH2O viscocity - STP
Truth is ever growing - but without context theres barely any such.

https://en.wikipedia.org/wiki/Solubility_table
http://www.trimen.pl/witek/calculators/stezenia.html
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[*] posted on 13-8-2023 at 09:53


Quote: Originally posted by Parakeet  
Here it is.
Haven't read everything yet, but the procedure looks amateur friendly.


Thanks but...
"Large-scale nitrite production (NaNO2) is achieved by bubbling gaseous N2O and NO through a solution of NaOH and Na2CO3".

Really?
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[*] posted on 13-8-2023 at 18:00


Quote: Originally posted by unionised  

Thanks but...
"Large-scale nitrite production (NaNO2) is achieved by bubbling gaseous N2O and NO through a solution of NaOH and Na2CO3".

Really?

Yeah. I've also read the third reference. It should be NO2.:(
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[*] posted on 14-8-2023 at 07:24


Quote: Originally posted by Sir_Gawain  
Would it work better with potassium nitrate? KNO2 is 10 times more soluble than KNO3.


In water it is, but in Alcohol the sodium nitrite salt is more soluble, and since alcohol is easier to remove thats why i went with the sodium salt instead of the potassium one.

I did try the reaction(didnt extract with alcohol yet)
When i mixed the sulfite and nitrate it was a bit endothermic, so it seem like it process very easily.

Also the sulfite/nitrate salt quantity on the video i posted are for the anhydrous sulfite salt, the user that made the video also made another video on how he made the sulfite salt and i think he is using the dihydrate like i did so the ratio is wrong. Probably why he has low yield(except from the fact he has lost a lot of product from overheating the mix and a broken beaker/kept the top part).

I think this method is very promising.

For 5G of NaNO3 its 9.187G of CaSO3.

$$Na2S2O5 + H2O \rightarrow NaHSO3$$ $$NaHSO3 + NaOH \rightarrow Na2SO3 + H2O$$ $$Na2SO3 + CaCl2 \rightarrow CaSO3 + NaCl$$ $$CaSO3 + NaNO3 \rightarrow CaSO4 + NaNO2$$ last one use heat - temp 230-300 C for 30-10 min.


[Edited on 14-8-2023 by fx-991ex]
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[*] posted on 14-8-2023 at 23:03


I ve never tried this way but i think its for bigger scale lab. Production the best methode. You dont need to melt something and you can make it in big Backers.
50 Parts Sodiumnitrate desolved in 150 Parts of Water. Add 350ml Ammonia solution with specific wight 0,96. After this give slowly 60 parts Zinc powder to it and hold temperature between 20C and 25C.After an short time all nitrate is converted to nitrite.
Its from an patent who improved an old lab methode from Poggendorfs Analen of Physic and Chemistrie.
Concentrations and Temperatures are improtand if it isnt exactly followed the nitrate gets reduced to hydroxide.
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[*] posted on 20-8-2023 at 04:54
Aluminium nitrite


aluminium metal doesnt react with nitric acid, and neither nitric acid vapors, but it will react with nitrate salts such as copper nitrate and iron nitrate- however if aluminium can be kept seperate from a nitric acid / nitrate salt mixture the NO2 may react rapidly with the aluminium to form aluminium nitrite
ideally the container is first flushed with CO2, butane maybe- as that is a very heavy gas

i did one attempt but the nitric acid/NOx fumes ate through the perforated plastic bag that was holding the shredded aluminium foil
(al foil flitter made by coffee grinding aluminium foil balls)

NO2 source was nitric acid and steel wool balls

i had a slight success proven by wettening the resulting mixture in IPAlcohol and adding dilute HCl which then formed IPNitrite, flame test faintly showed IPNitrite but it was largely yellow, probably due to hydrogen effervescence and presence of sodium nitrate as i used Na2CO3 + HNO3 to form the CO2 before i started the reaction, i shall attempt this further as this was a very poor attempt.
sadly it doesnt appear that aluminium carbonate exists, it decomposes into aluminium hydroxide which is difficult to filter as its a gel rather, IPNitrite might be best way to scavenge the nitrite from this reaction




~25 drops = 1mL @dH2O viscocity - STP
Truth is ever growing - but without context theres barely any such.

https://en.wikipedia.org/wiki/Solubility_table
http://www.trimen.pl/witek/calculators/stezenia.html
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[*] posted on 21-8-2023 at 04:09


Quote: Originally posted by Parakeet  
Here it is.
Haven't read everything yet, but the procedure looks amateur friendly.


Am I reading this correctly? The yield is about 50 micro moles per square cm per week.
So, with a 10cm by 10 cm electrode you would get less than a quarter of a gram of nitrite per week.



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[*] posted on 21-8-2023 at 04:11


Quote: Originally posted by Alkoholvergiftung  
I ve never tried this way but i think its for bigger scale lab. Production the best methode. You dont need to melt something and you can make it in big Backers.
50 Parts Sodiumnitrate desolved in 150 Parts of Water. Add 350ml Ammonia solution with specific wight 0,96. After this give slowly 60 parts Zinc powder to it and hold temperature between 20C and 25C.After an short time all nitrate is converted to nitrite.
Its from an patent who improved an old lab methode from Poggendorfs Analen of Physic and Chemistrie.
Concentrations and Temperatures are improtand if it isnt exactly followed the nitrate gets reduced to hydroxide.


Do you have a link to this patent or pdf?
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