Sciencemadness Discussion Board

Preparation of ionic nitrites

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Lionel Spanner - 18-3-2022 at 02:35

Quote: Originally posted by BAV Chem  
On another note, what is the calcium hydroxide even there for in this reaction?
Some include it while others leave it out completely. I had some success with starch and no hydroxide and got a 30% yield (see a few posts above).

It's so the reaction takes place in alkaline conditions, and produces an insoluble carbonate that can easily be separated from the nitrite. I should think you'd get very different reaction products in neutral or acidic conditions.

[Edited on 18-3-2022 by Lionel Spanner]

BAV Chem - 18-3-2022 at 09:13

It makes sense to run the reaction under alkaline conditions but doesn't that happen on its own without any hydroxide.
At the temperatures the reaction happens at some nitrite is inevitably gonna decompose to sodium oxide which ends up as sodium carbonate.
The calcium hydroxide will obviously deal with the carbonate but then you have NaOH instead of carbonate, which isn't much better. Also it'll turn into carbonate anyways upon standing.

Fantasma4500 - 18-3-2022 at 09:34

im gonna try the starch method asap. however it seems quite dark on my end as i have just tried using pure NaNO2, cold water and HCl with some IPA- the reaction quickly gets very hot and 2 layers seperate very quickly- quite contrary to anything ive achieved so far

i believe i also did flour, KNO3 and Ca(OH)2 - as well as without calcium hydroxide and got around same results, if not infact a bit worse
as for what im seeing, 30% yield seems incredibly optimistic. now i typically do use potassium nitrate, which is stable at much higher temperature (300*C higher?) and for now i can only pick between sodium sulfate contaminated NaNO3 or god forbid- sodium chlorate contaminated sodium nitrate. the latter can cause a spike in adrenaline. if sodium hydroxide and glucose is used, even in small amounts- didnt bother more with that yet.

as my results are not very interesting so far ill share what ive gone through

CaSO3 (well contaminated with Ca(OH)2 ) - seemed to do nothing at all
calcium sulfamate - total negative
steel wool- great potential
flour and hydroxide- great potential
lead- it should work but not having much success
ascorbic acid, reacted in solution - dont bother- very violent if mixed dry but does also work to some extent- maybe stoichiometry will make this work, or downing reaction speed with gypsum or similar.. sand?
MgAl- powder/chunks- some potential, mixture may with even very large pieces in closed atmosphere turn into a gigantic flare
sodium dithionate- total negative (solution)
iron oxalate- does work but not impressive yield

to be tried out: NaOH + NO2
NO2 from HNO3+ Cu, ascorbic can also work but will probably eat up a bunch of the formed NO2
it reacts accordingly: NaOH + NO2 = NaNO2 + NaNO3
oswald reactor leading into NaOH or even better KOH would work well, reference KNO2/KNO3 solubilities

as for crashing out KNO3, having a somewhat concentrated solution and added maybe 10-20% IPA (or ethanol) and then putting it in freezer should crash out mostly the nitrate

yet to try is Cr2O3 fused with KNO3, yielding K2CrO4 and KNO2

as for thermal decomposition- KNO3 is much more resistant to heat, i did in the past have some luck decomposing NaNO3 on hotplate (700*C hot)- ive tried this again with a new hotplate that only reaches 500*C and its just barely a slight positive for nitrite. ideally the sodium nitrite has to reach up to 800*C to efficiently decompose into sodium nitrite

purification of the nitrite can be done by converting it into IPN, distilling it off and reacting with a strong base, to create IPA and corresponding nitrite salt

cobalt and potassium combines with nitrite ions, or rather HNO2 to form a well insoluble complex called potassium cobolt hexanitrite. it does eat up some nitrite in the formation, and the complex appears very stable towards acids- so reaction with base is one of few ways ive found viable to make the nitrite useful in it- i didnt try concentrated sulfuric acid on it.

otherwise to have the crude portion precipitated out as KNO2/KNO3 in freezer with some alcohol added would remove most of the nitrate. KNO3 is 16g/100mL at 0*C in water, lower temperature aswell as alcoholic solution will further down the solubility while KNO2 is soluble at 200+g/100mL

clearly_not_atara - 18-3-2022 at 10:22

After some further consideration, it seems like the goal is a less energetic reducing agent. Perhaps a formate salt would be ideal:

2 KNO3 + 2 KHCO2 >> K2CO3 + 2 KNO2 + CO2 + H2O (g)

Or maybe even oxalate (if it reacts):

KNO3 + Na2C2O4 >> KNO2 + Na2CO3 + CO2

A sodium-potassium nitrate eutectic might be useful, but potassium formate already melts at 165 C with decomposition from 280-335 C:

[Edited on 18-3-2022 by clearly_not_atara]

Lionel Spanner - 18-3-2022 at 11:48

Quote: Originally posted by S.C. Wack  
Nitrite can be determined quantitatively by titration of a solution with permanganate in warm dilute sulfuric acid: 2KMnO4 + 5KNO2 + 3H2SO4 -> 5KNO3 + 3H2O + 2MnSO4 + K2SO4

[Edited on 17-3-2022 by S.C. Wack]

Another way is titration with ceric sulphate with ferroin as an indicator. This is advantageous if, like me, you don't have the patience for permanganate titrations.

(At my last job, I often carried out permanganate titrations to determine hydrogen peroxide content of peroxide cream, and it'd get painfully slow towards the end point; a single titration would typically take around 30-40 minutes.)

[Edited on 18-3-2022 by Lionel Spanner]

Keras - 21-3-2022 at 00:55

Quote: Originally posted by Keras  
Still, I’m surprised that the sodium dithionite method doesn’t work. If I follow the CRC Handbook, we have:

NO₃⁻ + H₂O + 2 e⁻ ⇋ NO₂⁻ + 2 OH⁻ e = 0.01 V
2 SO₃⁻ + 2 H₂O + 2 e⁻ ⇋ S₂O₄²⁻ + 4 OH⁻ e = -1.12 V

So logically, this reaction:

NO₃⁻ + 2 OH⁻ + S₂O₄²⁻ → NO₂⁻ + 2 SO₃⁻

should have e = 1.13 V, which is a pretty high K and should be almost quantitative, no?

Darn, that was totally wrong :(

Should be:

NO₃⁻ + S₂O₄²⁻ + 2 OH⁻ → NO₂⁻ + 2 SO₃²⁻ + H₂O

In any case, it doesn’t seem to work. I tried to heat the solution, and still no luck; so the problem is not kinetic. Why this does not work is beyond me.

clearly_not_atara - 21-3-2022 at 06:27

The ions electrostatically repel; there is no mechanism to bring them together. Kinetically, decomposition of dithionite is probably faster than rxn with nitrate under any realistic conditions. Maybe you could reduce an alkyl nitrate with dithionite, but that's a poor choice safety-wise.

Fantasma4500 - 21-3-2022 at 07:14

KNO3/NaNO3 mix have been used as a thermally stable salt mix for solar panels or something relating to solar panels
avoid KNO3.

ive had great success using lead, finally. but it requires sodium nitrate
you get about 200g Pb3O4 per 80g of nitrite, so you will need a bunch of lead if you wanna try it or utilize it
i havent looked much into turning Pb3O4 into Pb, which would be very handy..

procedure was as simply as just dump NaNO3 and Pb into a crucible, put on hotplate and cover with insulated pot, Al2O3 wool secured through a hole with a bolt+washer. my hotplate gets to 500*C which is plenty fine, you wanna stir the mixture up a few times as the reaction happens at the surface of the lead and Pb3O4 forms
it seems quite effective, best to do is to pour the molten content from crucible out into a metal tray, discard most of the lead for further use and then dissolve it in hot water, filter it and then turn it into IPN right away
then you suck out the IPN and store in bulk, distill it at 40*C and hydrolyze with NaOH to get- practically pure NaNO2

as it appears KNO3 + Pb doesnt work, i would steer clear of fractionally crystallizing nitrate/nitrite

if one was to construct oswald device and pump the NOx into KOH, one would get KNO3/KNO2 mix, which then can be fractionally crystallized, 60g is doable a day in industrial setup, maybe 30g with a really great homebuilt setup- loose number ive gotten from setups found on youtube with vague explanations of yield
NO would react with NaOH to form NaNO2- but forming NO doesnt as i see it seem viable, if NOx could be turned into NO that would make for very great yields with oswald

in any case, making nitrite it appears that you wanna avoid using potassium nitrate as starting material as its quite stable compared to sodium nitrate. i did also do 100g KNO3 + 50g Ca(OH)2 + 32g flour and it had some yields but not super impressive, i highly doubt it was as high as 30% yield, but it does appear to increase yields when you scale it up

clearly_not_atara - 21-3-2022 at 09:04

Nitrate salt eutectics are used as heat storage fluids with a wide liquid range and low corrosive potential in solar thermal installations to allow dispatchable electricity generation; could this be what you mean?

Nitrate salt decomposition is not exothermic in the absence of acids (incl. CO2) so as long as these salts are sealed they work great. In fact:
2 Na2O + 2 N2 + 5 O2 >> 4 NaNO3
is "exothermic" but kinetically impossible (for all we know).

Keras - 21-3-2022 at 11:57

Quote: Originally posted by clearly_not_atara  
The ions electrostatically repel; there is no mechanism to bring them together. Kinetically, decomposition of dithionite is probably faster than rxn with nitrate under any realistic conditions. Maybe you could reduce an alkyl nitrate with dithionite, but that's a poor choice safety-wise.

Of course, you’re right. I’m dumb.

By the way, I just tested sodium dithionite with ferrous chloride, and that gave me iron powder, which immediately oxidised when exposed to air.

[Edited on 21-3-2022 by Keras]

Keras - 22-3-2022 at 01:08

Is there any reason why zinc or iron could not reduced nitrates?

Since NO₃⁻ + H₂O + 2 e⁻ ⇋ NO₂⁻ + 2 OH⁻ e = 0.01 V, every metal that HCl oxidised should be able to reduce nitrates to nitrites, no?
Like: NO₃⁻ + Zn + H₂O → Zn(OH)₂ + NO₂⁻

Keras - 22-3-2022 at 11:47

Update: I did that. I just dropped a screw made of zinc plated steel into a solution of a few hundredths of milligrams of KNO₃.

Heated it up a bit, let it evaporate at r.t.

Got some beautiful transparent, needle like crystals. Does that correspond to potassium nitrate?

IMG_1011.JPG - 1.2MB

Fantasma4500 - 22-3-2022 at 15:01

@keras please read the other posts in this thread

interesting with dithionite and iron, this in solution could maybe work for nitrite- it could maybe work to produce zinc powder? i have a bit of it that stuff- honestly i feel done with experimenting with producing nitrites, i have found working methods by now

Fantasma4500 - 14-6-2022 at 03:21

NOCl + NaOH = NaNO2 + NaCl + H2O

HCl + KNO3 would form NOCl? this is much like NO2 + NaOH except- this one may be more selective as all the nitrogen compounds goes to NaNO2 where for instance NO2 + NaOH turns into i believe 1:1 nitrate/nitrite

i did also find other - seemingly automated saying:
NOCl + NaOH = NaNO3 + NO + NaCl + H2O
NaOH + NOCl = NaCl + HNO2

clearly_not_atara - 14-6-2022 at 06:07

I think it could work, but it's my understanding that the offgas from aqua regia is a lot more complicated than just NOCl. There's also some NO2, NO, and Cl2 in there.

Also, the variation I had heard most often was sending just nitric oxide -- NO without NO2 -- through NaOH, which apparently still gives nitrite (and nitrous? or something?). But I'm not sure how to generate this reliably in a controlled fashion -- the dilution of your HNO3 will change as it reacts, if you use metal reduction, so obtaining a steady stream of NO is non-obvious.

Fantasma4500 - 19-6-2022 at 03:43

i did a run where i used 100ish mL of HNO3 and maybe 15g copper, boiled down the solution a bit before reacting it with IPA + HCl, acquired about 25mL IPN - a slight bit tainted with IPA, it appears the IPA can both be in the water phase but also gets pulled out into the IPN
the IPN formed quickly pulls the IPA out of solution and seperates

i need to scale this up with proper dripping mechanism
i have 1 reaction flask with HNO3 being dripped into it, 1 airpump going in to force the gasses through (and eventually clean apparatus, and avoid suckback)
and then distillation bridge going to next flask to collect water, HNO3 and avoid suckback, then hose attachment comes out of this to bubble into NaOH

biggest challenge in this whole thing is keeping apparatus cold, getting the nitric acid and also having a working setup- dripping funnel, this can maybe be discarded by using lower conc HNO3 and dumping copper in like that

you get your second crop of NO2 by decomposing Cu(NO3)2 at 220*C, this gives same amount of NO2 as the nitric acid reaction does

clearly_not_atara - 19-6-2022 at 06:16

How are you differentiating isopropyl nitrite from isopropyl nitrate?

BAV Chem - 20-6-2022 at 02:54

Quote: Originally posted by Antiswat  
i did a run where i used 100ish mL of HNO3 and maybe 15g copper, boiled down the solution a bit before reacting it with IPA + HCl, acquired about 25mL IPN - a slight bit tainted with IPA, it appears the IPA can both be in the water phase but also gets pulled out into the IPN
the IPN formed quickly pulls the IPA out of solution and seperates

i need to scale this up with proper dripping mechanism
i have 1 reaction flask with HNO3 being dripped into it, 1 airpump going in to force the gasses through (and eventually clean apparatus, and avoid suckback)
and then distillation bridge going to next flask to collect water, HNO3 and avoid suckback, then hose attachment comes out of this to bubble into NaOH

Hey Antiswat, I can't quite follow this procedure here. I would think you added IPA and HCl to a solution of Cu(NO3)2 ...which produced isopropyl nitrite? Now that seems like it wouldn't work.
Or did you just react IPA, HCl, HNO3 and Cu directly to afford a whole slew of reactions which end up producing isopropyl nitrite amongst other things.

Either way this brings me to another idea which might just work to make nice and pure alkali nitrite.
I'm thinking if you react nitric acid with a large excess of EtOH you'll get a whole bunch of interesting reaction products: NOx gasses, HNO2, Acetaldehyde, Acetic acid, CO2... and eventually Ethyl nitrite. The latter is a gas and could easily be bubbled through an ethanolic solution of NaOH or KOH, affording the corresponding alkali nitrite which precipitates out.

I've already done some test with ethanol and nitric acid. If a large excess of EtOH is used the resulting gas is colorless and certainly smells like Ethyl nitrite. It also burns with the characteristic white flame of nitrite esters.
Quite possibly one could even substitute nitric acid for a nitrate salt and a mineral acid.

What do you think about this approach?

[Edited on 20-6-2022 by BAV Chem]

Fantasma4500 - 20-6-2022 at 06:53

the procedure was reacting Cu with HNO3 to form Cu(NO3)2 and NOx, was was pumped into NaOH to yield NaNO3/NaNO2- this solution was reduced and then reacted with IPA/HCl to form IPN

i have thought of directly pumping it into IPN but it might give other compounds as well that you then have to deal with seperating, volatile ones, potentially explosive.

HNO3 and EtOH? interesting. ascorbic acid also works despite the ascorbic acid reacts with the NOx

this would explain that NOx + hydroxide is indeed a quite viable method. burning dry ammonia is also an idea, but thats a more complicated apparatus, imo. oh i might also stress... ethyl nitrate may derive from HNO3+ EtOH. methyl nitrate is a high-explosive with decent sensitivity. thiscould be a dangerous reaction.

upon reading your writings on Etnitrite- thats very interesting... Etnitrite is like 13*C boiling point or something iirc, i would try this with IPA where the resulting nitrite is boiling at about- hm 40*C? this is much easier to handle, less volatile, you will need to react this chemical with NaOH to once again produce the desired nitrite salt.

NaHSO4 + nitrate could do yeah, though i think activating the NaHSO4 takes some heat- you might have to distill on it seperately first, thus arriving at HNO3
cant platinum or something catalyse the decomposition of nitric acid? there has got to be some more simple way to this

i would indeed say the nitrite flame is quite likely .... buuut that could also be the same with nitrate? so acquire an amount of this material and test boiling point maybe, i usually just put digital thermometer into a testtube and carefully heat it

maybe we should look further into the decomposition of HNO3? as i mentioned, Cu(NO3)2 - which can be made by Ca(NO3)2 + CuSO4 decomposes at 220*C to yield decent amounts of NO2 as it turns into CuO. if some NO forms thats actually much better than NO2, maybe we wanna also look into turning NO2 into NO? my last searches on this yielded not much

Hexabromobenzene - 22-6-2022 at 17:40

Sodium nitrite is a byproduct of the production of nitric acid from the air. Absorbing gases after water by sodium carbonate solution you will have pure nitrite with almost no nitrate

To get pure nitrite, you need an equivalent mixture of NO and NO2.

If you do not want to get nitric acid from the air, then you can get nitrite directly from the air. Reduce the maximum volume of the oxidation chamber for nitrous gases

You will need a transformer from neon lamps for 7-15 kilovolts.
You can easily have a nitric acid output of up to 10 grams per kilowatt hour or more.
Nitric acid and sodium nitrite can be free if you find a wire with 220 volts without an electric meter))

[Edited on 23-6-2022 by Hexabromobenzene]

clearly_not_atara - 23-6-2022 at 06:05

So what you're saying is, if you feed NOx through neutral water first, and the unabsorbed gases through alkali, the second flask will contain almost exclusively nitrite?

It makes sense, since the secondary flow should be mostly NO with little NO2. It may be relevant in this context that iron (III) nitrate releases practically all of its NOx at just over 150 C:

[Edited on 23-6-2022 by clearly_not_atara]

Fantasma4500 - 5-7-2022 at 00:52

it appears that youre right, about 150*C its effectively turning into NOx

im not sure if im reading this right though, it seems quite vague, they dont mention really what the weight is thats being lost- nothing about nitrous gasses? as for copper vs iron i would further argue that the decomposed copper nitrate can be recycled with nitric acid, while Fe2O3 might be more difficult to do like this, obviously iron nitrate has a lot of waterweight to it and they mention as much as 1300*C
i think i did come around the actual relevant decomposition temperature of iron nitrate last time i looked into the nitrate salt decompositions for NOx

the NO2 scrubbing with water is very interesting, but simply dump it into KOH solution, chill this well and you will have quite concentrated KNO2 solution- and then you can very easily turn this into nitrite of whatever sort, maybe ethyl nitrite directly used for isomeric re-jumbling into nitroalkyl, whatever the correct term be.

ANOTHER trick i was hinted by a fellow chemist in this would be to add some H2SO4 to the copper before adding nitric acid
this will react with the Cu(NO3)2 formed- making more HNO3, which then goes on to react once again

this would be very interesting to try with iron. drip HNO3 onto that while its already got H2SO4, the iron never becomes iron nitrate

the NO2 water scrubbing method seems rather silly to me because largely NO2 is produced- however if you recycle the freshly formed HNO3- albeit very dilute it can start to make sense, but you'd be running some serious bulk to make it work, KOH would be best bet at seperating out large amounts very fast.
forming NO2/NO out of air and pumping into KOH would be a better method than using precious HNO3, it may produce about 30 grammes a day with a small setup, that being maybe 15g of nitrite? in a month thats 450g.

we still have yet to test out how NOCl + NaOH does. if the nitrogen in that works entirely like NO would in NaOH that would mean very great yields, and one could just fractionally crystallize the sodium nitrite and HCl / nitrate salt may be possible for generating the gas

Fantasma4500 - 9-7-2022 at 02:02

im thinking of utilizing ethyl nitrite, as this can be isomerized into nitro by passing it through a tube thats heated to 130*C packed with mineral wool

if the temperature is kept at near BP of EtOH- ethyl nitrate should decompose? its a potentially dangerous procedure, i was thinking of just dropping nitric acid into EtOH/copper metal and using ethyl nitrite directly
but ethyl nitrate is quite dangerous if it doesnt decompose.

here they explain 10% HNO3 in EtOH is potentially explosive... because of vapors? and ethyl nitrate is commonly made by nitration of ethanol - but methyl nitrate is made by distilling HNO3 with MeOH
we need something on how to completely eliminate EtNO3 because just a single drop of this that goes off could wreck glassware and even your life
"The invention discloses a method for preparing ethyl nitrate by continuously nitrifying nitric acid steam, which belongs to the field of energetic materials ..."
they mention adding urea and trimeric cyanamide, its a bit confusing to read since its translated from chinese. they mention adding the ethanol directly to this nitric acid mixture which is.. 75% or 68%? at a temperature of 60*C
distillation at 90*C where the ethyl nitrate is left in the reaction vessel, which it can then be washed with simply water

"be warming up to 90 ℃ and start to reflux 5 minutes; Reflux, just steam product on one side, 90 ℃ of crude product upper stratas that obtain of distillation temperature are unreacted ethanol, and lower floor is thick ester and acid,"

this would be safer if you have excess EtOH and heat that to BP of ethanol?

99% EtOH reacts with 70% HNO3 within a minute at STP while 95% can take an hour of incubation time, where it in both cases vigorously boils up- likely ethyl nitrite. hydroxylamine is also mentioned. possibly ethyl nitrate carefully decomposing into nitrogenous gasses
now, nitroethane however, is also somewhat insoluble in water- i believe its 4g/100mL, but it floats ontop of water, while ethyl nitrate- assuming the two wouldnt end up dissolving in one another, would sink to the bottom
this would potentially make this reaction feasible
"When ethyl alcohol reacts with nitric acid, it forms: · Nitromethane · Nitroethane · Ethyl nitrate · Diethyl ether · Chemical Properties of Alcohols and Phenols - ...
Top answer:
Correct option is C) Ethyl nitrate"
no mentions on reaction conditions.

keeping the HNO3/EtOH mixture from suddenly increasing in temperature would be how to keep ethyl nitrate away, i believe
and this could maybe be done by reacting the HNO3 directly with Cu to form NO/NO2, which with methanol can in vaporphase directly form nitromethane
and this would also ensure no runoff if the nitric acid can react right away with the copper.. in excess EtOH maybe? unfortunately EtOH is miscible with practically all solvents, but the ethyl nitrite should come off if its just kept at 30*C, or room temperature even

[Edited on 9-7-2022 by Antiswat]

Pumukli - 13-7-2022 at 06:11

I would avoid ethyl-nitrite/nitrate altogether and would try to do something (optimize?) the HNO3/starch method.
E. Divers in 1899 (or so) wrote that this is the way to go if you want to make good quality NaNO2 with very low NaNO3 contamination.
The trick is he "forgot" to mention the right concentration of HNO3 and the right temperature of the reaction! :-) Only wrote about those that "you should adjust them in order to produce NO/NO2 mix where NO is in a slight excess". This is the key. This is what the optimization should aim for!
He only wrote that the gas mix should be absorbed into NaOH (or Na2CO3) solution, AND the excess, unabsorbed gas should be only NO! So it should be COLOURLESS what "bubbles away" from the absorbent solution!

Basically one should make a NOX generator which makes slightly more NO than NO2. I would rather "play" with this generator and determine the optimal working parameters of it than risk to work with EtONO/EtONO2. But for each of his own. ;-)

[Edited on 13-7-2022 by Pumukli]

Texium - 13-7-2022 at 09:02

It looks like this thread has been the main place for nitrite synthesis-related discussion for the last five years, so I've merged it with the original nitrite thread and stickied the resulting big thread for better visibility.

As it seems like sodium nitrite is becoming less available OTC in some areas, having a sticky thread on unconventional preparations seems warranted.

AJKOER - 16-7-2022 at 08:10

Reagent Friendly Path: Just a small starting amount of H2O2, aqueous NH3, Cu, air and some NaOH to maintain alkaline pH, and as this is an electrochemical cell, a small touch of a good stating electrolyte like sea salt, all resulting in NaNO2.

It is well known that the electrochemical dissolving of Cu with aqueous NH3 in the presence of air creates a side product of NH4NO2, which apparently, at one point (I suspect removal of ammonia resulting in a lower pH), creates a massive sudden N2 release per the decomposition of formed NH4NO2 (so, do employ a wide mouth and tall reaction vessel and still, at times, I get a spillage event).

See my prior comments on this event and details on the reaction previously posted here which also includes another reference on the electrolysis route to nitrite (not a galvanic cell).

Now, adding NaOH likely averts this NH4NO2 decomposition reaction forming the targeted NaNO2.

Also, the reaction with ammonia is depicted as proceeding both with H2O2 and O2 exposure along I suspect, it is not O2, but perhaps superoxide formed from solvated electrons (e- per the electrochemically cell) acting on oxygen (O2 + e- -> .O2-). Basis, quoting from a source at :

"Superoxide (O−2) also reacts with trace metals (Figure 1B) to produce H2O2; reactions with Cu(I) can proceed rapidly with rate constants (k) on the order of 2 × 109 M−1 s−1 (Zafiriou et al., 1998), although [Cu(I)] is relatively low (~0.1 nM; Moffett and Zika, 1988)."

which is my polite way of saying that a direct oxygen path may not be precisely correct.

[Edited on 16-7-2022 by AJKOER]

Fantasma4500 - 18-7-2022 at 04:02

you mention formation of nitrite in the same reaction vessel as.. superoxide, and also using H2O2? nitrite reacts with oxygen to form nitrate
if this is a membrane cell it adds difficulty to the operation, though its very possible to make your own membrances with PVC glue and plastic cloth
would be cool if it was a viable electrochemical route for making nitrite

i would go with oswald if i had space to have such a thing going, simply dumping the gasses through water and then have them pumped through hydroxide solution should do

the starch method is interesting, any mentions of what may be formed? nitrostarch is a thing..

im not quite sure of how one would gauge the quality of the gas mixture, amount of gas bubbles in second container after first water-scrub? im not even sure if NOx from nitric acid is viable for this.
i employed my IPN method, i got 100mL of presumably mostly IPN
so i set up for short path distillation and ended up with a majestic.... 35 milliliters of IPN, from maybe 300mL 62% HNO3 - hmmmm, im not sure if this is a good yield.
i did insert an airbubbler into my flask to counteract the ups and downs in temperature from my hotplate and the reaction- afaik nitrite isnt super sensitive to oxygen, but maybe im wrong, maybe i turned my poor nitrite into nitrate like that. maybe i messed up on making the IPN itself?

regardless if one goes with the NOx method, its ideal to pump the gasses into KOH solution, then dumping it in freezer will crystallize the KNO3 and youre left with mainly KNO2 - K2CO3 may also be used

how would one go about modifying the starch reaction, is the nitric acid dripped into solid starch, into starch solution, starch solution dripped in? controlled heating?

AJKOER - 23-7-2022 at 06:43

Correctly, there are claims of the direct action of oxygen gas on nitrite (and not superoxide) with NaNO2, as an example, a decomposition warning for the dry salt (see ), to quote:

"away from AIR, LIGHT, and MOISTURE."

However, I suspect, this trilogy are hardly independent agents, as nitrites are noted as photocatalysts.

NO2- + Light (Blue) -> .NO2 + e-

And, the hydrolysis of .NO2 is a known path to nitrous and nitric acid:

.NO2 + .NO2 + H2O -> HNO2 + HNO3

Again the action of light on nitrous acid:

HONO + Light -> .HO + .NO

And, the action of a solvated electron on oxygen:

O2 + e-(aq) --> .O2- (the introduction of the superoxide radical anion in the presence of an oxygen source)

which can also be sourced in the present of a transition metal impurity, in the so-called metal auto-oxidation reaction with dissolved oxygen:

M -> M+ + e-
O2 (aq)+ e- --> .O2- (aq) (again with the formation of superoxide)

And, the action of superoxide on nitric oxide is a well known path to peroxynitrite and eventually nitrate:

.NO + .O2- -> ONOO

So, once the photo or a transition metal inducing electron presence ceases (not just air presence, for example), a nitrite product could be stable. This is supported here in the copper, ammonia, O2/H2O2 reaction system with no citations of any ammonium nitrate creation.

Also, there is a source citing the presence of ammonium carbonate (see ) fostering the dissolving of copper ore with water, ammonia and oxygen. This is interesting as CO2/HCO3- appears to be a radical reaction promoter, to quote a source ("Radical production by hydrogen peroxide/bicarbonate and copper uptake in mammalian cells: Modulation by Cu(II) complexes" at ):

"It is well-known that the bicarbonate/carbon dioxide pair, the presence of which is important in maintaining physiological pH in extracellular body fluids, can accelerate the transition metal ion-catalysed oxidation of various biotargets. Despite of its relevance, however, most of the mechanisms that have been proposed to account for this important effect remain controversial [8], [9], [10], [11], [12], [13], [14], [15], [16], [17], [18], [19], [20], [21]. On the other hand, it is accepted that the bicarbonate/carbon dioxide pair can increase peroxynitrite-mediated one-electron oxidation and nitration via formation of the carbonate radical and nitrogen dioxide [22], [23].

So discounting possible underlying radical based reaction paths (aka, not just simply a direct elemental oxygen interaction) is not likely precisely correct, in my opinion.

As further evidence of the actual possible complexity in systems involving dissolved (not gaseous) .NO see my sources cited relatedly in this SM thread where dissolved .NO apparently behaves differently from gaseous .NO interacting with oxygen with a near total absence of eventual nitrate creation.

[Edited on 23-7-2022 by AJKOER]

[Edited on 23-7-2022 by AJKOER]

Lionel Spanner - 24-7-2022 at 07:41

Your man Experimental Chemistry has a nice video on preparing potassium nitrite via calcium formate.

Fantasma4500 - 29-7-2022 at 07:38

wow. 68% yield of KNO2 starting from Calcium Formate and KNO3, heated at about 300*C, no explosively exothermic reaction

sodium formate may be acquired in 25kg bags as a special de-icer, its used especially for areas that gets very cold but also to avoid salt getting onto vehicles, aircrafts etc.

sodium nitrate should also be very much doable

2KNO3 + (HCOO)2Ca = 2KNO2 + CaCO3 + H2O + CO2

this seems to be exactly what we have been hoping for to pop up in regards to reduction of nitrate salt. could formic acid maybe reduce nitric acid to nitrous acid?
The Reaction of nitric acid with formaldehyde and with formic acid and its application to the removal of nitric acid from mixtures

"Nitrites do not react with formaldehyde in neutral solution, " HNO2 may infact be formed in this reaction, it seems.

i see a comment in the calcium formate method video
"nitrite into HCL gives largely nitrosyl chloride"
would this not imply that it can maybe go the other way around again, so NOCl + NaOH = NaNO2 + NaCl?
HCl + KNO3 = NOCl (basically)

anyhow back to calcium formate method:

calcium formate is about 16g/100mL (roughly same 0-100*C)
sodium formate is 49-160g/100mL
CaCl2 is 60 to 160g/100mL

CaCl2 + NaForm = CaForm + NaCl
fractional crystallization would work- not that NaCl would really be a big issue as impurity by my assumptions

could we maybe directly use sodium formate for this reduction instead?

thank you very much for this input @Lionel Spanner

Fantasma4500 - 31-7-2022 at 09:17

Ca(NO3)2 + 2NaForm = CaForm + 2NaNO3

164g + 70x 2 140g = BOIL THIS DRY, combust

before / after total weight of dry substance
304g = 195g
i react it at about 360*C, it takes maybe 10 minutes before it starts to react, it doesnt evolve a lot of smoke but it does smell like a nitrate pyrotechnic composition so thats very annoying if you react too much at once.

(COOH)2Ca + 2 NaNO3 = 2 NaNO2 + CaCO3 + CO2 + H2O
130g + 170g = 138g + 100g

dissolve in water, boil NaNO2 dry or use as solution

i have tested a smaller sample of 10 grammes with IPA and HCl
when the IPN is formed it makes the polarity of the IPA seperate out so you get a very clear indication

thereafter i ignited the gasses in the flask and typical nitrite flame was seen. its also vasodilating.
its ideal to dissolve the soluble contents in water before adding acid as it causes a lot of effervescence, namely HCl

it may be possible to take a concentrated NaNO2 solution and dump into EtOH to precipitate out the NaNO2 for easy isolation as NaNO2 is 4.4g/100mL solubility in EtOH

this procedure can be done inside if done in small quantities, 100 grammes was too much
my heating device is a single hotplate, a stainless steel pot ontop of that which is isolated with Al2O3 ceramic wool, secured by a bolt + washer + nut going through top of the pot

i shall attempt further to simply react Ca(NO3)2 and NaFormate by direct decomposition so i dont have to dissolve that in water, and then boil that dry
i may add some Al2O3 because the fertilizer i get my Ca(NO3)2 is a mix of ammonium salt and Ca(NO3)2 roughly 90-10% ammonium salt being the latter, which can decompose... very rapidly on an unfortunate day- i might just do this reaction with homemade Ca(NO3)2 just to be completely sure, energetics can be very powerful especially if situated on a hotplate covered by a pot

i have no guesstimates for yields yet but this appears to be how were gonna be making nitrites in the future
only things i have to add is that calcium formate decomposes thermally into calcium oxalate- so calcium oxalate should maybe be attempted with sodium nitrate?

Myc - 24-9-2022 at 20:08

I tried a small batch of the above method (sodium formate and calcium nitrate) too. The reaction in the crucible was rather enthusiastic and measured over 320C. While some of the nitrite salt survived this I suspect there was also significant decomposition at this temperature (this supposedly happens above about 300C). So I guess the product has sodium oxide/hydroxide in there too?

Σldritch - 25-9-2022 at 04:38

Hm, is oxalic acid available dirt cheap as some product? I've haven't had much luck, but, I've considered making sulfuric acid from Ferrous Sulfate (which is dirt cheap itself) with it. The by-product would be Ferrous Oxalate Dihydrate. Maybe worth trying the reaction with that as well if one could end up with a lot of it as a byproduct...

EDIT: Huh, apparently Calcium Oxalate is almost three orders of magnitude less soluble than Calcium Sulfate.

[Edited on 25-9-2022 by Σldritch]

B(a)P - 25-9-2022 at 17:32

Quote: Originally posted by Σldritch  
Hm, is oxalic acid available dirt cheap as some product?

In Australia it is sold in hardware stores as rust and stain cleaner and costs about $15/kg.

clearly_not_atara - 26-9-2022 at 07:14

I was able to order oxalic acid per se, no questions asked, for the cement experiment I still haven't finished. It's sitting in a storage locker 1000 miles away right now for reasons. It's very useful for cleaning stuff; I think it was labeled as mold prevention for decks or something.

But I wouldn't use it for this because calcium oxalate is pretty inert and won't dissolve in anything. Plus, oxalic acid is strong enough (100 times stronger than formic acid) to protonate nitrate to some extent and may evolve some NO2.

Myc - 27-10-2022 at 16:24

Hi team,
It seems to me that we have a few different methods in this thread that yield a mixture of sodium nitrate and nitrite. However, there has been little discussion of separating the two. They seem to have similar solubilities in water. Crystallization by cooling a saturated aqueous solution would thus work to concentrate the salt that's present in a significantly larger quantity.

Does anyone have ideas for separating a mixture that has a more even ratio of nitrate/nitrite?

Sir_Gawain - 27-10-2022 at 20:37

I've thought of a way to make nearly pure NaNO2. First, you produce crude NaNO2 using one of the reduction methods (like the sulfur/NaOH/NaNO3 reaction). Then dissolve the impure product in water and drip it into a concentrated acid. The gasses (NO and NO2) are led into a cold sodium hydroxide solution. The reactions are:
2NaNO2 + H2SO4 = Na2HSO4 + NO + NO2 + H2O
Then; NO + NO2 + 2NaOH = 2NaNO2 + H2O
When the sodium hydroxide (now nitrite) solution pH reaches neutral, stop the gas production and evaporate the water to recover your pure product.

Myc - 28-10-2022 at 02:02

I believe I've read that this also produces nitrate...

Sir_Gawain - 28-10-2022 at 07:28

Bubbling just NO2 into sodium hydroxide produces both NaNO2 and NaNO3. (2NO2 + 2NaOH = NaNO2 + NaNO3 + H2O)
Using a mixture of NO2 and NO results in only NaNO2.

clearly_not_atara - 28-10-2022 at 20:06

Adding NiCl2 to a solution containing KNO2 and KNO3 should precipitate K4Ni(NO2)6*H2O selectively as a brown powder. It may be possible to return this to KNO2 by reaction with KOH.

Fantasma4500 - 17-11-2022 at 08:24

i just have some more input for this.
aluminium may be utilized for this reaction- very dilute HNO3 reacts with Al, somehow bypasses passivation layer- this could maybe imply NO formation over NO2? think of set&forget kinda reaction.

i got around this as i remembered Al reacts with NO2 to form- what? i didnt get to that, possibly it forms Al(NO2)2 but the Al2O3 which forms from the HNO3 might bump that into Al(NO3)2 - hm. more concentrated acid may be used with a small amount of HCl being added in- or H2SO4 maybe? HAc? Phosphoric?

this is a potential low cost method, otherwise birkeland eyde would be better, and yet better would be formate with nitrate

why would a MIXTURE of NO2 and NO produce pure NaNO2?
NaOH + NO = NaNO2
NaOH + NO2 = NaNO2+NaNO3
the NO doesnt act reducing

Sir_Gawain - 17-11-2022 at 16:57

NaOH + NO does not make NaNO2 (unless by NaOH + NO = NaNO2 + H). And in this case, NO almost does act as a reducing agent. My guess is that NO2 and NaOH react to form nitrate and nitrite, then NaNO3 + 2NaOH + 2NO = 3NaNO2 + H2O.

clearly_not_atara - 17-11-2022 at 21:17

You could see a reaction like

4 NO + 2 OH- >> N2O + 2 NO2- + H2O

As to whether this actually occurs, I don't know, but Wikipedia seems to think it does. One possible reaction pathway is:

NO* + OH- >> -ONOH

-ONOH + NO >> NO2- + HNO

2 HNO >> H2O + N2O (known reaction)

Lionel Spanner - 21-11-2022 at 13:39

Anyone else had grief trying to boil sodium nitrite solutions to dryness?
I've been able to concentrate it, and get a viscous, highly concentrated solution full of solids with a boiling point around 158 °C - the very last stage, which involves getting the last bit of water out by drying it an oven at 200 °C is where it's gone tits up, every single time.

1st attempt - used a porcelain dish; the salt crust was so hard I ended up cracking the dish and contaminating the product with bits of porcelain.
2nd attempt - used a dish lined with silicone coated oven-safe aluminium foil; after the concentrated solution was added, the coating lasted a whole 5 seconds before both the product and the foil committed suicide.
3rd attempt - used a small steel bowl; product reacted with the iron in the steel, causing it to evolve nitrogen and foam profusely, spilling much of it onto the base of the oven, and the final product was heavily contaminated with iron oxides.

Maybe something like an oven-safe silicone mould is the answer....

[Edited on 21-11-2022 by Lionel Spanner]

[Edited on 21-11-2022 by Lionel Spanner]

Myc - 21-11-2022 at 19:00

If you're boiling it down to concentrate it, stop at around 130 degrees then cool it to fridge temperature. The sodium nitrite will crystallise and you'll have a fairly small amount of water. Vacuum filter this paste while still cold. The filtrate won't have too much product in it, but you could boil and concentrate again if you have a fair bit. The solid can then be dried under a fan if you have low humidity. Finish the drying process in a desiccator (sealed box with plenty of dry CaCl).

Lionel Spanner - 22-11-2022 at 16:40

Quote: Originally posted by Myc  
If you're boiling it down to concentrate it, stop at around 130 degrees then cool it to fridge temperature. The sodium nitrite will crystallise and you'll have a fairly small amount of water. Vacuum filter this paste while still cold. The filtrate won't have too much product in it, but you could boil and concentrate again if you have a fair bit. The solid can then be dried under a fan if you have low humidity. Finish the drying process in a desiccator (sealed box with plenty of dry CaCl).

Thank you! Since posting my mini-rant I've had time to consider this problem in a calmer manner, and this proposition has a lot in common with the course of action I'm considering.

Solutions of sodium nitrite become supersaturated on boiling when their boiling point reaches around 126-130 °C, so today I made a fresh batch and transferred the concentrated distillate to another container once the boiling point got to around 130 °C, and cooled it in the fridge. Net result: a nicely mobile salt crust in a little bit of pale yellow water.

However, instead of vacuum filtration and dessication, I'm intending to remove the remaining water by azeotropic distillation with xylene (the azeotrope boiling at 92 °C, and comprised of 40% water/60% xylene at atmospheric pressure.) If this works, then the freshly precipitated salt could nucleate upon the surface of the existing solids, and the final product, being completely insoluble in xylene, could be separated by filtration and dried by heat. Fingers crossed!

[Edited on 23-11-2022 by Lionel Spanner]

clearly_not_atara - 3-12-2022 at 17:43

I would cross my fingers too when heating a flammable solvent with a somewhat unstable oxidant...

Lionel Spanner - 3-1-2023 at 16:44

Update: azeotropic distillation with xylene worked up to a point, removing all but about 5% of the water, but did not break up the concentrate as hoped, most likely due to the differences in density and polarity between the two.

The distillation also became less and less effective as the water content decreased. I managed to remove nearly all of the wet nitrite from the flask while it was liquid, and place it in a desiccator (read: tightly sealed Tupperware-style box) lined with solid caustic soda.

After 6 weeks in storage, no further weight loss was observed, and the final result was 22.7 g sodium nitrite of unknown purity, having started from 42.5 g sodium nitrate, representing a maximum yield of 66%.

Since dehydrating the product is such a chore, I will attempt to recover it by recrystallisation on the next attempt, most likely from alcohol with some added water.

[Edited on 4-1-2023 by Lionel Spanner]

[Edited on 4-1-2023 by Lionel Spanner]

Fantasma4500 - 12-1-2023 at 04:09

@Lionel Spanner
NaNO2 EtOH 4.4g/100mL, i bet acetone is much worse at dissolving it
precipitate it out by adding in a solvent, then vacuum filter
to speed things up you may blow hot air onto the mouth of the vacuum filter to remove the acetone faster
ideally you would place a beaker with the damp NaNO2 crystals in an oven, sealed up so air doesnt get into it- maybe hotplate. oven set above the boiling point of the solvent used
air can turn NaNO2 into NaNO3
i must also stress that using a solvent to flush with causes the water to be removed, and heating will now not so much cause the material to dissolve and gradually form a hard lump, this is especially important if you desire to dump it into maybe an erlenmeyer flask, youre skipping the part where you have to crush up hard lumps, and with luck you get fine powder right away.

anyhow, NO formation can be done with ammonia and oxygen, air pump into NH4OH solution, then this pumped through Cr2O3, maybe a glass tube of kitty litter silica gel with Cr2O3- then this lead into NaOH

ive flushed chlorates with acetone many times, a method for making pure NaClO4 is to use acetone as solvent
that solvent has to leave the NaClO4 eventually somehow. NaNO2 is barely an oxidizer, its so weak oxidizer its also a reducing agent

Lionel Spanner - 14-1-2023 at 14:56

Quote: Originally posted by Antiswat  
@Lionel Spanner
NaNO2 EtOH 4.4g/100mL, i bet acetone is much worse at dissolving it
precipitate it out by adding in a solvent, then vacuum filter
to speed things up you may blow hot air onto the mouth of the vacuum filter to remove the acetone faster
ideally you would place a beaker with the damp NaNO2 crystals in an oven, sealed up so air doesnt get into it- maybe hotplate. oven set above the boiling point of the solvent used
air can turn NaNO2 into NaNO3
i must also stress that using a solvent to flush with causes the water to be removed, and heating will now not so much cause the material to dissolve and gradually form a hard lump, this is especially important if you desire to dump it into maybe an erlenmeyer flask, youre skipping the part where you have to crush up hard lumps, and with luck you get fine powder right away.

I will definitely avoid unnecessary exposure of nitrite to air at high temperatures. It might well be worth removing the bulk of the water by vacuum distillation, in lieu of a rotary evaporator.
The solubility figure for alcohol is at 25 °C, and it's almost certainly higher at 70-80 °C; acetone could work very well as an anti-solvent, in order to get as much of it out of solution as possible. The catch is that any unreacted nitrate has a very similar solubility profile to nitrite.
Apparently, nitrite is a little less soluble in water than nitrate at high temperatures (160 vs 180 grams per 100 g water), so if push comes to shove, the two could be separated by fractional crystallisation, though I'm hoping it won't come to that.

Coming from a UK state school that spent its entire budget on languages, I never really understood the process of recrystallisation, and so I did a terrible job of it when I was at university. Now I can do it at my leisure, I now understand the process much better, and it just seems so much more simple.

Lionel Spanner - 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.

Myc - 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.

Lionel Spanner - 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]

Lionel Spanner - 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

A major breakthrough!

Lionel Spanner - 11-3-2023 at 09:29

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!

clearly_not_atara - 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!

Lionel Spanner - 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.

Lionel Spanner - 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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