Sciencemadness Discussion Board - Preparation of ionic nitrites

Preparation of ionic nitrites

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

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:

https://pubs.acs.org/doi/pdf/10.1021/es60044a002

[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

@clearly_not_atara
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?

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

https://pubs.acs.org/doi/abs/10.1021/jp972143m

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.

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:

https://link.springer.com/content/pdf/10.1023/A:101011281401...

[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

https://tkuir.lib.tku.edu.tw/dspace/retrieve/82495/-JMR-JMR5...

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.
https://en.wikipedia.org/wiki/Ethyl_nitrate

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

https://en.wikipedia.org/wiki/Nital#:~:text=A%20solution%20o...

https://patents.google.com/patent/CN102531909B/en
"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

https://www.toppr.com/ask/question/when-ethyl-alcohol-reacts...
"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 https://www.sciencemadness.org/whisper/viewthread.php?tid=18... 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 https://www.frontiersin.org/articles/10.3389/fmars.2016.0023... :

"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

AJKROER
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 https://nj.gov/health/eoh/rtkweb/documents/fs/2258.pdf ), 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 https://www.sciencemadness.org/whisper/viewthread.php?tid=14... ) 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 https://www.sciencedirect.com/science/article/pii/S016201341... ):

"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 https://www.sciencemadness.org/whisper/viewthread.php?tid=15... 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.
https://www.youtube.com/watch?v=BhBxDkU7AmI

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
https://onlinelibrary.wiley.com/doi/abs/10.1002/jctb.5010080...

"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.
https://gyazo.com/c414de838c4d0d7ac40e7399deb23d6f

(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
https://gyazo.com/95df06a3975a0ba352fecd3a3db5bebb

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

@sir_gawain
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
https://edu.rsc.org/exhibition-chemistry/chromiumiii-oxide-c...

@clearly_not_atara
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?

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?

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.

fx-991ex - 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]

Lionel Spanner - 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!

fx-991ex - 20-6-2023 at 08:18

Quote: Originally posted by Lionel Spanner

Quote: Originally posted 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

Lionel Spanner - 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!

Sir_Gawain - 22-6-2023 at 09:52

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

Lionel Spanner - 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]

Lionel Spanner - 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!

Parakeet - 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 378 times

Lionel Spanner - 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!

unionised - 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?

Parakeet - 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.

fx-991ex - 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]

Alkoholvergiftung - 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.

Aluminium nitrite

Fantasma4500 - 20-8-2023 at 04:54

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

unionised - 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.

Kloberth - 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?

Alkoholvergiftung - 21-8-2023 at 07:32

See Attachment. They are in German.

Attachment: NitritDE000000168450A_2.pdf (65kB)
This file has been downloaded 383 times

Alkoholvergiftung - 21-8-2023 at 07:33

First Part.

Attachment: NitritDE000000168450A_1 (1).pdf (78kB)
This file has been downloaded 392 times

Lionel Spanner - 22-8-2023 at 06:11

Quote: Originally posted by unionised

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.

That's with 0.1 molar ammonia, i.e. 0.17% w/v. I suspect a more concentrated solution would result in better yields.

[Edited on 22-8-2023 by Lionel Spanner]

Alkoholvergiftung - 24-8-2023 at 04:50

I ve found this crusicible methode.Englisch and later german copy.

Very finely divided metallic copper is first prepared by distillation of acetate of copper, and from this freshly prepared metallic powder, 2 equivalents, or even a slight excess, are taken to 1 equivalent of nitrate, according to the equation
2 Cu + KO, NO⁵ = 2 CuO + KO, NO³.
Persoz took 320 grm. saltpetre and 200 grm. from the copper represented in the manner indicated.
To prepare an intimate mixture, the saltpeter is first dissolved in the smallest possible quantity of hot water, and then the copper is added, which at first is difficult to wet. When the mixture has become quite uniform, it is heated in a porcelain dish, or better yet, in a cast-iron pan, in a sand bath, stirring constantly to prevent splashing. When the mass is completely dry, there comes a moment when, like a pyrophor, it catches fire and glows; when the combustion is over, which lasts but a moment, the reaction has taken place; it is allowed to cool, the melt treated with water, filtered, and the nitrate allowed to crystallize. If excess copper has been used, no nitrate is present, and crystallized nitrite is obtained at once, which is then melted and kept in well-closed flasks, as the salt is very hygroscopic. Any undecomposed nitrate of potash present is separated out by the first crystallization, since it is far less soluble than nitrate of potash. The copper oxide obtained as a residue during the operation can, after proper washing, be used for organic analysis, or at least for mixing with the organic substance to be examined, since, although it is very finely divided, it is nevertheless very dense and to a much lesser degree hygroscopic than the annealing of nitrate of copper.
It is to be remarked, however, that common copper, even in a very finely divided form, would not be fit for the foregoing purpose, as in its application
to bring about the intended reaction the temperature would have to be raised much higher, so that caustic potash rather than nitrate of potash would be obtained; on the other hand, when using the copper prepared from the acetate salt, the reaction occurs at as little as 200° to 250° C. (Annales du Conservatoire des arts et métiers, t. II p. 353.)

Man bereitet sich zunächst durch Destillation von essigsaurem Kupferoxyd sehr fein zertheiltes metallisches Kupfer, und nimmt von diesem frisch bereiteten Metallpulver 2 Aequivalente oder selbst einen geringen Ueberschuß, auf 1 Aequivalent Salpeter, entsprechend der Gleichung
2 Cu + KO, NO⁵ = 2 CuO + KO, NO³.
Persoz nahm 320 Grm. Salpeter und 200 Grm. von dem auf die angegebene Weise dargestellten Kupfer.
Zur Herstellung eines innigen Gemenges löst man den Salpeter zunächst in der möglich geringsten Menge heißen Wassers, und setzt dann das Kupfer hinzu, welches sich anfänglich nur schwierig benetzen läßt. Ist das Gemenge recht gleichartig geworden, so erhitzt man dasselbe in einer Porzellanschale oder besser, in einer gußeisernen Pfanne im Sandbade unter beständigem Umrühren, um Spritzen zu verhüten. Ist die Masse vollständig getrocknet, so tritt ein Moment ein, wo sie, gleich einem Pyrophor, Feuer fängt und erglüht; ist die Verbrennung vorüber, was nur einen Augenblick dauert, so hat die Reaction stattgefunden; man läßt erkalten, behandelt die Schmelze mit Wasser, filtrirt und läßt das salpetrigsaure Salz krystallisiren. Hat man überschüssiges Kupfer angewendet, so ist kein Nitrat vorhanden und man erhält sogleich krystallisirtes Nitrit, welches man dann schmilzt und in gut verschlossenen Flaschen aufbewahrt, da das Salz sehr hygroskopisch ist. Etwa vorhandenes nicht zersetztes salpetersaures Kali wird durch die erste Krystallisation abgeschieden, da es weit weniger löslich ist als das salpetrigsaure Kali. Das bei der Operation als Rückstand erhaltene Kupferoxyd kann nach gehörigem Auswaschen zur organischen Analyse, wenigstens zum Vermengen mit der zu untersuchenden organischen Substanz angewendet werden, da es, obgleich sehr fein zertheilt, dennoch sehr dicht und in weit geringerem Grade hygroskopisch ist, als das durch Glühen von salpetersaurem Kupferoxyd erhaltene.
Zu bemerken ist indessen, daß gewöhnliches Kupfer, selbst in sehr fein zertheilter Form, zu dem vorstehenden Zwecke sich nicht eignen würde, da bei seiner Anwendung
zur Hervorrufung der beabsichtigten Reaction die Temperatur weit höher gesteigert werden müßte, so daß man eher Aetzkali als salpetrigsaures Kali erhalten würde; bei Anwendung des aus dem essigsauren Salze dargestellten Kupfers hingegen tritt die Reaction schon bei 200 bis 250° C. ein. (Annales du Conservatoire des arts et métiers, t. II p. 353.)

woelen - 24-8-2023 at 05:11

Interesting quote. The english translation contains some small, but essential errors. Below follows the corrected translation.

Very finely divided metallic copper is first prepared by distillation of acetate of copper, and from this freshly prepared metallic powder, 2 equivalents, or even a slight excess, are taken to 1 equivalent of saltpeter, according to the equation
2 Cu + KO, NO⁵ = 2 CuO + KO, NO³.
Persoz took 320 grm. saltpeter and 200 grm. from the copper represented in the manner indicated.
To prepare an intimate mixture, the saltpeter is first dissolved in the smallest possible quantity of hot water, and then the copper is added, which at first is difficult to wet. When the mixture has become quite uniform, it is heated in a porcelain dish, or better yet, in a cast-iron pan, in a sand bath, stirring constantly to prevent splashing. When the mass is completely dry, there comes a moment when, like a pyrophor, it catches fire and glows; when the combustion is over, which lasts but a moment, the reaction has taken place; it is allowed to cool, the melt treated with water, filtered, and the nitrite allowed to crystallize. If excess copper has been used, no nitrate is present, and crystallized nitrite is obtained at once, which is then melted and kept in well-closed flasks, as the salt is very hygroscopic. Any undecomposed nitrate of potash present is separated out by the first crystallization, since it is far less soluble than nitrite of potash. The copper oxide obtained as a residue during the operation can, after proper washing, be used for organic analysis, or at least for mixing with the organic substance to be examined, since, although it is very finely divided, it is nevertheless very dense and to a much lesser degree hygroscopic than the annealing of nitrate of copper.
It is to be remarked, however, that common copper, even in a very finely divided form, would not be fit for the foregoing purpose, as in its application to bring about the intended reaction the temperature would have to be raised much higher, so that caustic potash rather than nitrite of potash would be obtained; on the other hand, when using the copper prepared from the acetate salt, the reaction occurs at as little as 200° to 250° C. (Annales du Conservatoire des arts et métiers, t. II p. 353.)

This is quite an interesting thing. I did not know that copper acetate can be converted to copper metal by heating the salt. This is definitely something to try on a small scale. Maybe this process also works with NaNO3 instead of KNO3. The resulting NaNO2 is much less hygroscopic than KNO2.

Also interesting to see that in this old writing potassium ion is considered divalent. They wrote KO,NO5 for potassium nitrate, which we now would write as K(NO3)2. Potassium nitrite is KO,NO3, which in modern notation would be K(NO2)2. Copper(II) at that time already was regarded as a divalent ion.

Alkoholvergiftung - 24-8-2023 at 06:07

It was from 1864.So the lag of knowlege.

woelen - 24-8-2023 at 23:14

This kind of old information can be really interesting. I believe that in old books and papers there is a lot of interesting information about common compounds, which is totally forgotten nowadays. This is exactly, why I love to read old books from around 1900 or so.

unionised - 25-8-2023 at 05:36

Quote: Originally posted by Lionel Spanner

Quote: Originally posted by unionised

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.

That's with 0.1 molar ammonia, i.e. 0.17% w/v. I suspect a more concentrated solution would result in better yields.

[Edited on 22-8-2023 by Lionel Spanner]

I suspect that a higher NH3 concentration would increase the yield of N2 at the expense of NO2- ions.
And the reasons for my suspicion are the law of mass action and the instability of ammonium nitrite.
But, until someone does the experiment, we won't know.

unionised - 25-8-2023 at 06:25

This tells you more than you probably want to know about decomposing copper acetate.

https://www.scirp.org/journal/paperinformation.aspx?paperid=...

Major success- breakthrough

Fantasma4500 - 27-8-2023 at 11:30

Quote: Originally posted by woelen

Very finely divided metallic copper is first prepared by distillation of acetate of copper, and from this freshly prepared metallic powder, 2 equivalents, or even a slight excess, are taken to 1 equivalent of saltpeter, according to the equation

you think thats an interesting quote? this is a marvellous breakthrough in nitrite preparation. i just tried it, mixed 3.2g KNO3 with 3.2g copper sponge i made from copper-copper electrolysis, NaCl + .. 1-3% HCl?
i believe CuCl2 + Al would also give copper fine enough for this
i folded up some aluminium foil and heated the mixture with a blowtorch until it melted and torched the top as well, maybe 5 seconds, it glowed from the flame
dipped that into 10mL water and shook it until the solid came off, black solid
it dissolved effectively, good sign. a few mL IPAlcohol in, a squirt of 10% H2SO4 and both vasodilatory + flame test proved success, i only took maybe 100mg sample to torch

no need for copper acetate decomposition, electrolytic copper sponge is plenty fine

a chemist proposed that Cu2O may also work, but that only makes sense if we can make Cu2O very easily.. CuCl2 + Al goes quite fast especially if you can dump in a bunch of ice as well to keep it under control

this reaction should go very well with a hotplate and pot ontop, metal crucible, insulated with aluminium wool

edit: NaNO3 works too

[Edited on 28-8-2023 by Fantasma4500]

Alkoholvergiftung - 28-8-2023 at 00:06

FeO very fine should work too.

Fantasma4500 - 28-8-2023 at 10:02

Quote: Originally posted by Alkoholvergiftung

FeO very fine should work too.

FeO is made from iron oxalate, i already tried NaNO3 + iron oxalate heated together, poor results
copper powder works. its solved.

Fantasma4500 - 1-9-2023 at 07:25

reporting grand success with electrolytic copper and NaNO3, 120g + 120g copper, heated together at about 400*C for an hour in closed container, black lowdensity crumbling spongy material achieved
120g NaNO3 used, 100g in 400mL H2O acquired, theoretically 120g NaNO3 turns into 97g NaNO2

Alkoholvergiftung - 2-9-2023 at 04:49

to bad if it was closed maybe it reacts like stated at 200-250C exotherm. Would much easier to reach with an hotplate.

Fantasma4500 - 2-9-2023 at 05:34

reporting in that Cu2O also works quite well, may even work better than copper metal
CuO is widely available, Cu2O was made by reacting solution of ascorbic acid with heating with CuO which i acquired from previous Cu + NaNO3 run, looking for alternatives to ascorbic acid as HCl + CuO + Al is also very easy, and aluminium in that case is the reducing agent

CuO and Cu2O is roughly same price, onyxmet lists Cu2O as 4.00 euros where CuO is 3.75 euro for 10 grammes

1188g KNO3 + 1681g Cu2O = 1000g KNO2 + 1869g CuO

CuO seems possible to reduce with sugar and some salt like sodium citrate and copper sulfate, maybe glucose and sodium hydroxide?

[Edited on 2-9-2023 by Fantasma4500]

Rainwater - 2-9-2023 at 06:40

Quote: Originally posted by Fantasma4500

reporting grand success with electrolytic copper and NaNO3, 120g + 120g copper, heated together at about 400*C for an hour in closed container, black lowdensity crumbling spongy material achieved
120g NaNO3 used, 100g in 400mL H2O acquired, theoretically 120g NaNO3 turns into 97g NaNO2

How was the product isolated for yield measurement?

Fantasma4500 - 4-9-2023 at 06:16

Quote: Originally posted by Rainwater

Quote: Originally posted by Fantasma4500

How was the product isolated for yield measurement?

well its a very clean reaction and zero nitrate is claimed by wolens reference, barely any "gunpowder" smell detectable inside the reaction vessel after a whole hour of operating, seeing that after i heated a 10mL solution sample in oven for an hour with fan going, and there was still aggressive amounts of nitrite in it, i wager that i have a solution is quite pure sodium nitrite, i could maybe get a bit extra scientific with using silver nitrate to precipitate it out but im okay with just believing its pure enough for my desires. its a very easy reaction so i invite anyone else more eager for data to repeat it, also for the sake of science and replication

Lionel Spanner - 5-9-2023 at 07:57

Below is a useful quantitative assay I've used in the past for testing the purity of nitrite.

1. Weigh accurately about 1.0 g of the sample of sodium or potassium nitrite to be assayed into a 250 mL volumetric flask, and record the weight (W).
2. Fill the flask about halfway with deionised water, and mix until all solids are dissolved.
3. Once all solids are dissolved, make up to the mark with deionised water.
4. Weigh 5% sulphuric acid (75 g) into a 250 mL conical flask.
5. Add 0.02M (0.1N) potassium permanganate solution (25 mL) to the conical flask via a pipette.
6. Titrate the contents of the flask with the nitrite solution. The end point is reached when the mixture just goes colourless; record the volume of solution delivered (T).
7. The % purity of the assayed sample can be determined by the following equation:
%P = 2156.25 ÷ T ÷ W (for sodium nitrite)
%P = 2626.25 ÷ T ÷ W (for potassium nitrite)

Fantasma4500 - 7-9-2023 at 01:17

just wanted to add CuO as from the reduction reaction can be reduced to Cu2O using glucose and NaOH, boiled for a few hours
the Cu2O reacts with HCl to form a water soluble salt which is black, upon dilution it becomes CuCl, this is the washed and dried Cu2O so no reducing agents leftover in it, cool little bonus.
100g NaNO3 needs 160ish Cu2O, NaOH or H2SO4 may be used to turn sucrose into glucose with heating in solution, straight up glucose seemed to work better tho, ascorbic acid works very well but not as OTC.
KNO3 seems dreadful for making nitrite for some reason, mostly didnt react with some of the copper powder i used, NaNO3 seems much better- probably because its just a bit more reactive

Kevlar - 9-3-2024 at 02:21

I've read this 21 year long thread and deduced one thing, that by far the best method I have seen is a displcement method -

(Ca(NO3)2) + (Na2SO4), The balanced equation is:

Ca(NO3)2 + Na2SO4 → CaSO4 + 2NaNO3

The stoichiometry of the reaction is a nice 1:1 ratio in moles, both of the chemicals are OTC and cheap & the end clean product being pure 2NaNO3. Which is easy to sep from the insoluable CaSO4.

I will buy some of both and come back with pictures, I have got a ref sample of the 2NaNO3 to compare with.

I have been saving up a lot of eggshells and might make the Ca(Na3)2 along with the nitric acid from KaO3 + NaCl reaction.

fx-991ex - 9-3-2024 at 05:34

We trying to make nitrite NO2 not nitrate NO3.

Kevlar - 10-3-2024 at 03:38

fx-991ex I will make the NaO2 from oxidation of the NaO3 using starch, I just got carried away with doing some inorganic chem.

But thanks for pointing that out to me!

RU_KLO - 27-5-2024 at 06:46

Im trying to make NaNO2 for first time.

I choose NO2+NO bubbling in NaOH.

So I made a 4M NaOH solution (100 ml), in which a mix of NO2 - NO gas was bubbled.
NO2-NO gas was made from copper + 7M (aprox) HNO3 dripping (100ml). The dripping was adjusted, as so the inner color of the reaction flask. Mild - almost transparent - red brown color was tried. (i.e. one drop every 1-2 seconds).

All this was in a negative pressure setup (a cheap 12V water/air pump was used).

It took a while - maybe 2 hours and copper needed to be replenished.

After this, ph was meassured, it was highly alkaline (NaOH excess).
(on a test, after adding some acid to a sample, till acidic, there were some red fumes from this solution. So probably some NaNO2 was in solution)
Next HNO3 was added to neutralize the NaOH. Ph was taken to 6-7.

Next a solution of Na2SO3 (unknown molarity) was added.
The idea was to transform the NaNO3 (from the neutralization) to NaNO2.
(some sulfur was in the Na2SO3 and some sulfur was precipitated - or I think it was

At the end was left with 250 ml which was boilled down till 50-40ml which crashed "a lot" of salt. (check picture)

Now, I need to somehow purify and get NaNO2.
Whats in the "soup":
1) 2 NaOH + NO2 + NO → 2 NaNO2 + H2O
2) NaOH + HNO3 = NaNO3 + H2O
3) Na2SO3 + NaNO3 = NaNO2 + Na2SO4 (maybe?) is this correct? (or maybe some sulfur is made?

what are my best route?

1) add more Na2SO3 to convert any NaNO3 left. (will an excess of Na2SO3 hurt?

regarding purification by recristallization, what the best route, because Im not having good results - maybe bad procedure.

NaNO2
Solubility in water
71.4 g/100 mL (0 °C)
84.8 g/100 mL (25 °C)
160 g/100 mL (100 °C)

Na2SO4
Solubility in water
Heptahydrate:
19.5 g/100 ml (0 °C)
44 g/100 ml (20 °C)

Na2SO3
Solubility in water
27.0 g/100 mL water (20 °C).

If I add little water, heat to dissolve everything, and let it crash out, will get the same result.

So how is it done?
a) add boiling water instead of cold till everything dissolves?
b) add cold water till everything is dissolves, remove X% of it (lets say 30%) take it to freezer or ice (0ºC) filter. then remove the last portion till some cristals are seen?
c) filter what you got, dry, weight and from there, add enough ice cold water to dissolve the NaNO2 but not to the Na2SO4
Example: 10grs (dry mixed salt) was obtained,
(solubility per ml NaNO2 = 0.71gr (0ºC); Na2SO4 (0.195gr (0ºC))
So for NaNO2 is 14 ml of water. and for Na2SO4 is 51.28ml.
So if I add 14ml of H2O 0ºC, NaNO2 will be dissolved and excess of Na2SO4 will not.

Or maybe is simplier o more complicated than this (because of common ion effect)

Any help or step by step procedure for purifying is well received.

As side note, a small sample of this salt was tested with HCl and a lot of red fumes were watched.

and as bonus: If you have Sulfite and nitrate, you could directly made nitrite.... Is this correct?

thanks

fx-991ex - 27-5-2024 at 08:11

I dont think mixing Na2SO3 and NaNO3 in sln will reduce to SO4 and NO2. This usually is done in a crucible under high temp(300C) using CaSO3 instead of the Na salt.(for ease of purification)
When you add acid it do make red fumes but that do not mean its NaNO2.

The Na2SO3 react with acid to liberate SO2 gas which probably react with NaNO3 to release NO2 while making Na2SO4.

I did something similar last year, mixing NaNO3 and CaSO3 gave red fumes upon addition of Acid. but after removing by filtration the CaSO3(No CaSO4 as expected), it was only unreacted NaNO3, no red fume upon addition of acid.

What you want to do is, react your Na2SO3 With CaCl2, filter and keep he solid CaSO3. then react this CaSO3 with your NaNO3 under high temperature(300C for 15 minute).
Then you can filter out the CaSO4 and evaporate water to keep the NaNO2.

Also the CaSO3 crystalize as hydrates so more is needed if its not dehydrated and having a slight excess dont hurt things. Most people seem to overlook this part giving bad yield because they dont use enough CaSO3.

[Edited on 27-5-2024 by fx-991ex]

[Edited on 27-5-2024 by fx-991ex]

RU_KLO - 28-5-2024 at 07:27

After "bad" news from fx... and that is possible that I do not have any nitrite, performed the following test.
diferent reagents (used in the procedure) where added to a potassium permangante solution to see if they reduce it.

the nitrite? (prepared solution) and the sodium sulfite reduced it, the rest did not. (check picture)

So as fx said, probably I got no nitrite and the excess sulfite is doing the reduction.

So I will try to take a sample of the nitrite? and add some BaCl2 to precipitate BaSO3 (solubility 0.0011 g/100 mL) This will remove the sulfite and left with nitrite "only" in the solution and check if this new solution reduce the KMnO4.

(note: already tried BaCl2 + Na2SO3 and a fine white precipitate precipitates)

RU_KLO - 2-6-2024 at 14:04

Last experiment was failure.

So i tried the Pb - Nitrate procedure. It seems it worked.

Tested the supernatant, few drops on potassium permanganate in H2SO4, it cleared, so there must be some nitrites.

Currently the Nitrite is drying. Im following Nurd rage permanganate standarization (with Sodium oxalate), so I can follow (previuos in this post) titration of the nitrite/nitrate, so to know how good it was.

attached some pictures.

The sodium nitrite I got is white, not yellow as stated in this post from others.

The PbO I got is creamy white , but pictures show it is yellow...

Are this colors I got correct?

Lead was from UPS batteries, maybe some contaminats are there.

Rainwater - 3-6-2024 at 13:33

PbO is usually black, red, brown or yellow depending on the oxidation state.
That looks like white lead.... googling 2PbCO3·Pb(OH)2_wiki
, heating a sample over a flame it should cause a color change, or react it with an acid to evolve CO2.
Heat it to about 650c and it should turn black if its PbO. Just remember its toxic and will contaminate all your gear

Master Triangle - 8-7-2024 at 14:14

The aluminium reduction of potassium nitrate absolutely works to some extent, I dried a sample and the unpurified salt fumes strongly with HCl.

I used 20g KNO3 in 100ml water, with 4g of Al grinding dust, 6060 extrusion alloy, unwashed so it still had a tiny wax contamination and years old so it is fully oxidised. Stirred and heated on a 55C hotplate. Added a knifetip of Cu2O as a supposed catalyst (no idea if this works in alkaline conditions, or without chloride). A few drops of isopropyl were used to get floaters to be wetted. NO3- to Al ratio is just above stoichiometric if we forget about all the reactions we don't want to happen.

Added ammonia and small amounts of NaOH (to dissolve oxide) with minor effect for several hours until it suddenly took off and started boiling from the very exothermic reaction (this might be a bad idea with Al dust). Continued bubbling for about an hour. It didn't produce nasty amounts of ammonia, so not sure how much of what came off was added by me and how much was from the reaction.

Ended up with a white precipitate with a few Al particles remaining, filtered that away and dried a sample which filled a vial with NO2 from a drop of HCl. Just evaporating the rest down now, hopefully the aluminate is mostly precipitated before nitrite+nitrate begins coming out.

Anyone got ideas for removing aluminate, aluminium hydroxide and any excess potassium hydroxide? Not sure about this system, I added the remaining dried sample back in for the evaporation and it seems there was a fair bit of Al(OH)3 since some would not re-dissolve on heating. The solubility of nitrite in methanol might do it, and might even help separate some nitrate.

Also, the solution was totally clear and has only taken on the typical yellowish tint as I evaporate it. Is nitrite coloured at all? Or does it just get the yellow tint from oxidation?

For repetition, there are a few things unclear that could be important: the potassium cation, whether Cu2O does anything, the alloy used. The speed of the reaction could be important, but a rapid reaction certainly did not destroy all NO2- and probably formed most of it judging by the amount of Al at that stage.

I'll do a permanganate titration at the end, pretty sure aluminate can't interfere?

Master Triangle - 10-7-2024 at 02:08

Yup, a pretty good success!

From 20g of starting nitrate I crystallised out 8g of unreacted nitrate and titrated what was left against permanganate, getting 5.78g or a yield of ~29%. If you go by just the nitrate that was reacted that would be a nearly 50% yield, or even higher if there was a significant amount that I did not crystallise (likely).

I believe ammonia could interfere with the titration but that would have all boiled off during evaporation.

I think I'll try to do some optimisation runs, I could have certainly used more aluminium, and there are a lot of other variables to sort out too.

A lot to sort out with purification too; if you are all good with having some nitrate contamination it might be a good idea to add nitric acid to turn all the KOH into KNO3 which is easy to crystallise out and won't attack your glassware, would probably help with precipitating out Al(OH)3 too. That would have to be done after boiling off the ammonia to ensure you don't make ammonium nitrate.

jackchem2001 - 11-7-2024 at 04:54

I have attempted reduction of KNO3 with CaSO3 a couple of times. I will leave some observations here:

You can't qualitatively monitor the reaction progress by taking samples at various points in time and adding sulfuric acid to them (the intent being to look for the extent of nitrogen dioxide evolution from any nitrite present). However, the initial reaction mixture gives NO2 on addition of sulfuric acid - clearly nitric acid generated in-situ is reduced by sulfite, so this method is no good.
The wiki claims heating sodium nitrite to decomposition (300°C - note a much lower temperature than the potassium salt) leaves behind sodium hydroxide (?) and sodium oxide. In the case of the potassium salt, after leaving the reaction mixture in an open glass dish in an oven at 210-250°C (using an IR themometer) for perhaps 4 hours, when water was added to the cooled reaction mixture some bubbling was seen. I expected this bubbling to be from potassium oxide reacting with water to give hydrogen and potassium hydroxide, but the pH was only very slightly alkaline (i.e. the pH of my tap water).

This suggests to me there may be no disadvantage to leaving the potassium reaction mixture at this temperature range for very long periods of time. If reaction with atmospheric oxygen is feared, then an open container could be using during the initial heating period which could later be closed once the system was up to temperature. I think a glass dish inside of a cast iron pot could be used to accomplish this.
My earlier run of this reaction failed - I suspect due to insufficient heating. I am yet to work up my current run.

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