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Preparation of ionic nitrites

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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





nitirte.jpg - 437kB

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)


test nitrite.jpg - 233kB

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.




NITRITE.jpg - 135kBPBO.jpg - 123kB

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)2wiki
, 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:

  1. 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.

  2. 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.

  3. My earlier run of this reaction failed - I suspect due to insufficient heating. I am yet to work up my current run.

Sir_Gawain - 11-7-2024 at 06:51

I’ve also had a failure with the KNO3 and CaSO3 reaction. I suspect it only works with sodium nitrate.

fx-991ex - 11-7-2024 at 08:07

I did it with NaNO3 only(i read somewhere that it could be dangerous with KNO3, am glad its not, i will give it a try next time).
You need to use alot of heat 300C very minimum, make sure both reactant are very fine and well mixed, slight excess of CaSO3 seem to help.

I think the reaction is very slighly exothermic if you can monitor the temperature closely, i did it in a porcelain crucible over a bunsen burner and after the temp climbed a bit from 200something to over 300 i poked in it with a stainless steel wire and i could see a slight change of the matter, its hard to explain but it didn't "feel" the same(probably just melted a bit from reaching close to 320C boiling/decomp temp). Or maybe thats just me.

Am thinking the CO2 from the bunsen burner probably help to displace some Oxygen to avoid re-oxidation(along with the excess SO3), but for bigger batch the bunsen burner is not practical(hard to heat evenly) so better use an oven at 300-315 C.


[Edited on 11-7-2024 by fx-991ex]

jackchem2001 - 16-7-2024 at 17:54

Thanks for all of the comments
Quote: Originally posted by Sir_Gawain  
I’ve also had a failure with the KNO3 and CaSO3 reaction. I suspect it only works with sodium nitrate.

I hope not, as the decomposition temperature of potassium nitrite is much higher (>500°C) as compared to sodium nitrite (300°C), so would hopefully give a better yield.
Quote: Originally posted by fx-991ex  
You need to use alot of heat 300C very minimum, make sure both reactant are very fine and well mixed, slight excess of CaSO3 seem to help.

I think so too, as leaving potassium nitrate and calcium sulfite at 210-250°C for around 4 hours gave no detectable nitrite (adding sulfuric acid to boiled down extracts gave only nitric acid and no NO2, and adding sulfuric acid to the material that did not dissolve in workup gave SO2 indicating unreacted sulfite). At some point I will try using a loosely closed crucible (avoid oxygen) at much higher temperature and see if I get any nitrite. I am not sure how to deal with the problem of heating the mixture evenly since I don't have an oven that can get any hotter than 250°C.

[Edited on 17-7-2024 by jackchem2001]

Fantasma4500 - 18-7-2024 at 02:29

i thought i would add into the copper metal powder + sodium nitrate method..
CuCl2 is a bit annoying to make, CuSO4 is very readily available and at very cheap price, i was at one point in possession of about 250kg of it for instance.

CuSO4 + NaCl = CuCl2 + Na2SO4
now seperating this is very tedious, so i never took advantage og this, its evidential that this reaction does occur in solution as the NaCl granules when added to CuSO4 sol turns a vibrant green and more so when the next reaction occurs:
Al + CuCl2 = Al(SO4)2 + Cu

this Cu can then be reacted with NaNO3 at melting point temperatures to yield virtually 100% NaNO2
and to recap, the CuO formed can be reacted with NaOH+glucose for an hour to yield Cu2O which can work just as well as Cu powder, alternatively HCl is reacted with CuO and Al is once again added- this method can also be used to produce hydrogen as a byproduct

NaNO3 + Cu = NaNO2 + CuO

i hope the general community will maybe in 2025 be able to grasp the simplicity and blessing that this method, as brought to us by woelen from an ancient chemistry book is, a forgotten gem and a true blessing in current times where nitrites are scarce due to it offering an easy way out of life

Master Triangle - 19-7-2024 at 00:41

Turns out the copper is essential and probably catalytic in the aluminium reduction!

I just re-ran the aluminium reduction with twice the aluminium but no dash of cuprous oxide and the yield is only 15% of the original run. It did not even seem to consume more of the nitrate despite using more aluminium.
I still have not tried the copper reduction despite having some copper powder around. I'll probably give a go eventually for comparison but I really like the idea of using aluminium since it is so much cheaper.

I'll definitely have to try it again with more copper, probably as a Cu(II) salt since I doubt Cu2O is the ideal form for it. I feel like there must be a connection to the Al/CuCl2 reductions that are done for organics like nitroalkenes, which is suspected of being an electrochemical reaction arising from the partial copper plating of the aluminium surface, but could just be from the formation of urushibara copper or some heterogenous Cu/Al(OH)x catalyst.

The fact this is strongly alkaline probably rules out some pathways like acidic dissolution of a copper oxide, so I suspect it would end up relying on ammonia complexation of a copper hydroxide to return the oxidised copper to the aluminium for reduction, since you are not going to end up with a copper chloride or such in a high enough concentration for the reaction rate observed.

I did some test tube reactions and aluminium in ammonia can rapidly reduce copper ammonia complex to the metal but only when enough KOH was added, presumably enough to solubilise the Al(OH)3. I get a bright yellow colour on addition of the KOH too, a stark contrast against the blue ammonia complex, any idea what this is? Only occurs where the KOH is concentrated.

So I guess I just have to find the ideal conditions for reducing copper with aluminium under alkaline conditions, but preferably without adding near-stoichiometric hydroxide. Any ideas?

Master Triangle - 19-7-2024 at 00:59

Quote: Originally posted by Fantasma4500  
i hope the general community will maybe in 2025 be able to grasp the simplicity and blessing that this method, as brought to us by woelen from an ancient chemistry book is, a forgotten gem and a true blessing in current times where nitrites are scarce due to it offering an easy way out of life


I very much hope so too, and am working tower some part of it!

Can't recall what their reference was, but I know some of the early methods of producing alkyl nitrites used nitric acid and sulfuric acid with copper turnings, which would react with the chosen alcohol to produce the nitrite. Ref is Organic Medical Chemicals by M. Barrcliff and Francis H. Carr. I presume this made nitrous acid in-situ, and hopefully dominated over nitration since they didn't make any recommendations on how to not die from distilling ethyl nitrate, but who knows with books that old.

Also, I keep getting "form not secure" notifications and can't post anything unless I have logged out and logged back in for every post. Anyone know what that is?

jackchem2001 - 19-7-2024 at 21:14

An oxidizing agent is one which can accept a lone pair and eliminate an adjacent proton. Nitric acid acts as an oxidizing agent on alcohols - it accepts a lone pair to become nitrous acid and gives the carbonyl. Hence (ignoring acid concentration effects - see later in the post) with one equivalent of nitric acid two equivalents of alcohol ought to be used (to give 1 equivalent of carbonyl and nitrite ester, with the latter being easily seperated by distillation due to its volatility). Further reduction of nitrous acid is not said to occur (the product nitroxyl is unstable).

Boiling points of relevant compounds:

This oxidiation reaction is quite violent and only seems to occur above a certain concentration of acid. The nitric acid may be accessed by sulfuric acid and a nitrate salt, but if you add enough water to dissolve everything then the nitric acid concentration seems to be too low to affect oxidation. Nitrogen dioxide evolution is seen and no external heat is needed to distill something from the mixture.

The reason I say something is because when I did this (using sulfuric acid, potassium nitrate, and isopropyl alcohol without any copper) I got a yellow water immiscible liquid which matches isopropyl nitrite. However, the flame color was orange not white, and its boiling point was 54-58°C. When the crude material was left in contact with water overnight, the boiling point rose to 72-74 degrees - which I might hope is diisopropyl ether forming through SN1. I did not do any other tests with this procedure like fractionating the mixture.

To reiterate the stoichoimetry from earlier: if you draw a mechanism using sulfuric acid with one mol of nitrate salt, one 'proton mol' is needed (but it might be best to use more acid to give the needed concentration for oxidation) and two mols of water are produced. Two mols of alcohol are needed, giving one mol of carbonyl and one mol of nitrite ester. I think it would also make sense to add the isopropyl alcohol dropwise to the acid mixture and not the other way around.

[Edited on 20-7-2024 by jackchem2001]

BAV Chem - 21-7-2024 at 03:43

Quote: Originally posted by Master Triangle  

Can't recall what their reference was, but I know some of the early methods of producing alkyl nitrites used nitric acid and sulfuric acid with copper turnings, which would react with the chosen alcohol to produce the nitrite. Ref is Organic Medical Chemicals by M. Barrcliff and Francis H. Carr. I presume this made nitrous acid in-situ, and hopefully dominated over nitration since they didn't make any recommendations on how to not die from distilling ethyl nitrate, but who knows with books that old.


I tried a similar method of making isopropyl nitrite earlier this year with hydrochloric acid and sodium nitrate instead of sulfuric and nitric acids as those are quite valuable to me. The reaction seems to work quite well at first but after a while it only produces NO instead of the alkyl nitrite. I think this is because after a while water from the HCl starts to build up in the reaction and that leads to side reactions producing NO. In the end I never got yields better than 20% but I might try again at some point with NaHSO4 instead of HCl. I also tried a small scale test with oxalic acid instead of HCl and that did afford quite a bit of isopropyl nitrite but sadly I screwed up later on and lost everything from that run so I don't know about the yield.

Interestingly this reaction isn't even all that dangerous. Yeah it likes to run away in the beginning with HCl but other than that it seems to run along just fine. The isopropyl nitrite always distilled over nicely without decomposing.

Btw i also got a message saying this web form is not secure when i was trying to log in yesterday. Turns out I was on
https://www.sciencemadness.org/talk/misc.php?action=login and changing it to https://www.sciencemadness.org/whisper/misc.php?action=login did the trick.

Master Triangle - 25-7-2024 at 23:24

The source I was reading did use surprisingly little water, so little that I thought they would produce some alkyl nitrates, but I am not familiar with nitrations, maybe that little bit of water was enough to prevent it.

It could be the HCl though, wikiped says HCl can reduce NO2 to nitrosyl chloride, although I thought that would react with alcohols to produce the nitrite anyway (and HCl) in much the same way as nitrosylsulfuric acid. Also, did two layers form? If there is a low concentration of alcohol in the aqueous layer I guess that could lead to NO escaping before the nitrous acid can react with the alcohol, but the production of nitrite certainly shouldn't stop.

Interesting that it might work with oxalic, since copper oxalate is super insoluble. Did it consume the copper, or did it bubble? I kinda wonder if the nitrate just allows the copper oxalate to dissolve or if the oxalate could actually be the main reductant there. Either way it could be a superior synth, was that just with nitrate salt too?

I'll try the alkyl nitrite to nitrite salt reaction once I have enough nitrites, I really want to see if it is efficient within a practical timeframe (sounds like it takes ages), still continuing with the aluminium reduction but getting the salt with a method that uses an alkyl nitrite intermediate might lead to a very pure product free of nitrates and other salts if you can get the NaOH out.

BAV Chem - 26-7-2024 at 14:58

As far as I know the copper itself acts as the reducing agent in this reaction.
The HCl liberated nitric acid from the nitrate salt, that's for sure:
NaNO3 + HCl ---> NaCl + HNO3

This should have then reacted with the copper giving copper nitrate, water and nitrous acid:
Cu + 3 HNO3 ---> Cu(NO3)2 + H2O + HNO2

The latter would of course react with the alcohol, forming a nitrite ester:
ROH + HNO2 ---> RNO2 + H2O

With excess HCl more nitric acid should be generated from the copper nitrate:
Cu(NO3)2 + HCl ---> CuCl2 + HNO3

Now the actual chemistry happening in solution is probably a lot more complicated and involves a few more reactions.

I also got this process from Organic Medical Chemicals by Barcliff and Carr. The amounts of reagents were roughly equivalent to their method with a bit of tweaking to match the new stoichiometry a little better. I wrote some of this stuff down so I can roughly tell you what I was doing there.

Trial 1) HCl:
10g NaNO3
55ml HCl (~23%)
100ml iPrOH
10g Cu (finely cut wire)

These reactions were carried out near the boiling point of the alcohol, so the nitrite would boil right out. It briefly acted up a little at the start, producing some NO2 fumes but then calmed back down. Over time two layers formed, a light green layer of alcohol on top and a nearly black layer of salt solution beneath it. There was some bubbling coming from the copper and I assume that was NO gas arising from side reactions like
2 HNO2 ---> H2O + NO2 + NO
which i guess could have happened in the bottom layer where there wasn't a lot of alcohol to react with the nitrous acid. Iirc the production of nitrite ester never really stopped but just got really slow and inefficient at some point.

Trial 2) Oxalic acid:
10g NaNO3
25g oxalic acid
20ml HCl (~23%; added later)
100ml iPrOH
10g Cu (finely cut wire)

Same reaction conditions as before. This time it didn't run away but was a bit slower than the other run. Interestingly the oxalate doesn't seem to act as a reducing agent itself, it just displaces the nitrate ions from their salts and ends up precipitating copper oxalate, which is just as insoluble as you'd expect it to be. No layer separation occured this time. At some point I added some HCl to hopefully liberate more HNO3 and also because I found an error in my calculations but that didn't really do much (if memory serves).

In both cases quite a bit of the copper dissolved, especially with HCl because a lot of copper(I) chloride formed as well. Anyways the reaction is a low yielding mess with HCl with all the side reactions you get.

Lionel Spanner - 28-9-2024 at 12:18

As many posters in this thread are no doubt aware, highly soluble inorganic nitrites (and nitrates) are very difficult to recrystallise, being extremely soluble in water and nearly insoluble in everything else. I recently found that a 30% aqueous solution of sodium acetate is an excellent solvent for recrystallisation, with the caveat it doesn't work if there's a significant amount of another highly soluble salt present, e.g. sodium hydroxide - in this instance, it has a nasty tendency to freeze solid at fridge temperature due to sodium acetate precipitating out.
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