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Author: Subject: Preparation of ionic nitrites
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[*] posted on 30-8-2009 at 14:16


Some time ago I've tried reductions of NaNO3 with Pb granules, Cd powder and Zn powder (or was it Sn? Sorry, can't remember.) under varying conditions (but always in air, maybe I should try in vacuum or inert gas one day). Never did I get anything close to single phase material. Pb was especially annoying. This method sucks.
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thumbup.gif posted on 2-9-2009 at 08:53
No lead/iron/carbon needed


OK, after reading through this thread, I have concluded that the classic route to nitrites was cumbersome and difficult requiring significant skill and manipulation. There is enough literary material both new and old that stated that simply heating the nitrate of sodium or potassium past fusion until oxygen is evolved produces significant nitrite.

I have made a brief video of me testing the result of heating roughly 1 gram of sodium nitrate recrystallized from the fertilzer"nitrate of soda". This is actually the second time I attempted this pilot run. The first run I tested the solidified mass with concentrated HCl and got copious fumes. Mainly concerned that evolved chlorine could be falsifying the bulk of "the brown cloud of death" so the second attempt was tested with battery grade sulfuric acid.

Here is the link for the video.

I will repeat the test with potassium nitrate since its melting point is higher. Also, the potassium nitrate I will be using was made from ACS grade nitric acid so the possibility of unseen organic contaminates will be eliminated.




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[*] posted on 2-9-2009 at 09:34


Nice experiment. Do you have any commercial nitrites to use as a basis for comparison? It could be my faulty memory but I seem to recall considerably denser fumes when acidifying commercial NaNO2. A titration would be better yet but just an A/B acid test with commercial/homemade should give some idea of conversion.

I agree that the various high-temperature partial reductions of nitrates seem to be unwieldy. I made a few attempts some years ago with charcoal, tin, and lead but didn't try too hard to optimize them after I found a cheap source of technical NaNO2.




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chloric1
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[*] posted on 2-9-2009 at 12:04


If I ever get around to buying that centigram digital scale I could perform this on a 10 gram test basis and do a weight basis for lost oxygen. For example 10.1 grams of KNO3 would yield 8.5 grams of KNO2 if conversion was 100%. Like wise 8.5 grams of NaNO3 would become 6.9 grams of NaNO2. This can be followed with a standardized KMnO4 titration on a solution with a known solids content. With this method I would be happy with a 40% yield because it is cheap and easy. Alot like the women I used to chase when single:D:D. The nitrite from the sodium salt would easily separated by added concentrated silver nitrate solution til no more precipitate forms and chilling in ice salt mixture. The remaining silver nitrate solution decanted and stoichiometric sodium chloride be added to regenerate pure sodium nitrite. The silver chloride processed with washing soda and a oxy-MAPP flame to a pure silver button. Barium nitrite could be prepared in a simular method. The potassium salts can be separated mostly by recrystallization.

I need to get a buret set up too!:(




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[*] posted on 4-9-2009 at 04:50


One of the more interesting articles about decomposition of NaNO3/NaNO2 in the air:
J. Phys. Chem., 1956, 60 (11), pp 1487–1493

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[*] posted on 11-9-2009 at 20:50
NaNO2 from Ethanol, HNO3 and NaOH


It amazes me why so many people persevere with making sodium nitrite using sodium nitrate and lead because, to me, lead is so ugly. Anyway, it seems to work. Way back at the start of this thread there are two methods that are attractive, the first method by madscientist and the sodium nitrate and aluminium method by kingspaz (all the way back in 2002), both need blistering high temperatures to proceed though. Here I present a method that is a low temperature method but I've never tried it practically so I don't know if it works. This method uses nitric acid which may not be attractive to the people who are using sodium nitrate.

Sodium nitrite is made in two steps, firstly ethanol and nitric acid is used to make ethyl nitrite, then ethyl nitrite is reacted with sodium hydroxide to make sodium nitrite.

2C2H5OH + N2O3 = 2C2H5N02 + H2O
C2H5ONO + NaOH = C2H5OH + NaNO2

This is how you can make ethyl nitrite:
Equal volumes of ethanol and common nitric acid are mixed together, and copper turnings (or wire) added, when a quiet action commences by itself, and the distillation is completed almost without heating.

I got this from A Treatise on Chemistry by H. E. Roscoe and C. Schorlemmer, Volume 3 Part 1 (available from the internet archive and library on this site).

Quote:

The compound formed by this action of nitric acid on alcohol is, however, not ethyl nitrate as was formerly supposed, but ethyl nitrite, one part of the alcohol being oxidized, and the nitrogen trioxide, thus formed, combining with another part of the alcohol in the following way :

2C2H5OH + N2O3 = 2C2H5N02 + H2O

Ethyl nitrite thus obtained always contains oxidation-products of alcohol, especially aldehyde, and this turns alcoholic potash brown when shaken up with the liquid.

Ethyl nitrite free from aldehyde is prepared by leading nitrogen trioxide, obtained by heating one part of starch with ten parts of nitric acid, of specific gravity 1.32, into a cold mixture of two parts of 85 per cent, spirit and one part of water. During this operation the heat evolved is so great that the retort must be cooled by immersion in cold ; water, and then the nitrous ether distils over spontaneously. The vapours are condensed in a well-cooled receiver, washed with water in order to remove alcohol, and dried over chloride of calcium. According to Schmidt and Duflos, a small quantity of ethyl chloride is formed at the same time, and for this reason it is better to dry the substance over carbonate of potash.

Instead of leading nitrogen trioxide into the liquid, the gas may be evolved in the liquid itself. In the method proposed by E. Kopp, equal volumes of spirit of wine and common nitric acid are mixed together, and copper turnings added, when a quiet action commences by itself, and the distillation is completed almost without heating. Carey Lea distils 90 cc. of nitric acid of specific gravity 1.37 with 150 cc. of 90 percent spirit and 40 grams of ferrous sulphate, the distillate being freed from ether by shaking with water.

Ethyl nitrite is a mobile liquid possessing a pleasing and yet penetrating ethereal smell, resembling apples or Hungarian wine, and a peculiar pungent taste. It boils at 18, and has a specific gravity of 0.900 at 15.5 and a vapour density of 2.627 (Dumas and Boullay). When ignited in contact with air it burns with a bright white flame. The pure ether can be kept for many years without undergoing any change, but if impure, and especially if it contains water, it soon becomes acid and gradually evolves oxides of nitrogen in such quantities that the bottle containing it frequently bursts. Alkalis, especially in alcoholic solution, decompose it quickly with formation of alcohol.


Sodium nitrite is then made from ethyl nitrite by heating it with a hot solution sodium hydroxide. I'm guessing that the author means an aqueous solution of sodium hydroxide but I'd use an alcoholic solution (see the end of the above quote) and maybe the reaction will occur at room temperature.

I got this from Practical Organic Chemistry by Frederick George Mann and Bernard Charles Saunders, Edition 4, Page 131

Quote:

Since aliphatic hydrocarbons (unlike aromatic hydrocarbons, p. 155) can be directly nitrated only under very special conditions, indirect methods are usually employed for the preparation of compounds such as nitroethane, C2H5NO2. When ethyl iodide is heated with silver nitrite, two isomeric compounds are formed, and can be easily separated by fractional distillation. The first is the true ester, ethyl nitrite, C2H5ONO, of b.p. 17°: its identity is shown by the action of hot sodium hydroxide solution, which hydrolyses it, giving ethanol and sodium nitrite.

C2H5ONO + NaOH = C2H5OH + NaNO2

The second compound is nitroethane, CaH6NOa, of b.p. 114°: its identity is clearly shown by the action of reducing agents, which convert it into ethylamine.
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[*] posted on 12-9-2009 at 04:45


Great first post! The nitric acid could be made in-situ with nitrate on hand and sulfuric acid. Really handy for one who can buy sulfuric at $20 for 15L and 25kg bags of KNO3 :).

This method also saves the trouble of having to separate nitrites and nitrates. Swamped with exams, trying this in around a week or so (along with the 100s of other things I have to try!).

Formula409.
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[*] posted on 13-9-2009 at 06:57


NaNO2 from Ethanol, HNO3 and NaOH
It is hard to find equally useless method of preparation NaNO2.
First, you have to prepare "N2O3" (from HNO3 and starch ) and then C2H5ONO, with boiling point 17 C, next hydrolysis of nitrite.
Long and low-yield method.
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[*] posted on 13-9-2009 at 08:29


It looks like from the reference madcedar presented, N2O3, is prepared in-situ by the action of ethanol and copper on nitric acid. This would make the reaction less tedious, but even still I am inclined to agree with kmno4.



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[*] posted on 15-9-2009 at 12:29


I do not agree with Kmno4 and ammonium isocyanate. The method may have its troubles, but I like the new and uncommon point of view of this method. All the other high-temperature methods have their own serious troublesome parts and especially when you want to do things on a small non-industrial scale then things may be very hard.

I am inclined to try it on a microscale and see how well it fares. Making ethyl nitrite without the need of having nitrites around looks like an interesting thing on its own. The most doubtful part of this process seems to me the reaction of ethyl nitrite with sodium hydroxide. How fast is that reaction?




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[*] posted on 18-9-2009 at 06:59


Quote: Originally posted by madcedar  


I got this from Practical Organic Chemistry by Frederick George Mann and Bernard Charles Saunders, Edition 4, Page 131

Quote:

Since aliphatic hydrocarbons (unlike aromatic hydrocarbons, p. 155) can be directly nitrated only under very special conditions, indirect methods are usually employed for the preparation of compounds such as nitroethane, C2H5NO2. When ethyl iodide is heated with silver nitrite, two isomeric compounds are formed, and can be easily separated by fractional distillation. The first is the true ester, ethyl nitrite, C2H5ONO, of b.p. 17°: its identity is shown by the action of hot sodium hydroxide solution, which hydrolyses it, giving ethanol and sodium nitrite.

C2H5ONO + NaOH = C2H5OH + NaNO2


Hmm anyone ever tried adding ethyl iodide to silver nitrite? I certainly know the author of this hasn't (i'm referring to Mann and Saunders not you Madcedar), which undoubtedly makes woelens query regarding the speed of the caustic hydrolysis of the ethyl nitrite pertinent because the author hasn't tried it himself.

Instead of all that fuss of making the ethyl nitrite in the manner described (which i'm not con vinced would work super well anyway) just make the ethyl nitrite ala Vogel 3rd edition, one only requires ethanol, H2SO4 and NaNO2. Then you can decompose ethyl nitrite and obtain the NaNO2 you require.

Hey Chloric there's a new product on the market called a tripod, what it does is holds a camera, so your video doesn't end up looking all Blair Witch and giving everyone watching it on motion sickness.:P:P:P


[Edited on 18-9-2009 by Panache]




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[*] posted on 18-9-2009 at 07:58


Quote: Originally posted by Panache  
[
Hey Chloric there's a new product on the market called a tripod, what it does is holds a camera, so your video doesn't end up looking all Blair Witch and giving everyone watching it on motion sickness.:P:P:P


[Edited on 18-9-2009 by Panache]


Actually I got a functional, but obviously used, tripod for $4 at Goodwill:cool: I just would not be as much fun to use it if I was not able to illicit a smartass comment on this forum.:D

No but seriously I did not use it because the video was rather short. And besides the Blair witch effect adds to the "mad science" appeal.

@ woelen I feel that the simplest ways to nitrite are NOx to sodium base solution. Maybe liquify some NO2 and bubble NO through it and add combined NOx to soda ash solution. The NO is easily made with diluted nitric acid on copper scrap out of contact with air.




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[*] posted on 18-9-2009 at 09:43


I read in my lab textbook that Cu(I) can reduce Ammonium Nitrate into the Nitrite. The NH4NO2 then decomposes into N2 and H2O in the presence of acid, hopefully not all at once!

Could this also work on KNO3 to form a stable nitrite? The Cu(II) formed could be recycled with sulfite.

I will find it again when I get home to see what the important details are.
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[*] posted on 20-9-2009 at 01:18


I agree with kmno4 and ammonium isocyanate when they say this method may be low yielding. Like I said before, I haven't tried it so I don't know. It seems to me to be a really easy procedure, the ethyl nitrite can be lead straight from the reaction mixture and bubbled through a solution of NaOH in alcohol. This is where this part of my first post is important:

Quote: Originally posted by madcedar  

Quote:

...

The pure ether can be kept for many years without undergoing any change, but if impure, and especially if it contains water, it soon becomes acid and gradually evolves oxides of nitrogen in such quantities that the bottle containing it frequently bursts. Alkalis, especially in alcoholic solution, decompose it quickly with formation of alcohol.


From that it seems to me that the worries of woelen and Panache may not be a problem. It looks like the NaNO2 will fall off the ethyl nitrite easily and as luck would have it, NaNO2 is not very soluable in alcohol so it should drop out as a solid which makes isolation easy.

This is from the Merck Index
Quote:

Sodium Nitrite
Sol in 1.5 parts cold water, 0.6 part boiling water, slightly in alc.

Sodium Nitrate
One gram dissolves in 1.1 ml water, 0.6 ml boiling water, 125 ml alcohol, 52 ml boiling alcohol, 3470 ml abs alcohol, 300 ml abs methanol.

Sodium Hydroxide
One gram dissolves in 0.9 ml water, 0.3 ml boiling water, 7.2 ml abs alcohol, 4.2 ml methanol, also sol in glycerol.


I'd love to try this method out myself but I'm not set up to do such things at home. I'd also love the try the sodium nitrate heated with aluminum (foil) because that seems like a very attractive procedure to use at home.
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[*] posted on 31-10-2009 at 07:56


Polverone in reply to the method put forth on the very first page of CaSO3 + Nitrate this should work as well using the Sodium salt as well and since the solubility of Sodium sulfate is so low compaired to something like
Potassium nitrite seperation should be no problem. Will this reduction take place as well in an (aq)solution?

I want to ask anyone that has tried various methods what there favorite reaction for reducing to the nitrites is. I have tried Carbon reduction but that always seems to go two far and I end with quite a bit of lost nitrite/nitrate. I have melted the Potassium nitrate and slowly added carbon to it and I have grinded the two together igniting the mixture and they both produce less then par results.

I have also performed the lead reduction on a couple different occasions but I find it impossible to work with this shit responsiblely. No matter what I do I get a metallic taste in my mouth when finished even though I have taken every precaution such as a dust mask, gloves, coveralls, shower afterwards ect... nothing seems to work well at all and I am quite frankly done screwing around with the moltan lead method even if it does produce the best results I have seen yet. When it comes down to it this produces to many hazards such as hot molten lead, extremly finely divided litharge and nitrite solutions non of which are remotely healthy to be around.

I want to try your sulfite method Polverone since it seems as though it would be straight forward and high yeilding but right now I have no CaCl2 and only Sodium Metabisulfite so I am going to look and see what I can do here in terms of finding an alternative means other then lead.


PS: Came across this showing a visual indicator of nitrate reduction thought might be of some interest to some

Source:http://web.clark.edu/tkibota/240/Unknowns/Nitrate.htm

[Edited on 31-10-2009 by Sedit]





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[*] posted on 31-10-2009 at 09:52


Sodium Nitrite
from
Inorganic Chemical Preparations
Erdmann & Dunlap ©1900



NaNO2a.jpg - 79kBNaNO2b.jpg - 97kBNaNO2c.jpg - 74kB
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[*] posted on 31-10-2009 at 13:56


@Sedit

Zinc does indeed reduce nitrate in aequos solution but it won't stop at the nitrite unfortunately :(

I also hate fiddling with lead, especially since fusing the mixture and extracting with boiling water produces copious amounts of toxic aerosols.

Fe is a promising candidate as an alternative reductor. I have an old chemistry book that claims NaNO3 is reduced *quantitatively* to sodium hyponitrite by fusion with iron powder. This is questionable but for sure iron can reduce nitrate, as shown in the preparation of ferrates from KNO3. If the fused mass is dissolved in boiling water, the ferrate will decompose and insoluble Fe hydroxide/oxide(?) will settle leaving a yellowish solution of - well either nitrite or hyponitrite, or possibly both.

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[*] posted on 31-10-2009 at 13:58


Some info on sodium hyponitrite.

Attachment: HYPONITRITE.pdf (514kB)
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[*] posted on 31-10-2009 at 14:37


Here is a cryptic reference to reduction of NaNO3 by adding lime and then SO2 and filtering off the CaSO4. Sorry if this has been discussed, but it seems worth trying on a small scale.
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[*] posted on 31-10-2009 at 14:45


My main concern is purity of the formed Potassium nitrite hence the reason I like the lead reduction leaving no soluble biproducts like sulfates and such. When starting with Potassium nitrate however this isn't as big of a problem then when trying to reduce Sodium nitrate because the solubility difference between the nitrate and nitrite of Potassium is pretty large.

There are a few different metals other then lead that should perform the reduction pretty well but Lead has the low melting point which makes exposure of fresh metal surface very easy with a little stirring. I once tried Lead in a ballmill and this indeed reduced a good deal of Sodium nitrate but there is still the issues of Lead around which I didn't care for. Perhaps I will try to ball mill a slightly damp slurry of Copper turnings, Silica, and Potassium nitrate to see how well that works. The Silica should aid in exposing fresh copper to the nitrate and allow further reduction.

I have also tried in the past to reduce an (aq) solution of Sodium nitrate with Al in a mildly alkaline solution before. It seemed to be going well but it would run away after a while and when the temperature reached to high ammonia would bail out of the solution by the ton. I think Aluminum may just be able to reduce it well in solution if kept cool enough to not allow further reduction.

I just tryed a little bit ago for shits and giggles to mix an equal volume of Sodium Metabisulfite with Potassium nitrate and as soon as it started to melt and react LARGE amounts of NOx fumes begun to bail out of the mix so I may just try this again and capture the formed fumes in water or alkaline solution. Honestly haven't given the reaction taking place much thought since I haven't really had time to between trick or treating.

Taoiseach thanks for the PDF but my computer crashed and it will be a few days before I am able to get adobe back to read a damn thing on here. I lost hundreds of GB of files and about 3GB I was going to add to my library.





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[*] posted on 31-10-2009 at 15:17


Sedit, I have used the nitrite test you showed above and it works well, but it is so sensitive that it will detect any small amount of nitrite. That test is used to identify bacteria that reduce NO3 to NO2, which raises two points:

Some bacteria reduce NO3 to NO2, some reduce it all the way to N2. The test is done using a growth medium containing KNO3. So a culture of the right bacteria might even be used to produce NaNO2 or KNO2.

If the test for NO2 is negative, the procedure is to add Zn powder. If NO3 is still present, it is reduced to NO2 and the tube turns red, indicating that the bacteria did not reduce the KNO3. So, Zn definitely reduces KNO3 to KNO2 in aqueous and might be used in a prep under the right conditions.
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[*] posted on 31-10-2009 at 15:33


I came across a patent a while ago(just looked for it again to link it but I can't seem to locate).It claimed quantitative conversion of nitrate to nitrite.Basically a zinc(or cadmium) copper couple is set up in column.Multiple passes of the nitrate in solution (about 10 from memory) is said to do the trick.



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[*] posted on 2-11-2009 at 06:34


Will sodium sulfate displace calcium nitrite? Calcium nitrite is widely available and cheap as chips:)
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[*] posted on 2-11-2009 at 06:38


Calcium nitrite? where can one obtain that 'as cheap as chips'?
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[*] posted on 2-11-2009 at 07:13


unome: I believe you mean calcium nitrate, as that really is cheap as chips. Sodium sulfate will displace this, as the reaction is driven forward by the precipitation of the insoluble calcium sulfate. If however, you do actually mean calcium nitrite, then the reaction will work the same way, and you'll end up with a solution of sodium nitrite and a precipitate of calcium sulfate.
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