Sciencemadness Discussion Board

Preparation of ionic nitrites

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guy - 17-8-2007 at 11:16

Yes Pb(NO2)2*H2O is yellow (reference: http://journals.iucr.org/c/issues/1985/10/00/a25000/a25000.p... )

I suppose this may be removed by addition of sulfate.

--------

I think PbO can be used as a catalyst for nitrate -> nitrite.

3PbO + KNO3 --> Pb3O4 + KNO2

Pb3O4 --500C--> 3PbO + 1/2 O2
------------------------------------
KNO3 --PbO/500C-->KNO2 + 1/2 O2

Is this plausible?



[Edited on 8/17/2007 by guy]

12AX7 - 17-8-2007 at 14:22

What about 1/2 O2 + KNO2 = KNO3?

Tim

JohnWW - 17-8-2007 at 14:58

Quote:
Originally posted by guy
Yes Pb(NO2)2*H2O is yellow (reference: http://journals.iucr.org/c/issues/1985/10/00/a25000/a25000.p... )
(cut)

That file, and indeed the whole site (of crystallography journals) requires a log-in name and password, i.e. it is censored!! Please give these to us, Guy.

guy - 17-8-2007 at 15:53

Quote:
Originally posted by JohnWW
Quote:
Originally posted by guy
Yes Pb(NO2)2*H2O is yellow (reference: http://journals.iucr.org/c/issues/1985/10/00/a25000/a25000.p... )
(cut)

That file, and indeed the whole site (of crystallography journals) requires a log-in name and password, i.e. it is censored!! Please give these to us, Guy.


Oh...Well I was accessing from my university so I didnt need to type in anything. Sorry!

guy - 17-8-2007 at 16:28

Quote:
Originally posted by 12AX7
What about 1/2 O2 + KNO2 = KNO3?

Tim


Oh yea you are right. darn.

Filemon - 23-8-2007 at 13:23

Quote:
Originally posted by Polverone
Quote:
Originally posted by Filemon
Quote:
Originally posted by guy
Na2CO3 and KNO3 does not make NH3.


I have also surprised. But smelled ammonia. You prove.

I would guess that your KNO3 is impure and contains some ammonium salt. Do you smell ammonia if you mix it with a small amount of NaOH or KOH and add a couple drops of water?


Not this impure one the KNO3 but Na2CO3. I don't know that impurity has but I have discovered that this quite impure one.

JoPann - 15-12-2007 at 05:55

Quote:
Originally posted by guy
Quote:
Originally posted by 12AX7
What about 1/2 O2 + KNO2 = KNO3?

Tim


Oh yea you are right. darn.


Hmm I don't think this would happen, because at higher temperatures KNO3 decomposes on its own to KNO2

I also made some experiments by pyrolysis of KNO3 without any reducing agent and got pretty bad results. The molten mass coloured greenish yellow but when adding some HCl to a solution of the product not much gas was produced.

I also noticed, the molten mass dissolved the glass layer of my crucible.

And I tried to separate the nitrite from the nitrate by dissolving the salts in ethyl alcohol. In some MSDS I've read, there is a difference in the solubility between the KNO2 and KNO3. So the KNO2 goes into the solution.
But i haven't got more detailed informations about this process and in my case the Nitrate doesn't disturbs, because at the moment i just need the nitrite for Gas production and some "Azo" syntheses.

Btw.: I'm about to modify the Mushpratt method in adding charcoal to molten KNO3, so I think the procedure is calmer.

[Edited on by JoPann]

Jome - 27-5-2008 at 15:36

I once tried to make some nitrite through incomplete combustion of a mix of sugar and saltpeter. With HCl a simple test on the residue after 10g of random mix gave a voluminous amount of brown gas.

A week ago, I tried again, using three mixes. The first was 1:1, the second 1:2 (close to stoich) the third 1:3, which was hard as hell to get to burn due to the low amount of fuel.

This time I used a solution of KMnO4 on the extract and assuming the reaction went
H2O + 2KMnO4 + 3KNO2- ---> 2KOH + 2MnO2 + 3KNO2, the amount of KMnO4-soln to cause a permanent color change was measured by adding 1ml at a time.

Yields turned out to be very low, something like 1.5-2% based on original KNO3. YES it was a sloppy half-assed experiment but IMO the procedure does not show promise at all. I theorize strongly colored ClNO from the reaction with HCl tricked me into believing it was a generous amount of NOx.

Did an experiment today on the liquid alkali-nitrate reduction of carbon (60/30/5 KNO3 KOH C)
Water was added to homogenize (it turned into a goop) but the mix kept foaming even after being "dry" the first time. Was this CO2? Or just water from the KOH? Gradually it evolves slower, whatever it is.

Then all of a sudden a cloud of smoke rises, and within 2 seconds the (stainless steel!) beaker is rescued from the fire.

The cake was dissolved in 150ml of boiling water. 4ml of this was diluted to 50ml, and a 2mg/ml solution of KMnO4 was added two ml's at a time (stupid dripper)

Color change immediately, meaning a maximum of 0.3% of the intended 50g KNO2 is present. Only option here that could "save" this method would be if the smoke-evolution was some kind of runaway which destroyed all nitrite.

Time to aquire some lead, I guess.

Sedit - 16-3-2009 at 09:25

[I moved this to where is should be]

I have attempted performing the reduction of sodium nitrate using NaOH+S in nothing more then an informal test run. Mainly just to see what I would be getting into trying to reduce about and ounce and a half of nitrate.

All materials where pretty much eyeballed and a table spoon of each was mix and ground in a morter and pestle.
A small sample was taken out and put into a small stainless steal crucible and melted with a small hand torch with which it promptly ignited and sustained it self for a few seconds.

First thing to note as mentioned a few post above is that it is imperative to ball mill all this together else there is mini explosions that will splash molten melt everyware. Even grinding it as fine as possible it would still hit occasional "hot" spots and produce a small bang.

When it is gone to compleation there is a ring of redish substance and a dark grayish black solid lump.
HCl was driped on a piece of the lump I broke off and there was a puff of brown NO2 so that is a good sign I assume.


Question:Is there any substance that can be added to this mix to quell the pyrotechnics of it and still allow it to perform the reduction? It looks like it could be promising to me but the 7 inch high rapid flames comming of a small pile the size of a quarter makes me think that this performed on an ounce of nitrate could get ugly a little to quick for my liking.:o

Sedit - 19-3-2009 at 19:42

Sorry for the double post here but I can not edit any longer.


Ok Im on a little quest here to make some sodium nitrite and Iv noticed a wiki artical on ammonia nitrite that speaks of ammonia nitrites being produced thru the oxidation of ammonia with H2O2 or ozone.

Does anyone think that it would be feasable to produce sodium nitrite from the reaction of ammonia nitrite and NaOH?

Im thinking along the lines of feeding dry NH3 into my 35 percent H2O2 and gently boiling it down till I have a concentrated enough solution of nitrite and then reacting this with NaOH to yeild sodium nitrite and ammonia hydroxide.

Any one have any thoughts on this method before I have a go at it within the next few days? It would seem to me that it would yeild the purest of any nitrites discussed so far.

497 - 20-3-2009 at 02:38

That seems like a waste of good H2O2.. But other than that I don't see why it wouldn't work, if in fact H2O2 does oxidize NH3 at a reasonable rate. The yields might suck too, since NH4NO2 is so unstable.. I have doubts about whether it could even survive being boiled down from a dilute solution.

I think it would be more interesting and novel to build an ozone generator and mix dryish NH3 with the O3 stream to directly precipitate the NH4NO2. That might be an interesting route to HNO3 from only ammonia and electricity, if it worked well at all.. Hehehh, then decompose it to NO2 and run that back through the ozone stream to get some N2O5!

Sedit - 20-3-2009 at 06:46

As for wasting H2O2 this really isnt a concern of mind considering I can aquire this for about $20 for a quart of it very easyly yet nitrates and nitrites are becomming problematic.

Is there any reason that the NaOH and H2O2 can not be premixed and the NH3 feed into this solution to produce the more stable NaNO2 directly without having to boil the solution? I think that saturation of an H2O solution with NH3 would may be the best way to proceed then react the proper amount of H2O2 followed with the addition of Sodium hydroxide to form the sodium nitrite.

.Iv considered the use of O3 because I have a few table top homemade tesla coils and these will bail out ozone as theres no tommorow but then theres the issue of getting two gas generators to run perfectly to produce a solid. Im thinking it may be best to feed these into some sort of solvent to allow them to react there. Possibly some Alcohol or perhaps even H2O

Sedit - 21-3-2009 at 13:30

Starting small and improper as usual I attempted to add some 35% H2O2 to over the counter ammonia solution just to see if what I expected to happen would, It didn't.

First of all I expected to have the smell of ammonia disappear after the addition of H2O2 and the ammonia was oxidised to ammonia nitrite. Im not really sure but the smell may have subsided quite a bit but it is still noticeable and the solution took on almost a slight halogen smell like a dilute bleach smell.

Seeing as the faint NH3 smell could possibly be due to decomposition of the unstable nitrite I am going to attempt to scale it up and do it properly.

One concern I have at the moment is the so called NaOH I have is a mixture of NaOH/KOH which will make proper ratios impossible so until I obtain some pure NaOH I am not going to be able to do it the correct way.


Question: If oxidisers such as O3 and H2O2 work to oxidise the ammonia would it not be possible to use other oxidizers such as permagnates and the likes?

[edit] because after sitting over night the smell is greatly reduced.... but still there.

[Edited on 21-3-2009 by Sedit]

Sedit - 30-5-2009 at 14:19

I requested this paper sometime ago with no response so instead of wait I decided to just attempt a process that I felt may work so I did not have to hassle with hot lead and splattering PbO.
Quote:
Reduction rate of sodium nitrate by lead drops with wet-ballmilling

Affiliation(s) du ou des auteurs / Author(s) Affiliation(s)
(1) Department of Applied Chemistry, Graduate School of Engineering, Tohoku University, Sendai 980-77, JAPON

Localisation / Location
INIST-CNRS, Cote INIST : 21521, 35400006896014.0130

http://cat.inist.fr/?aModele=afficheN&cpsidt=2808050


If someone can get this for me it would be great.


Anyway it appears to work great although I never isolated yeilds because... well Im stupid and let it sit out to dry just to have it liquify on me.

I misplaces my lab notes and it was about a month ago yet I just tested it for NOx production and the flask was filled with a dense brown gas.

20 grams of NaNO3 where added to a ball mill with a 2x molar excess of lead shot smashed flat for more surface area. Also a small amount of course hardware store sand was added to assist in the removal of any PbO that will form on the lead surface and a little bit of h2o just enough to make a sludge.

I milled this for about two days which is more then likely over kill but with out the text I have no idea with out extensive runs how to optimise this. After I added hot water and filtered the mix (which I still hold the filtrate of and may be able to determine yeild based on PbO to Lead acetate and weigh). The filtered solution was slowly dried over a very gental heat and left out overnight to liquify again and lose some of it. :P

Either way when I just found the remains I set aside again as I was testing to see if NH3 feed into H2O2 formed nitrites(They do a little bit but thats another post all together) I droped a small amount into H2SO4 and a thick dark brown cloud came out only slightly less then when it was formed from molten lead. This slightly lighter color could also be due to its damp nature also.

Jimmymajesty - 24-8-2009 at 07:10

I tried the starch + diluted nitric acid and the sugar+KNO3 mehod as these ones seemed the most conveninet to make nitrites.
Maybe I missed something, but I failed in both cases, the yields were terrible or none at all.

In case of sugar+KNO3 I simply can't believe that someone could make a decent amount of nitrite, it seems that the product decomposes at that temperature and mainly K2CO3 forms. And you will definitely piss off the neighbours as a lot of fumes will be generated even with test tube quantities.

In case of starch+diluted nitric acid, it is hard to make 1:1 NO+NO2 gas mixture, so a lot of NO bubbling through the scrubber unused + I think you mainly generate CO2 which react with the NaOH giving carbonate instead of nitrite.

I dunno what to do next, sniffing lead vapors isn't something that my bone marrow would thak me... maybe the sulphur method.

Taoiseach - 29-8-2009 at 07:39

I had a go at the lead+NaNO3 method also. A coal fire in an old saucepan was used as heat source.

I first melted the NaNO3 and then added Pb according to

NaNO3 + Pb ---> NaNO2 + PbO

The Pb came in the form of 1mm lead shot. The reaction proceeded smoothly and quite some bubbling was observed.

n001.jpg - 69kB

Taoiseach - 29-8-2009 at 07:41

This is the end product. The mixture became thicker and thicker and in the end it was impossible to keep it thoroughly melted. I kept heating for another 30 minutes and let it cool down.

n006.jpg - 72kB

Taoiseach - 29-8-2009 at 08:08

Altough difficult to see in the image, quite a lot of lead remained unreacted. I extracted the fused mass with hot water, filtered and started boiling.

The filtrate was quite yellowish. A lot of unreacted lead was recovered tough.

First a huge amount of clear crystals formed which I filtered and tested with HCl. No reaction so this must be NaNO3 :mad:

The final yield was pathetic, maybe 80g of slightly yellowish crystals. They do fizzle and bubble a bit when thrown into HCl but its not nearly as intensive as with reagent grade NaNO2. Also the color is just slightly off-white whereas the reagent grade stuff is much more yellowish. I discarded it as I have no used for impure crap of unknown composition.

At least I could use the PbO to make Pb(NO3)2.

----

I wonder what I have done wrong here. Possible sources of error I have come up with so far:

-too low a temperature (not likely as the reaction proceeded very quickly in the beginning and then slowed down rapidly)

-NaNO2 oxidized to NaNO3 by O2 from air (maybe air must be excluded from the cooling mass and the filtrate as its boiled down). HOW PRONE IS NaNO2 TO AERIAL OXIDATION?

-Not enough stirring to react all Pb (quite possible - the mixture became very thick and unstirrable after the initial reaction subsided)

Anyways this is not a promising route to NaNO2 IMHO.

Possibly the easiest way to reasoably pure NaNO2 is to bubble NOx into a sodium peroxide solution/NaOH + H2O2 mixture.

ALSO a C+Ca(NO3)2 mixture might work well. Yields might be shitty but the CaCO3 formed is insoluble so at least the nitrite would be easy to isolate.

not_important - 29-8-2009 at 08:55

This is the classic route to sodium nitrite, used for a long time and documented in many books on inorganic chemistry, including many lab methods and preparative chemistry books.
Quote:

INORGANIC CHEMICAL PREPARATIONS
BY
FELIX LENGFELD

SODIUM NITRITE
Hampe, A. 125, 336.
Melt 85 grammes sodium nitrate in a
small, flat, iron dish (sand-bath), and add
slowly 206 grammes granulated lead. Stir
until the lead is completely oxidized (30-50
minutes), and pour the molten mass upon
an iron or slate slab. Powder and wash
thoroughly with boiling water. Cool the
filtrate, pass in carbon dioxide for a few
minutes, heat to boiling, and filter hot.
Evaporate until crystals separate on cooling.
Unless the evaporation has been carried
too far, the crystals are almost pure sodium
nitrate. Pour off the solution, again evaporate
and cool, and repeat until the solid
shows a decided test for nitrite. Then
evaporate the mother liquor to dryness, and
extract repeatedly with boiling absolute
alcohol. Evaporate the alcohol and recrystallize
from water. To determine the
strength of the nitrite, dissolve a weighed
portion in considerable water, add the quantity
of potassium permanganate it would
require if pure, acidify with sulphuric acid,
and titrate the excess of permanganate.


Things that may have gone wrong include adding the lead too quickly which can chill the melt, failing to maintain the needed temperature, and the NaNO3 being too moist (it does absorb moisture) so the the melt contains too much water and is too cool to start with.

The reaction is somewhat exothermic, on a larger scale it sometimes glows where the lead is being added. The lead does need to be finely divided, or added slowly with good stirring. In a related preparation, barium nitrite is made from a boiling solution of barium nitrate with addition of powdered lead made by reacting a solution of lead acetate with magnesium ribbon.

Calcium nitrate tends to go all the way to the carbonate/oxide.


Sedit - 29-8-2009 at 09:08

I am confused Taoiseach. Personaly I feel that this is the most promising of all nitrite routes I have seen. Perhaps its about order or something. Opposed to how the text N_I quoted I find it best to melt an excess of Lead in a crucible and slowly add the nitrate with stirring. I have plans on performing this later tonight time permitting and as long as I can find some spectricide at my local stores.

When adding the nitrate to the lead it will quicly melt and turn the lead bright orange I stir this for a few minutes before adding another addition about a gram at a time. It is a little more extensive I assume but in the end addition to H2SO4 produces a dark cloud of NOx.

Perhaps you are evaporating the solution at to high of a temperature. I prefere to do it at a low temperature to prevent oxidation back to the nitrate.

Sedit - 29-8-2009 at 09:08

double post?????

[Edited on 29-8-2009 by Sedit]

kmno4 - 29-8-2009 at 14:00

This procedure works, read my post on 4-th page in this thread.
NaNO2 is dirty cheap, about 4-5 dollars per 1 kg.

JohnWW - 29-8-2009 at 14:20

What do you want to use the NaNO2 for? I suppose you could use it as a meat preservative, or for diazotization of aromatic amines (and thence formation of aromatic halides and phenols and nitriles, and aromatic diazo compounds including dyestuffs), or formation of nitrosylamines.

Taoiseach - 29-8-2009 at 14:55

I'm also confused. Obviously I did something wrong - but what?
I did everything as described in literature. Even passed CO2 thru the filtrate but no cloudiness appeared. I always wondered how plumbate complexes could possibly form as they require strongly alkaline conditions.

Maybe the mistake was that I let the melt cool down without covering it. It might have drawn air. Also I did not extract the mass with hot water but rather filled water into the soupcap and boiled it for quite some time. Maybe the PbO re-oxidized the NaNO2 or it was oxidized by air.

I could not find any information on how easily NaNO2 is oxidized to NaNO3. Does NaNO2 have to be protected from air?

@not_important

The NaNO3 was definetely dry as I first heated it for quite a while until a clear liquid was obtained.

Also the fact that obviously quite a lot of PbO was formed suggests that reduction did take place, so I start to believe the NaNO2 was re-oxidized/destroyed later.

kmno4 - 29-8-2009 at 15:54

I remember that somebody asked for thermal stability of NaNO2 and NaNO3 in the air but I cannot find this post.
There is article about it (from Springer), but I do not have it at hand.

turd - 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.

No lead/iron/carbon needed

chloric1 - 2-9-2009 at 08:53

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.

Polverone - 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.

chloric1 - 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!:(

kmno4 - 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

Attachment: j150545a005.pdf (296kB)
This file has been downloaded 4830 times

NaNO2 from Ethanol, HNO3 and NaOH

madcedar - 11-9-2009 at 20:50

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.

Formula409 - 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.

kmno4 - 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.

ammonium isocyanate - 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.

woelen - 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?

Panache - 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]

chloric1 - 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.

mr.crow - 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.

madcedar - 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.

Sedit - 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]

Vogelzang - 31-10-2009 at 09:52

Sodium Nitrite
from
Inorganic Chemical Preparations
Erdmann & Dunlap ©1900



NaNO2a.jpg - 79kBNaNO2b.jpg - 97kBNaNO2c.jpg - 74kB

Taoiseach - 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.

Taoiseach - 31-10-2009 at 13:58

Some info on sodium hyponitrite.

Attachment: HYPONITRITE.pdf (514kB)
This file has been downloaded 878 times

entropy51 - 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.

Sedit - 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.

entropy51 - 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.

starman - 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.

unome - 2-11-2009 at 06:34

Will sodium sulfate displace calcium nitrite? Calcium nitrite is widely available and cheap as chips:)

Picric-A - 2-11-2009 at 06:38

Calcium nitrite? where can one obtain that 'as cheap as chips'?

DJF90 - 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.

Sedit - 2-11-2009 at 07:22

Calcium nitrite 30% solution to inhibit steel rebar corrosion in concrete.

Attachment: eucon_bcn.pdf (135kB)
This file has been downloaded 1345 times

Taoiseach - 12-11-2009 at 02:45

Some interesting tidbits from "Chemistry - inorganic and organic" by Charles Loudon Bloxam (1883 lol :D):

-solutions of nitrites are readily oxidized back to nitrate by heating in contact with air.

-sodium nitrate is reduced by fusion with iron fillings, the main product being sodium hyponitrite which can be obtained as needle-shaped crystals



no.JPG - 26kB

roamingnome - 12-11-2009 at 22:03

I was reading about the iron filling bit a few days ago,
but just looking at wiki for basic information
about Hyponitrite one finds ...
can act as a reducing agent for example reducing iodine, I2,:[2] N2O22− + 3I2 + 3H2O = [NO3]− + [NO2]− + 6HI

that seems like a slick way to HI

dann2 - 30-12-2009 at 04:41

Hello Folks,

Just wondering if I use roofing Lead + Na Nitrate for to make Na Nitrite, will the Tin (I don't know if there is/is not any) in the roofing Lead interfere with the reaction or what products may I expect.
Thanks, Dann2

DJF90 - 30-12-2009 at 10:27

Roofing lead should be pretty pure lead. Beware the "Adhesive flashing strips" though, as these are mostly bitumin with an aluminum foil; not lead as I'm sure you could tell from that description. The brand name of these escapes me at this present moment in time.

Taoiseach - 1-1-2010 at 09:36

Tin/antimony is no problem at all as they get oxidized to white insoluble crap that is easy to remove by filtration.

dann2 - 5-1-2010 at 19:40

Hello,

Attached patent describing seperation of Sodium Nitrite from Sodium Chloride. May be useful.

Dann2

Attachment: US3965247.pdf (284kB)
This file has been downloaded 952 times

Nitrate + Charcoal + Starch = Nitrite

dann2 - 9-1-2010 at 15:23

Hello,

Link below states that you can make Sodium Nitrite from Nitrate, Charcoal and Starch.

http://chestofbooks.com/health/materia-medica-drugs/Manual-P...

Anyone tried this? I don't see it mentioned in the thread.

The Sodium Nitrite that I made from Nitrate heated with molten Lead appears to be only approx. 10 to 20% Nitrite. It does not really matter to me actually, just saying.
I measured its 'purity' using a Nitrite detector kit for aquariums.
I had not realized you are supposed to fuse (melt) the Sodium Nitrate when mixing/stirring with the melted Lead.
Shall have to fire up the cauldron again........

Will Methylene blue detect Nitrite?

Cheers,

Dann2

Taoiseach - 17-1-2010 at 13:04

Another idea for the preparation of nitrites:

Calcium formate, Ca(HCOO)2, could be an excellent reducing agent for nitrates:

Ca(HCOO)2 ---> Ca(2+) + 2CO2 + 2H(+) + 4e(-)
2[NO3(-) + 2e(-) ---> NO2(-) + O(2-)]

thus:

2NaNO3 + Ca(HCOO)2 ---> 2NaNO2 + H2O + CO2 + CaCO3

the only products being calcium carbonate, sodium nitrite and carbon dioxide. The latter is a gas and calcium carbonate is one of the least soluble calcium compounds (much less soluble than calcium sulfate even) - hence separation of the nitrite should be a breeze :)

I dont know if this reaction proceeds with reasonable speed in aequous solution but it'd be worth a try! Calcium format is quite soluble (6.6 g/100 mL @20°C) and the slightest amount of carbonate produced would cause turbidity, thus showing that the reaction proceeds.

Under more vigorous conditions the reaction would certainly proceed, as formates make quite energetic fuels in oxidizer-fuel kind of mixtures. Maybe by carefully adding Ca(HCOO)2 to molten nitrate in small(!) portions the temperature could be controlled not to exceed the decomposition point of NaNO2 (271 °C).

Taoiseach - 22-1-2010 at 09:58

Looks like I'm talking to myself again :)

Reduction of nitrate to nitrite by formate is confirmed by attached paper. According to this, melted nitrate reacts vigorously with formate, evolving carbon dioxide and turning from colorless to yellow. Apart from the reaction I suggested above, another competing reaction is reported to take place. In case of calcium formate:

Ca(COOH)2 + 2NaNO2 ---> CaCO3 + Na2CO3 + N2O + H2O

This reaction would contaminate the nitrite with carbonate but the paper suggests its neglible.

Another problem mentioned is that the fused mass tends to suddenly explode :o

Adding Ca(OH)2 to dilute the mixture as well as to precipate any carbonate formed might remedy the situation.

No useful information could be found on the reduction potential of formate in aequous solution. There's a little chance it'd work tough. At least sodium chlorate is known to be able to oxidize carbon in hot solution.

Attachment: Reactions between sodium carboxylic acid salts and molten sodium nitrate and sodium nitrite.pdf (925kB)
This file has been downloaded 1245 times

Taoiseach - 1-2-2010 at 01:27

"Absorption of nitrous gases" (p.315), available from the library section, confirms the calcium formate route, stating that it quantitatively reduces sodium nitrate to nitrite. Details are in French pat 388,563 which I was unable to find tough.





ca.JPG - 36kB

Jimmymajesty - 6-3-2010 at 03:53

It has not been mentioned before IIRC so I decided to post this method, as it seems the easiest method to make pure N2O3 to me.

Excerpt from brauer preparative inorganic chemistry page:488

NOHSO4 made by generating SO2 by potassium metabisulphite (OTC) + HCl and lead it into ccHNO3

SO2+HNO3=>NOHSO4

Then add water dropwise into the slurry to form N2O3 in a slow way.

I am not sure if it works, but will definitely try. :)

Nicodem - 6-3-2010 at 04:44

Quote: Originally posted by Taoiseach  
Details are in French pat 388,563 which I was unable to find tough.

You probably mistyped it or something as it is right where it is supposed to be: FR388563 (also available in German version as DE203751)

JohnWW - 6-3-2010 at 20:18

Quote: Originally posted by Nicodem  
Quote: Originally posted by Taoiseach  
Details are in French pat 388,563 which I was unable to find tough.

You probably mistyped it or something as it is right where it is supposed to be: FR388563 (also available in German version as DE203751)

When one tries to save PDFs of those French and German patents of 1908, v3.espace.com throws up popup windows that immediately cause Firefox to crash. Has anyone been able to download them, somehow? If so, please post them here are attached files.

Jimmymajesty - 13-3-2010 at 06:15

I tried out the method for N2O3 generation by means of NOHSO4.

Unfortunately the SO2 reacts very slowly with icecold HNO3, and dry SO2 generation by K2S2O5 is a pain...

While bubbled SO2 through HNO3 I also made some K2S2O5 solution, it had a smell of SO2, so I poured diluted HNO3 on it and had a lot of (I presume) N2O3 generated which dissociated at the heat of the reaction to NO and NO2..

Next time I will put K2S2O5 in a gas generator and drip HNO3 it and lead the gases into NaOH sol.

Would anyone be so kind to back me up that the following equation is correct: (I did not bother balancing it)

K2S2O5+HNO3+H2O=>K2SO4+HNO2+H2SO4

Thanks a lot!

starman - 13-3-2010 at 18:24

Quote: Originally posted by Sedit  
Calcium nitrite 30% solution to inhibit steel rebar corrosion in concrete.


Jimmy I don't know your intended use of N2O3 or whether you're just interested in the chemistry of the route you're looking at.
Calcium nitrite solution quoted,available fairly cheap commercially(a hardware stare ordered in for me) plus dilute HCl worked very well for my purposes.You get that lovely nitrous acid blue immediately on addition.

[Edited on 14-3-2010 by starman]

Jimmymajesty - 14-3-2010 at 09:32

Actually I don't know what to do with it at the moment but it is an interesting chemical, so I decided to make some. The calcium nitrite is not available in my country, maybe only for some industrial bastards, sulfuric acid based drain opener also unknown.

The problem is I don't want to use lead as it is poisonous and also I have not succed in any thermal method to make it eg:sulphites+nitrates, sugar+nitrates.

I have only made NaNO2 by starch + HNO3 so far, but a lot of NaOH ended up as Na2CO3 as a result of the evolving CO2.

I carried out a test yesterday as follows:

200g K2S2O5 was put in a flask and HNO3(85°%) were being dripped to it until no gas evolution was observed.
The problem was that owing to the heat generated by the oxidation, the formed SO2 and HNO3(gas) reacted in the condenser forming a solid mass possibly NOHSO4 and something else that color was deep green. (I also get that green color when added H2O to the NOHSO4+HNO3, I don't know offhand what is this material.. I am too tired to look it up :P)

I am going to repeat the reaction but, saturated K2S2O5 solution will be dripped into ~50% HNO3.

starman - 14-3-2010 at 17:53

Jimmy you seem to be making more complicated than necessary.You know that dilute HNO3 + Fe will do the job right? Sure your NO/NO2 ratios will change later on slowly evolving excess NO in the end,but if its just to look and see,whats it matter?

bluetrain - 20-3-2010 at 08:11

I made sodium nitrate from ammonium nitrate and sodium hydroxide. A solution of the salt had a pH of ~5-6 which I guess means there still is some ammonium ions in there. Would this be unsafe to melt? It will decompose at around 210'C into oxides of nitrogen and water so if I use a large enough container I think I'll be safe. Impure ammonium nitrate can however explode when heated so I'm not quite sure.

Also, does anyone know if airgun bullets are pure lead or if there is added some other metal to make them harder? The lead I can source is from lead batteries or fishing weights with the latter being the easiest to acquire. The stores are closed until Monday so if I want some nitrite this weekend I'll have to use airgun bullets.

Anders Hoveland - 5-7-2010 at 23:46

Would NaNO3 mixed with HCl solution, and then adding aluminum foil, make nitrogen dioxide and nitric oxide?

I have see somewhere that Zinc with HCl can reduce nitrobenzene to analine so I would think it could reduce a nitrate. The nitrite would immediately react with HCl, giving off nitrogen oxides.

NaNO2 + 2HCl --> 2NaCl + H2O + NO + NO2

The NO and NO2 could be bubbled into baking soda solution to make sodium nitrite.

Polverone - 6-7-2010 at 08:48

Quote: Originally posted by Anders Hoveland  
Would NaNO3 mixed with HCl solution, and then adding aluminum foil, make nitrogen dioxide and nitric oxide?

I have see somewhere that Zinc with HCl can reduce nitrobenzene to analine so I would think it could reduce a nitrate. The nitrite would immediately react with HCl, giving off nitrogen oxides.

NaNO2 + 2HCl --> 2NaCl + H2O + NO + NO2

The NO and NO2 could be bubbled into baking soda solution to make sodium nitrite.


Why don't you try it? None of the needed materials are rare or expensive. I know that a mixture of hydrochloric acid and nitrate salts gives red nitrous fumes with zinc.

turd - 6-7-2010 at 10:00

This thread shows that there are many ways to obtain nitrites. The challenge is to make *pure* nitrite. So far this thread didn't provide convincing results (I might have missed something though).
The proposed N2O3 method does not seem to be advantageous:
http://www.erowid.org/archive/rhodium/chemistry/pseudonitros...
(Read the part about generating N2O3 from metals)

Sedit - 6-7-2010 at 13:04

Whats so hard and impure about the molten Lead method? if you have concerns of Pb contamination there have been suggestions of feeding C2 into the solution and filtering off Pb carbonate other then that I find the Pb method easy and top not but a tad tedious due to alot of filtering and waiting for the cake to dissolve.

turd - 6-7-2010 at 14:14

Like I posted upthread: I tried with different metals and never got anything close to pure NaNO2. There was always more or less (typically ~50%) NaNO3 and other impurities. Particularly, I found the Pb method hard to control with either incomplete reaction or runaways. Maybe I'm just too clumsy for this kind of chemistry, but it certainly isn't *easy* to get good yields with molten NaNO3/metal. Maybe one day I will try in vacuum and with an excess of a metal that doesn't melt. But it doesn't have a hight priority since I managed to acquired some NaNO2 p.A..

Sedit - 6-7-2010 at 14:35

Its very easy esp when reducing NaNO3. KNO3 is abit harder.

Take a LARGE amount of lead and melt it. Once its melted completely start adding gram quantities of nitrate and stir. Keep adding till complete addition and the stir as much as you can giving ample time for the reaction to finish. KNO3 is a bit harder because it likes to clump up with the Litharge but all thats needed is extra heat.

Once you feel its done stop stirring cover loosely and let it set so that your nitrite and litharge float to the top. THEN turn off the heat

Cover and allow to cool.(This I feel is many peoples downfall is they allow reoxidation).

Add excess cold water and cover again and allow the cake to break up. The nitrite is highly soluble in H2O. Filter and filter again to remove litharge and slowly evaporate or do so rapidly under inert atmosphere or positive pressre of steam. When its concentrated Allow it to cool and placeit in the freezer where any unreacted nitrate will preciptate out. Filter and your remaining liquid is a concentrated Nitrite solution of purity I hav never seen a problem with. The only issue I hae ever had is keepin the nitrite dry.

Its not hard at all just a process that sounds more complicated on paper then it really is.

Taoiseach - 7-7-2010 at 01:14

The lead reduction method is useless crap IMHO. I once read in an old chemistry book that several runs are required to obtain nitrite of acceptable purity. That is, melt nitrate with lead, extract the fused mass, remove nitrate by crystallization, evaporate, melt with lead again... repeat at least three times. This might be acceptable on a industrial scale but not for the amateur experimenter. I have tried this method twice now and each time there was a HUGE amount of nitrate crystallizing from the liquid upon cooling. Nitrite has such a high solubilty that as a rule of thumb, if anything crystallizes from the solution then the reduction was not even close to completion. I never cared to dry and weight the remaining nitrate but I roughly estimate that at least 50% of the nitrate remained unreacted.
The lead/nitrate melt was stirred excessively in the beginning but the mixture quickly solidified and made stirring impossible.
As a sidenote, reduction of thiocyanate with several metals yields cyanide, and this was once proposed by Erlenmeyer as a industrial process. It was noted that lead performs rather shitty compared to other metals due to its low melting point. The blobs of molten metal has a much smaller surface and intimate mixing is not possible to the same extent as with powdered zinc or iron. I believe the same effect comes into play when nitrate is reduced with lead. If higher yields and better purity is the goal, then we need a procedure where the reactants can be mixed prior to melting.

I suggest doing some research into alternative reducing agents such as carboxylic acid salts, iron powder or copper(I)oxide. Dont waste your time with the lead reduction. Using powdered lead might remedy the situation somewhat but lead shot/plumbing lead is definetly a dead end.
A promising experiment using calcium formate was described here:

http://www.versuchschemie.de/topic,13865,-Herstellung+von+Ni...

The reaction is vigorous but not explosive. Intimate mixing of the reactants is required to obtain a good yield. Using Ba/Ca formate/oxalate makes sure that oxidation products are insoluble and nitrite of high purity is obtained by simple filtration.

Iron formate/iron oxalate are also worth a try.

Powdered iron also reduces nitrate, producing some ferrate which can be decomposed into insoluble iron hydroxides by boiling.

[Edited on 7-7-2010 by Taoiseach]

Anders Hoveland - 7-7-2010 at 10:19

I did some more experimenting. A solution containing dissolved KNO3 and about 5-10% HCl was made and aluminum foil added. The smell of nitrogen dioxide became obvious immediately above the solution, but a brown color was not vissible. A variation was tried with more concentrated HCl solution and a brownish color developed in the solution and bubbles appeared with a strong burning smell, but still no brown gas. Next, solid HH4NO3, as much as would dissolve, was added into HCl (5-10%) solution. A small strip of Foil was added, and this time some brown gas developed over the solution, but more excessive generation of NO2 remained elusive. Hydrogen could still be ignited above this last solution, indicating that
2Al + 6 HCl --> AlCl3 (hydrate) + 3H2 was still the primary reaction

I think Nitrogen dioxide oxidizes hydrogen at ambient temperatures, so this probably limits much NO2 from developing.
Nitric oxide might be safe from reaction with hydrogen. I am unsure how to get rid of the hydrogen though. Nitric oxide would react with oxygen in water to form nitric acid, so using oxygen to oxidize all the hydrogen (through the nitric oxide that acts like a catalyst) would not work.
Perhaps by using dry chlorine (passing it through baked/powdered CaCl2), the hydrogen will be oxidized to HCl gas and nitrosyl chloride will be left over. If excess chlorine is not used, and the gas allowed to sit for a minute, so that all the chlorine is able to be reduced, this could then be bubbled into pure anhydrous acetone to possibly obtain nitroso-acetone ONCH2COCH3. Or it could be bubbled into a Na2CO3 solution to obtain NaNO2, NaCl also forming. These are just ideas.

If I could get the generation of hydrogen to be more diffuse, this might prevent any HNO2 or NO2 that forms from immediately getting reduced. This would require aqueous metal ions that would react with acid to give off hydrogen, as the reaction on solid metal is probably too concentrated.
Perhaps Cu+ and H2S would form CuS , H+ and 1/2 H2, driven by precipitation of insoluble CuS ?

I remember reading that reaction of dilute HNO3 on Cu, actually produced traces of NH4NO3.

[Edited on 7-7-2010 by Anders Hoveland]

Sedit - 7-7-2010 at 10:49

""The lead/nitrate melt was stirred excessively in the beginning but the mixture quickly solidified and made stirring impossible.""""

Then your issue is too low of a heat source. You must make sure that your Nitrate/Nitrite is melted as well else little is going to happen. Keep in mind that you ARE dealing with a duel phase reaction here and without rapid stirring and making sure to add small portions at a time you get Nitrite and Nitrate encased in Litharge.

Keep in mind these melting points and see why NaNO3 is so much easier to reduce the KNO3. It sounds to me that you tried to reduce KNO3 without using enough heat. NaNO3 proceed pretty rapidly and with ease but its harder to work up so its a trade off of sorts. Lead dont boil till 1749 °C so heat that shit up as high as you can get it and it will go along just fine.

327.46°C Pb
334°C KNO3
308°C NaNO3

440.02°C KNO2
271°C NaNO2

I honestly feel if someone can't get the Pb method to work its because there just not doing it right.
Quote:
I believe the same effect comes into play when nitrate is reduced with lead. If higher yields and better purity is the goal, then we need a procedure where the reactants can be mixed prior to melting.


Try ballmilling Nitrate with Pb shot prior to the melt then. I performed an experiment a while back because I read an abstract that mentioned the reduction performed this way successfully but I never received the full paper so I went ahead and added excess Pb shot to damped nitrate and Course silica in ballmill and let it run for a couple days. I added the silica in hopes that it will expose more surface area but I never quantified the results because they seemed to show a marked reduction in NOx production on addition to H2SO4. I still think that route warrents attention but I have not gotten around to experimenting with it more and completing it with a melt could be exactly what it needs. It was to messy for me to workup hence the reason I never really attacked it from all angles.

@anders

I will have to look at my notes but I did a few simular experiments a while back IIRC and if you use NaOH/Al NH3 is produced in abundence so there is a good chance your not getting your desired results because you are reducing the nitrate all the way to amine and locking it away as Ammonium chloride.

[Edited on 7-7-2010 by Sedit]

Taoiseach - 8-7-2010 at 14:07

@Sedit

>Then your issue is too low of a heat source. You must make sure that your Nitrate/Nitrite is melted as well else little is going to happen. Keep in mind
>that you ARE dealing with a duel phase reaction here and without rapid stirring and making sure to add small portions at a time you get Nitrite and
>Nitrate encased in Litharge.

The oxidation product of lead is PbO and its not going to melt at any reasonable temperature. The mixture will solidify towards the end of the reaction and there's nothing you can do about it - except using a huge excess of nitrate.

>Lead dont boil till 1749 °C so heat that shit up as high as you can get it and it will go along just fine.

No it wont. The nitrite will decompose and leave you with useless sodium oxides/hydroxide.

>I honestly feel if someone can't get the Pb method to work its because there just not doing it right.

Nobody has done any sort of analysis/estimation of purity so far. I believe that what has been produced so far by this method is nothing but a bunch of nitrate with *some* nitrite in it. Adding HCl says almost nothing about purity, as even very small amounts of nitrite will produce red fumes.

>Try ballmilling Nitrate with Pb shot prior to the melt then. I performed an experiment a while back because I read an abstract that mentioned the
>reduction performed this way successfully but I never received the full paper so I went ahead and added excess Pb shot to damped nitrate and
>Course silica in ballmill and let it run for a couple days. I added the silica in hopes that it will expose more surface area but I never quantified the
>results because they seemed to show a marked reduction in NOx production on addition to H2SO4. I still think that route warrents attention but I
>have not gotten around to experimenting with it more and completing it with a melt could be exactly what it needs. It was to messy for me to workup
>hence the reason I never really attacked it from all angles.

This is a very interesting method and much more promising than the high-temperature reduction.

Sedit - 8-7-2010 at 15:11

The solidity of NaNO3 melt is only semi riggid at completion and small portions on addition take care of that so even though there is a lumpy mass at the end much of it is nitrite. Honestly when it comes to the reduction of KNO3 however I really do not like this method at all. I have to have yeilds somewhere around either here or The Vespiary for the Pb reduction. If I never posted them or can't find them i'll just have to perform the reduction again. I could use some nitrite anyway as I want to replenish depleting supplys of chemicals. What about adding SiO2 to the lead prior to the addition to help breakup the melt?

The ball milling is just not up to scale for me due to the small ballmill I have so its for nothing more then an experimental. If there was a way to ballmill and melt at the same time it would get my seal of approval over the range of nitrate cations.
Quote:
No it wont. The nitrite will decompose and leave you with useless sodium oxides/hydroxide
.
Ihave allowed the melt to sustain melted for extended period of time but I have never see a marked decrease although it would be interesting to find a means to yeild pure hydoxide to see how much decomposition takes place. The Na should decompouse in the mix as its being oxidised but it does not appear to do so.

Adding HCl is not a great sign of nitrite and H2SO4 seems to provide a better color from the mixture when dripped in.

As a side note a side product of the oxidation iv been messing around with using HNO3 and EtOH appears in all shapes and formes to be a mixture of nitros acid and acetic acid it takes on the distinct blue color of nitros acid but smells strongly of AcOH. Not quite GAA status but not really wet either. Its honestly quite a problem for me at the moment. Large amounts of NOx is produced in the first stage of the oxidation.

Formatik - 8-7-2010 at 23:36

Quote: Originally posted by Taoiseach  
I suggest doing some research into alternative reducing agents such as carboxylic acid salts, iron powder or copper(I)oxide.


Be careful what reductant you use, sodium acetate and potassium nitrate will explode violently when melted together according to the Ber. article posted in the cyanides thread on the nitrite method of NaCN. The same reference noted a delayed explosion in a melt consisting of a molecular mixture of sodium formate and NaNO2.

There is also some old method on forming it by heating nitrate alone: the KNO3 is glowed strongly, then the mass is solubilized in hot water. After 24 hours, the liquid is poured off from precipitated KNO3. Then neutralized with dilute acetic acid, and mixed with the double volume of ethanol. After several hours, three layers form, the middle layer contains a yellowish oily liquid, from which the KNO2 is obtained by evaporating over H2SO4, N.W. Fischer (Pogg. Ann. 150 [1848] 116).

Taoiseach - 9-7-2010 at 03:33

>There is also some old method on forming it by heating
>nitrate alone: the KNO3 is glowed strongly, then the mass
>is solubilized in
>hot water. After 24 hours, the liquid is poured off from
>precipitated KNO3. Then neutralized with dilute acetic acid,
>and mixed with
>the double volume of ethanol. After several hours, three
>layers form, the middle layer contains a yellowish oily liquid,
>from which
>the KNO2 is obtained by evaporating over H2SO4, N.W.
>Fischer (Pogg. Ann. 150 [1848] 116).

KNO3/NaNO3 doesn't decompose cleanly; there is an equilibrium formed between nitrate decomposing into nitrite and nitrite decomposing into various oxides. That must be the reason they add acetic acid - it converts the oxides/hydroxides into acetate which is decently soluble in EtOH whereas nitrite is not.

>Be careful what reductant you use, sodium acetate and
>potassium nitrate will explode violently when melted
>together according to the >Ber. article posted in the
>cyanides thread on the nitrite method of NaCN. The same
>reference noted a delayed explosion in a melt
>consisting of a molecular mixture of sodium formate and
>NaNO2.

I recently had a go at the Ca formate reduction and its not explosive tough exothermic and a bit violent. I tried again with an equal volume of Fe dust added. It made the reaction more controllable and also improved thermal conduction. When the reaction set it the whole pot started glowing dark red for a few seconds; lots of fumes were given off and some hissing but no flame. The mass was left to cool, extracted with hot water, filtered to remove CaCO3 and precipated with EtOH. The nitrite formed a heavy yellow oil with EtOH; this was evaporated to dryness. Btw cyanides also form this "oil" (probably some sorta gel) with EtOH.
Yield is poor tough because the solubility of nitrites and cyanides in EtOH increases drastically with the amount of water present.

>The ball milling is just not up to scale for me due to the
>small ballmill I have so its for nothing more then an
>experimental. If
>there was a way to ballmill and melt at the same time it
>would get my seal of approval over the range of nitrate
>cations.

Lead sponge can be precipated from lead acetate by adding zinc metal strips or zinc powder. After torough washing and drying it should be possible to intimately mix it with the nitrate prior to melting.





[Edited on 9-7-2010 by Taoiseach]

Formatik - 9-7-2010 at 08:50

Quote: Originally posted by Taoiseach  
KNO3/NaNO3 doesn't decompose cleanly; there is an equilibrium formed between nitrate decomposing into nitrite and nitrite decomposing into various oxides. That must be the reason they add acetic acid - it converts the oxides/hydroxides into acetate which is decently soluble in EtOH whereas nitrite is not.


Sodium and potassium nitrate are very few of the alkali compounds that readily yield their oxides on thermal decomposition. It would probably also be a good idea not to do this in glass or platinum (who uses it anyway?).

Quote:
Lead sponge can be precipated from lead acetate by adding zinc metal strips or zinc powder. After torough washing and drying it should be possible to intimately mix it with the nitrate prior to melting.


Concerning a method where spongy Cu-powder made from CuSO4-soln. and zinc dust, is melted together with KNO3 is described below. Then there is also something by O.L. Erdmann (J. pr. Ch. 97 [1866] 387 footnote) where KNO3 with an excess mass of iron filings is molten in a cast iron crucible and heated to moderate glow.

Attachment: J. chem. Soc. Lond. 36, 595.pdf (176kB)
This file has been downloaded 820 times

Jimmymajesty - 21-7-2010 at 11:03

Hi folks!

I looked through the thread a couple of months ago, and based on the comments I had not even tried the molten lead method.. It seemed so messy and troublesome.. I like to reproduce hard to reproduce processes @ home:) but grinding lead, filtering lead, melting lead, three things that I would avoid at any cost, I just stay away from anything that involves lead, and I suggest this attitude to all of you!

Nitrosyl sulphuric acid(NOHSO4, chamber crystals) replaces NaNO2 in most of the cases, so make that instead of XNO2, It also stores well in a well stoppered glass bottle, and relatively easy to make, (easier than any XNO2 shit)

Take an aspirator and connect it to the end of a setup that consists of a heated quartz tube, in which you burn sulphur, the quartz tube is connected to a bubbler full of H2SO4 the escaping SO2 from the bubbler is led to cc HNO3, it takes some time and the reaction mixture requires cooling but NOHSO4 is not that hard to make afterall. After about a hour yelow cristals starts to precipitate if the solution is properly cooled. One drop of the resulting yellow liquid makes a shitload of N2O3 if it is added to water.

I found a book on www.archive.org search for: The Aromatic Diazo-compounds And Their Technical Applications, its a nice book with preparative examples, including the use of nitrosyl sulphuric acid.

Sedit - 21-7-2010 at 11:52

Thats is interesting Jimmy no doubt about it deserves more attention.

For some people Bisulfite's might be a better means at SO2 production then a sulfur burner even though in the long run sulfur would be cheeper. I assume the conentrated H2SO4 is a means of drying the SO2 but that should pose no problem at all.

The concentrated HNO3 could be made with H2SO4 and KNO3 then frozen and cold filtered to remove as much KHSO4 as possible so I see much potential for this being highly over the counter.

Its there any risk involved if the nitric acid used is red from NOx contamination?


[Edit]
This statement in a patent on the synthesis of nitrosyl chloride seems to contridict your statement that addition of the nitrosylsulphuric acid to water generates the desired gas.
Quote:

Solutions containing more than 40% of nitrosylsulphuric acid crystallize when the water content exceeds 10%. To overcome this disadvantage, it is then necessary either to reduce the content of nitrosylsulphuric acid, which is reflected by a decrease in the productivity, or to heat the solution to a temperature of the order of 50 to 100° C., which is expensive.

The presence of water causes, in the more or less long term, hydrolysis of the nitrosylsulphuric acid, which correspondingly decreases the yield of nitrosyl chloride. Hydrolysis becomes more significant as the temperature rises.


However even though it seems to suggest water will crystalize the nitrosylsulphuric acid it seems that it does so slowly and does not 100% hinder the formation if this patent is to be trusted.

Reference: http://www.freepatentsonline.com/6238638.html



What about reacting KNO3 in excess H2SO4 and then chilling the solution significantly and slowly adding Sodium Metabisulfite to generate SO2 insitu.

H2SO4 + Na2S2O5 -->> 2 SO2 + Na2SO4 + H2O

H2SO4 + KNO3 --> KHSO4 + HNO3

2HNO3 + 3SO2 + 2H2O --> H2SO4 +2NO


Yes I know the Stoichiometry leaves alot to be desired at this point but im just museing here so bear with me. I dont know much about this compound right now and am speaking as I learn.

[Edited on 21-7-2010 by Sedit]

Sedit - 21-7-2010 at 18:15

My scale has had abit of a meltdown due to a H2SO4 accident a few weeks back so it will be a little bit before I can give anything conclusive but I thought I would share this.

I took KNO3 and dissolved it in a test tube of 97% H2SO4 making sure to have excess H2SO4 there. I then c chilled this in an Ice/NaCl bath to lower the temperature down as much as I could with ease. There was a slight reaction as the addition was slowly made and at this point is became a thick syrupy consistancy. Unaware if anything at all had happened I let it sit in the ice back for a few hours as I went about my own things. I came back to a clear'ish solution with needles of more then likely bisulfites crystalising out. I decided to drip H2O in at which point large amounts of light yellow gas as formed and quickly oxidised to dark NO2.

This looks very promising to me but alot more test need to be done to determine if the gas was indeed N2O3.

[EDIT]

In the morning there is a thin layer of yellow liquid, resting on top o all the precipitated salts that where in the reaction, that resembles all photos I have seen of nitrosylsulphuric acid. Addition of H2O indeed generates a large amount of gas production that by all means appears to be N2O3.

Is there any solvent I could use to isolate this compound from the salts to run better, purer test with it?

[Edited on 22-7-2010 by Sedit]

Jimmymajesty - 22-7-2010 at 08:27

SO2 generation by means of K2S2O5 is not that good idea IMHO, if you drip ccHCl to solid K2S2O5 then the KCl cristallizes on the surface of the K2S2O5 crystals because of the endotherm reaction, you can heat the flask of course but a cake will form and it will prevent agitation, magnetic stir bar is useless here. If you dissolve the K2S2O5 in water and drip cc HCl into that, the SO2 will remain in the water, you have to heat the solution to push out the SO2, but it means more H2O vapour in the resulting gas, and your bubbler with H2SO4 will quickly become uneffective. So burn sulphur with an aspirator, I also tried to burn sulphur in a direct air flow, by means of an aquarium pump, but is also poses many problems, aspirator is the way to go!

You have to use dry SO2 here, or you will continously hydrolize the NOHSO4 formed during its production.
I do not have experience with H2O+NOHSO4 system, but I noticed that, in the book the H2SO4 was always used in large excess to the NaNO2, this is probably due to the water formed in the reaction, which has to be 'absorbed' by the excess of H2SO4!

It should be noted however when NOHSO4 added to water it makes a very exotherm and vigorous reaction. I made some N2O3 volcano by injecting 1ml NOHSO4 in H2SO4 to water, I bet If you could see that you would hardly believe that NOHSO4+H2O mix even exists in any proportion(it also new to me:).

In brauer preparative inorg. chem fuming nitric is used, and I also used fuming nitric without any problem.

You probably start with 95% HNO3, in the initial stage of the reaction the NOHSO4 formed will react with the 5% water present to give H2SO4 and N2O3, as the reaction proceeds the H2O will be consumed, and H2SO4 wont form anymore till the end of the reaction, depeding on the moisture content of your SO2 of course. You can follow the temp. of the reaction mixture by the colour, if it becomes intense orange you have to cool it till it go back to yellow.

I do not really know how could you isolate the NOHSO4 from H2SO4+NaNO2. Prepare relatively conentrated NOHSO4 from fuming HNO3+SO2 and filtrate the crystals on a glass frit, but you do not have to isolate it to use it.

Sedit - 22-7-2010 at 09:21

Dripping H2SO4 onto slightly damp Metabisulfite is working wounders for generating SO2 with ease. Theres good odds that this will also remove a large portion of the water from the reaction.

Iv been running small scale test tubes since you posted this and im pretty confident this can be done with Bisulfite as the SO2 source and quite possible insitu. The substance that appeared over night was light golden liquid and fumed vigerously with water generating a blue solution if only a small amount of water was used. A light yellow gas was formed but on heating the solution formed dense brown NOx fumes.

When given a chance I want to run it the right way as a control to be sure further experiments are what I am looking at.

[Edited on 22-7-2010 by Sedit]

Jimmymajesty - 23-7-2010 at 10:12

Sedit Yes the bisulphite method works, it is in the book:)

Today I made isopropyl nitrite by NOHSO4 and it worked well.

I took about 100ml izopropyl alcohol and diluted it with 20ml water, I chilled the mix to 5°C started to agitate while slowly dripped the NOHSO4+H2SO4 into the mixture (I dunno what concentration, I sucked it up from the white NOHSO4 crystals) every drop made a hissing noise but no nitrogen oxides formed, all the gasses absorbed by the alcohol.
After the addition of about 10ml NOHSO4 solution I poured the whole mixture into cold water, and a yellow layer separated, as usually.. I smelled it got headache etc. so the method works, but needs to be optimised ofc.

cnidocyte - 12-2-2011 at 00:28

Quote: Originally posted by kmno4  

NaNO2 content is easy to determinate using KMnO4/H2SO4.
(add slowly [with stirring], NaNO2 solution to KMnO4/H2SO4 sol. untill it becomes colourless)

I've been looking into decent spot tests to differentiate between nitrites and nitrates. How does this one work? Does the MnO4- ion oxidise nitrites but not nitrates?

cnidocyte - 13-2-2011 at 08:27

Quote: Originally posted by Polverone  
I tested a slightly modified version of Muspratt's nitrite preparation method this afternoon. 100 g of KNO3 were mixed with 12.1 g of charcoal and the whole thing ball milled for about an hour. I tested a little bit of the mix and found that it would burn without an external heat source. I poured the mix into a stainless steel dish and ignited it.

The reaction wasn't terribly fast, due to the great excess of oxidizer, but it was fairly vigorous. There was a lot of bubbling and splashing of the molten salt since my dish was barely large enough to hold the charge of powder. When burning finished I placed a sheet of copper over the top of the dish and waited for it to cool.

I had different results. 10g KNO3 was grinded and mixed with 1.2g activated charcoal with a mortar and pestle until a fine gray powder was obtained. I tried igniting the mixture with a lighter but all that happen was the powder being flamed crackled and turned black. I had to blowtorch the mix and keep heat on it to keep the reaction going. Balls of white molten salt rolled around on the mix and ultimately I ended up with a dirty gray coloured solid clump. This is the same colour that the KNO3 turned when I melted it by itself in a boiling tube. This is even after I recrystallised it so I don't really know whats going on.

entropy51 - 13-2-2011 at 09:12

What was the source of the KNO3? How did you recrystallize it?

(I guess I'm wondering if it really is KNO3. People have used stump removers that were not KNO3, thinking that they were KNO3.)

It sounds as if you ground KNO3 and charcoal together in a mortar. Please tell me you didn't do that.

[Edited on 13-2-2011 by entropy51]

cnidocyte - 13-2-2011 at 09:44

Quote: Originally posted by entropy51  
What was the source of the KNO3? How did you recrystallize it?

It sounds as if you ground KNO3 and charcoal together in a mortar. Please tell me you didn't do that.

My friend gave me about 20g of KNO3 that he got on ebay. I recrystallised by adding boiling water until the KNO3 was fully dissolved then cooled to around 5C and collected the crystals.

I did ground the nitrate with the charcoal in the mortar. It was outdoors and I was ready to deal with it if it ignited. This isn't very different to ball milling is it?

Rosco Bodine - 26-8-2011 at 04:32

Quote: Originally posted by JohnWW  
Quote: Originally posted by Nicodem  
Quote: Originally posted by Taoiseach  
Details are in French pat 388,563 which I was unable to find tough.

You probably mistyped it or something as it is right where it is supposed to be: FR388563 (also available in German version as DE203751)

When one tries to save PDFs of those French and German patents of 1908, v3.espace.com throws up popup windows that immediately cause Firefox to crash. Has anyone been able to download them, somehow? If so, please post them here are attached files.


Here are the patents attached FR388563 and DE203751

Check the archive for the old E&W forum as there were numerous additional patent processes listed there.
I'll try to dig those up again and repost them here.

Attachment: DE203751 Nitrite from Nitrate and Formate.pdf (59kB)
This file has been downloaded 776 times

Attachment: FR388563 Nitrite from Nitrate and Formate.pdf (71kB)
This file has been downloaded 708 times

terraxus - 30-8-2011 at 10:58

why not just heat NaNO3?

Alastair - 7-10-2011 at 12:33

Quote:
Quote: Originally posted by madcedar  


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




Why not just react N2O3 with NaOH directly?
And btw this method isnt bad for any impurities, you can purify gas from excess NO or NO2 by putting it into your freezer (one becomes solid, the other stays gas while N2O3 remains in liquid form)

rstar - 8-10-2011 at 09:44

Can reaction of (molten) Potassium nitrate with Zinc make Potassium nitrite ??

If yes, then i think it would be like :

KNO3 + Zn = KNO2 + ZnO :D

ZnO is insoluble in water, while nitrite is soluble

KNO3->[RED]->KNO2 effective reductions

niertap - 16-4-2012 at 18:41

Has anyone tried the nitrate to nitrite reduction using iron filings? I've heard heating nitrates with lead works well, but heating them with iron filings seems like a better and quicker method. Would it work? If not what would be formed?

[Edited on 4-17-2012 by Polverone]

chochu3 - 16-4-2012 at 21:07

There are no experimental details of using iron, hence asking if anyone has tried and not posted here about it. Any way tried the reduction with lead you do have to stir frequently and is superior because of lower mp than if you were to use iron. But if iron is all you have go for it. Post back whatever findings you have, would love to hear the outcome.

[Edited on 4-17-2012 by Polverone]

niertap - 16-4-2012 at 21:27

I saw a little bit about it in a thread about the creation of K2FeO4, but nothing concrete.

I've been thinking about that and from user experiences it seemed like K2ferrate, iron oxides, and or nitrates could form. I'm assuming higher heats cause two nitrites to react with an atom of iron to form the ferrate; lower heats possibly just causing the nitrate -> nitrite reduction, producing Fe2O3.

The completely thermal approach appears to the reaction seems unpredictable, so I am currently trying an aqueous approach. Transition metals seem to be very good at catalyzing oxido/reducto reactions.

[Edited on 4-17-2012 by Polverone]
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