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Author: Subject: Preparation of ionic nitrites
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[*] posted on 2-6-2002 at 09:57
Preparation of ionic nitrites


I have been looking into various methods outside of the NaNO3/Pb method for preparing sodium nitrite. The first method was heating calcium sulfite and sodium nitrate together. This had seemingly good yields of sodium nitrite; however, most of the sodium nitrite was destroyed when I forgot that I was boiling off the water on an electric burner outside, and it was fried for around 10 minutes... I prepared the CaSO3 from NaHSO3 and CaCl2. The procedure for preparing CaSO3 was mixing stoichemical amounts of NaHSO3 and CaCl2, then wetting the mix; SO2 gas was liberated, CaSO3 was formed, as well as NaCl. The sludge left over was scraped into a filter, then was run through the sludge, which removed the sodium chloride. Of course, SO2 gas, generated by reaction of the NaHSO3 with an acid, could have been bubbled through a solution of Ca(OH)2 to form CaSO3 as well.

This is what I tried:
I placed 15g of CaCl2 in a beaker, along with 28.1g NaHSO3; then added 50mL of hot water. It fizzled, and liberated a good-sized amount of SO2 gas. A few minutes after SO2 gas was no longer being liberated, I dumped the contents of the beaker into a filter. I poured an additional 100mL of water through the filter, to insure that very little NaCl was left. I lost some CaSO3, due to CaSO3 being somewhat soluble in acids (and the solution formed after reaction of the CaCl2 and then NaHSO3 was obviously sulfurous acid, due to dissolved SO2 gas). I then heated the contents of the beaker with a "hotplate" that warms to about 120C, to dry it. I soon had a yellowish sludge, which is Ca(HSO3)2, which is simply CaSO3 and H2SO3 bonded together. After a while, all of the sulfurous acid was finally driven off, and I was left with dry CaSO3. I had 13.3g of CaSO3; an 82% yield. I then proceeded to heat the 13.3g of CaSO3 with 9.4g of NaNO3, on a propane burner. I stirred constantly to insure even heating. After about five minutes, I had a deep-yellow solid mix; the color didn't change any more after that point. I dumped the contents of the beaker into 500mL of water, and filtered. I began heating the filtered solution (which was a golden yellow), but I forgot about it... and ended up toasting the NaNO2 for quite a while with that electric burner. I'm confident it was destroyed by that, because when I try to dissolve that solid in water, I just get a brownish mix; the color of the solid reminds me of soil high in clay content.

This is my second idea for preparing sodium nitrite; this one has not been tested as of yet, but will be shortly. It requires sulfamic acid (HSO3NH2), calcium oxide, and sodium nitrate.

CaO + 2HSO3NH2 --> Ca(SO3NH2)2 + H2O

This reaction will have to use alcohol as the solvent, not water. Sulfamic acid hydrolizes to ammonium hydrogen sulfate; calcium sulfamate hydrolizes to calcium ammonium sulfate...

Ca(SO3NH2)2 + CaO + 3NaNO3 --> 2CaSO4 + 3NaNO2 + N2 + 2H2O

Ca(SO3NH2)2 is very water soluble. CaSO4 is not. Since the calcium sulfate formed in the second reaction would not dissolve in water, while the sodium nitrite would, extracting the sodium nitrite would be very easy.

[Edited on 2-6-2002 by madscientist]




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[*] posted on 6-8-2002 at 21:20


I made some sodium nitrite today. This is the process I used:

2HSO3NH2 + 2CaO + 3NaNO3 ----> 2CaSO4 + 3NaNO2 + 3H2O + N2

I thoroughly mixed 20g of powdered NaNO3, 8.8g of powdered CaO, and 15.2g of powdered HSO3NH2 (sulfamic acid). I then began heating it on a propane burner, stirring rapidly to insure even heating. After a few minutes, the mixture suddenly began fizzling and billowing water vapor. I stopped heating the beaker, and the reaction continued to accelerate. After about three minutes, it stopped reacting. The reaction certainly was more exothermic than I expected. I suppose a stoichemitric amount of Ca(OH)2, or even CaCO3 could be substituted for the CaO - but the reaction would be much less exothermic.

After the contents of the beaker cooled down, I then poured 200mL of hot water into it. It fizzled quietly (I'm hypothesizing that was a small amount of leftover sulfamic acid hydrating, forming the ammonium ion which was being oxidized to nitrogen gas by the nitrite ion), but soon stopped. I poured it through a filter. A considerable amount of CaSO4 was caught in the filter. The filtered solution was a beautiful light, golden yellow. I'm currently heating it gently in a flask (don't want the nitrite ion being oxidized by oxygen gas), at about 90C, to drive off water. Report on the yield is on the way.




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[*] posted on 7-8-2002 at 12:48


The theoretical yield was about 16.2g of NaNO2. I managed to scrape about 11g of NaNO2 off of the bottom of the flask. The NaNO2 is a very pale yellow.



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[*] posted on 9-8-2002 at 04:19


i make mine like this:
2Al + 3KNO3 ---> Al2O3 + 3KNO2
it needs to be molten about 40 minutes and its done. i think its a decent method just for the ease of doing it. my KNO2 is a very pale yellow/white.
madscientist i like your method by the way.
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[*] posted on 9-8-2002 at 11:49
another way


The other night I was looking at Muspratt volume 1 under "nitrous ether" (ethyl nitrite) and I found what seems like an interesting, simple method. Mix 100 parts of very fine KNO3 with 12.07 parts of lampblack (I'm guessing any finely divided carbon would work - I hope so), heat it in a crucible, and cover it and remove it from the heat source when the reaction is concluded (http://bcis.pacificu.edu/~polverone/muspratt1/c-834.html). According to the book some carbonate is formed but it is mostly pure enough to be used as-is (at least by the standards of 1860). It also mentions the formation of silicate - not sure if this is because they started with impure KNO3 or because of the crucible's composition. I am going to try this soon, not because I need more nitrite but because every other preparative method I've come across has been such a pain.

BTW, kingspaz, what physical form is your aluminum in? Powder, granules, foil, wire?

[Edited on 18-2-2003 by Polverone]
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[*] posted on 9-8-2002 at 13:22


my Al is in spherical powder form about 300mesh i think. i got it from some fibreglass place. it doesn't burn easy so can't be used for flash. i'm sure coarse powder made from Al foil would do the job. i just chose powder as it has a large surface area so should react better.
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[*] posted on 9-8-2002 at 15:35
The Muspratt method


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.

While I waited, I chipped some of the molten material (now solid) off of the concrete where it had splashed out. This material had gas bubbles and pockmarks in it. It was a pale yellow and had a pearlescent sheen to its surface. Adding a bit of this to water and tossing in a splash of HCl, I was rewarded with bubbling and thick orange-red fumes.

After it had cooled, I broke the bulk of the charge out of the dish. It had a mottled appearance, with brighter yellow spots mixed in with a majority of pale brown material. The pale brown material looked like what I had obtained when I had unsuccessfully attempted to make KNO2 by thermal decomposition in the past. When I added lumps of this substance to HCl, I also had considerable bubbling, but the orange gas production was much weaker than before. I am guessing that the yellow is the nitrite and that the brown is... not.

These were promising preliminary results. It seems that the stuff that spent more time in a heated state was "overdone"; perhaps the charcoal should be incrementally decreased or the charge should be ignited in a larger vessel (where heat would be better dissipated). Or perhaps, unknown to me, the metals in stainless steel decompose KNO2 at higher temperatures, which would also explain the superiority of the material that splashed out.

But I'm probably not going to be the one doing the further research. I buy NaNO2 by the pound and it's cheap. However, for anyone desperately seeking nitrites, I suggest you try this method first. It is simple, the reaction is fast, and the materials are readily available.
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[*] posted on 9-8-2002 at 17:44


I added a small amount of my product from the sulfamic acid process to a sulfuric acid solution (about 25% concentration, I estimate - I wasn't being exact, it wasn't necessary). I didn't add enough to create a visible amount of NO2 (since I have so little NaNO2), but the odor of the gas that was rapidly given off was unquestionably that of NO2. I'm very familiar with that scent.

Polverone, I was pulling your leg; I guess I was a bit too sincere-sounding when joking around. :P




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[*] posted on 3-9-2002 at 02:49


Polverone: Do you have a guess on the yeild of the Murspratt method?
And if the yeild isn't great, what can be done to eliminate the "pale brown material"?
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[*] posted on 3-9-2002 at 10:53
Sorry, no estimate


of the Muspratt-method yield. I haven't yet tried repeating the experiment since (like I've said before) I have easy access to NaNO2. However, if you want nitrites, I heartily recommend that you try the Muspratt method. KNO2 is much more soluble than KNO3, so it should be relatively easy to get mostly-pure KNO2. NaNO2 and NaNO3 have nearly identical solubilities (at 20 C; I don't have temperature/solubility curves for them) so they can't be separated so easily.

I'm not sure what the brown stuff is, so I'm not sure how it could be eliminated. I'd suggest making relatively small (50 g) batches of mixture with more and less fuel, igniting under similar conditions, and seeing which appears to yield the least brown crud (or dropping fragments of cooled material into acid and visually determining which gives the richest color).

In any case I would suggest crushing/powdering the material that is left behind and leaching it with a small quantity of water. The most soluble portions of the mixture will contain more KNO2. Or, taking the opposite approach, you could boil a small amount of water with the material, cool, filter out any crystals (should be unreacted KNO3), and then evaporate the remaining liquid. If you use too much water, of course, you'll achieve no separation.

If you want to find impurities, try mixing a concentrated solution of your KNO2 with one of calcium nitrate; cloudy precipitate indicates carbonate and/or hydroxide.
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[*] posted on 3-9-2002 at 11:40


dissolving NO2 in H2O gives
HNO3 + HNO2, the latter dedcomposing and the former being completely ionized.

2NO2 + H2O => HNO3 + HNO2
(no news for anyone bewanderd in inorganic chemistry)

However , dissolving NO2 in an alkaline solution (preferably KOH) yields
NO2- and NO3-

2 OH- + 2NO2 => H2O + NO2- + NO3-

separation is quite easy since KNO2 is more soluble at low temps.

Making NO2 is the difficult part since it so toxic (as is NO). One easy method is dissolving metalls in HNO3, HCl+ XNO3 or H2SO4 + XNO3.

High conc of H+ favours the forming of NO2, low conc NO.

/rickard

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[*] posted on 10-12-2002 at 07:18
For god's sakes, don't let it touch aluminium!


I finally got another five kilos of KNO3 and was able to try the muspratt method of preparation. I used 100g KNO3 and 12-14g of lampblack (which was probably not pure at all)(which is funny because I was guessing how much to use). I roughly mixed the ingredients and put them in a tin - assuming that homogeniality[sp.] didn't matter much once the reactants were molten.

As you said, the reaction was slow but vigorous, liberating a hell of a lot of white crap (much to my neighbour's apparent distaste). The solid was concurrent with what you had described. I made a solution (heaps of colloidal crap) filtered it, and intend to dry it out in the sun.

A couple of things I noted with the experiment.
- Small amounts of gas was liberated when making the solution. not sure what to make of this.
- The solution was basic, I know this for sure.
- When in contact with aluminium foil for just a short period of time, a strong odour of ammonium was noted.

This last observation is worrying, because we all know that when nitrites combine with amines they form nitrosamines, which are really carcinogenic, and the only way I can imagine ammonia got into the cycle is via NH2

I think something like...

KNO2 + 2H2O ===> KNH2 + OH(-)

This also accounts for the basic solution.

Please correct me!
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[*] posted on 18-2-2003 at 17:45


Here's an idea that occurred to me a while back for preparing nitrites (this should also work for preparing nitrous acid, of course).

2NaHSO<sub>3</sub> + 2NaNO<sub>3</sub> --(heat)--> 2NaHSO<sub>4</sub> + 2NaNO<sub>2</sub> ----> 2Na<sub>2</sub>SO<sub>4</sub> + H<sub>2</sub>O + NO<sub>2</sub> + NO

The sodium bisulfate won't be able to liberate a significant quantity of nitric acid, as the pK<sub>a</sub> is too high (not acidic enough). The generated gas should be almost entirely composed of equamolar quantities of NO<sub>2</sub> and NO.

Basically the idea is to mix equamolar quantities of sodium bisulfite powder and sodium nitrate powder, heat in a flask, and bubble the generated gasses into a strongly alkaline solution, yielding relatively pure nitrite (or into water to yield metallic cation-free nitrous acid). The generated gasses must not be allowed to come into contact with atmospheric oxygen, or much of the nitric oxide will be oxidized to nitrogen dioxide, rendering the gasses nearly useless for preparing relatively pure nitrite (unless one plans on attempting fractional crystallization).

Edit: The edit button works! :D

Ramiel, this is my hypothesis as to what the reaction was between potassium nitrite, aluminum, and (I assume) water that generated ammonia:

2KNO<sub>2</sub> + 4Al + 6H<sub>2</sub>O ----> 2KOH + 2Al<sub>2</sub>O<sub>3</sub> + 2H<sub>2</sub>O + 2NH<sub>3</sub>

[Edited on 19-2-2003 by madscientist]




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[*] posted on 19-2-2003 at 02:02
Easy NaNO2?


I had heard NaNO2 could be prepared by heating, and boiling, NaNO3 for 15 minutes or so. I assume this liberates the loosely held Oxygen but not the more strongly bonded ones? Anyone tried this, or is it that you have to either buy it or reduce NaNO3 with something other than just heat?

Also, some preparations call for distilled water - can bottled water (Evian etc) be used? I assume it's of a suitable purity?
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[*] posted on 20-2-2003 at 15:27


xoo, my method above seems to work quite well. the trick is to keep the KNO3 just above its melting point (334*C if i remember correctly). this can of course be judged by the KNO3 melting :)



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[*] posted on 20-2-2003 at 19:58


On the Muspratt method, the procedure mentions using 'parts' instead of grams. The people who tried this used grams instead of 'parts'. Might this be the problem as the density of the two chemicals are most likely not equal. Perhaps I'm missing something here though.
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[*] posted on 11-9-2003 at 12:34


Some more ideas:

Na<sub>2</sub>S<sub>2</sub>O<sub>4</sub> + 3NaNO<sub>3</sub> ----> 2Na<sub>2</sub>SO<sub>4</sub> + NaNO<sub>2</sub> + NO + NO<sub>2</sub>

Na<sub>2</sub>S<sub>2</sub>O<sub>4</sub> is the ingredient in many solid toilet bowl cleaners (known as "sodium hydrosulfite.";)

3NaNO<sub>3</sub> + S ----> Na<sub>2</sub>SO<sub>4</sub> + NaNO<sub>2</sub> + NO + NO<sub>2</sub>

Those reactions could be of interest if they proceed as predicted due to the possiblity of providing a dry equamolar mixture of NO and NO<sub>2</sub>. Isolating the solid nitrite produced surely would prove to be difficult.




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[*] posted on 12-9-2003 at 05:47


3NaNO3 + S ----> Na2SO4 + NaNO2 + NO + NO2

How did you plan to carry this out? Seems like a pyrotechnic composition to me that'll just burn leaving SO2 and K2O.




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[*] posted on 12-9-2003 at 07:35


I'm not sure how well it would work. If I remember correctly, a significant quantity of the sulfur in black powder ends up as a sulfate salt after deflagration. But I've never tried it, and I guess it probably wouldn't be too useful since it'd be so difficult to eliminate SO<sub>2</sub> from the product.



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[*] posted on 12-9-2003 at 07:52


Could it be that there is also a significant amount of SO3 present? We're using a powerful oxidizer and high temperatures here, so why not? Maybe that's where the sulfate comes from.



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[*] posted on 12-9-2003 at 07:55


Definitely. I was expecting the SO<sub>3</sub> to displace the other acid anhydride though:

NaNO<sub>2</sub> + SO<sub>3</sub> ----> Na<sub>2</sub>SO<sub>4</sub> + NO<sub>2</sub> + NO




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thumbup.gif posted on 24-1-2004 at 00:56


Whether, heating Sodium sulfite (Na2SO3)and calcium nitrate(Ca(NO3)2.4H2O )together will yield Calcium nitrite?

if yes at what temp.?
i am trying to reduce calcium nitrate to nitrite, using S, Na2S,Na2SO3
have u tried these ?
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[*] posted on 24-1-2004 at 06:35


" When in contact with aluminium foil for just a short period of time, a strong odour of ammonium was noted. "

I don't think you need to worry about nitrosamines - Aluminium easily reduces Nitrate/Nitrite to Ammonia - especially in basic so,ution - thats where your NH3 is coming from.
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[*] posted on 13-8-2004 at 22:27


I just wanted to add that I can now confirm: CaSO3 + NaNO3 -> CaSO4 + NaNO2 works fairly well. My CaSO3 was prepared from CaCl2 and sodium metabisulfite, my NaNO3 from NaOH and NH4NO3. There was some water of crystalization left in one or the other of the reactants, since I saw water vapor in the test tube before any sort of reaction.

There were various sulfury smells during it. The mass of reactants melted but never formed a homogeneous melt or fluid; it bubbled and sluggishly slumped. During cooling, as always happens, the test tube cracked :mad:. I used one of my largest tubes, too. I will have to run a larger-scale reaction in a steel can and see how that turns out.

For this proof-of-concept run, I didn't separate the NaNO2 but acidified the cooled reaction mass as it was. I was rewarded with a rich bubbling of NO2.




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[*] posted on 14-8-2004 at 01:54
Re: Nitrites


If you have a means of generating large amounts of the gaseous lower oxides of N2, particularly NO and N2O, nitrites should be easily obtainable by bubbling the gas into a concentrated alkali solution.

What do you want to use the [expletive deleted] stuff for? I have heard of nitrites, particularly nitrite esters such as amyl nitrite (presumably obtained by reacting the corresponding alcohols with nitrous acid or inorganic nitrites plus sulfuric acid), being used as antidotes for blood-circulatory or heart poisons like cyanide.

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