Sciencemadness Discussion Board

Preparation of ionic nitrites

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madscientist - 2-6-2002 at 09:57

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]

madscientist - 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.

madscientist - 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.

kingspaz - 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.

another way

Polverone - 9-8-2002 at 11:49

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 ( 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]

kingspaz - 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.

The Muspratt method

Polverone - 9-8-2002 at 15:35

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.

madscientist - 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

Ramiel - 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"?

Sorry, no estimate

Polverone - 3-9-2002 at 10:53

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.

rikkitikkitavi - 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.


For god's sakes, don't let it touch aluminium!

Ramiel - 10-12-2002 at 07:18

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!

madscientist - 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]

Easy NaNO2?

AngelEyes - 19-2-2003 at 02:02

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?

kingspaz - 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 :)

Haggis - 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.

madscientist - 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.

vulture - 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.

madscientist - 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.

vulture - 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.

madscientist - 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

karandikarmv - 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 ?

kryss - 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.

Polverone - 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.

Re: Nitrites

JohnWW - 14-8-2004 at 01:54

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.

John W.

Polverone - 14-8-2004 at 10:40

I buy my NaNO2. This was just to see if there was an easier way than the traditional reduction of molten nitrate with lead to prepare it on a lab scale. It's useful for preparing hydroxylamine, azides via hydrazine and alkyl nitrites, and diazo compounds. It can also be used to perform a few organic oxidations or to prepare small quantities of pure nitrogen.

JohnWW - 14-8-2004 at 12:58

Sodium nitrite is also widely used as a food preservative, mostly in meat small-goods like sausages, bacon, and "corned" and smoked and canned meats. However, some studies have shown that excessive consumption can be carcinogenic, possibly through its formation of nitrosylamines when metabolized in the presence of amino-acids.

John W.

BromicAcid - 14-8-2004 at 17:35

There was actually an entire chapter on that very effect in my college toxicology class. But it is determined that the amount used is a 'safe' dose only contributing a few hundred (possibly thousand) cases of cancer world wide a year.

Other method

chloric1 - 14-8-2004 at 19:40

I would not know percentage of yeilds but...Potassium Nitrate and granulated sugar in approx 6:4 ratio. This is a potent pyro mix and leaves molten nitrite behind! I know because Conc. HCl gave lot of NO2 with traces of Chlorine from HCl oxidation. be very careful becuase the smallest spark sets this mix off and it burn vigorously at WHITE HEAT:o Also ignition of KMnO4 with sulfur produces the green form of MnS!

frogfot - 21-9-2004 at 02:34

I recently got some strange results with Pb/KNO3 procedure..

Melted 207 g 99,95% lead (bought from a metal supplier) in an iron crucible, and while stirring mechanically I added 101 g KNO3 in small portions (5-10 min). This was then continiously stirred mechanically, mix foamed a little. Foaming increased a bit after about 12 min and reaction mix started to glow from beneath. After a while whole mix started to glow red and it started to give of thin white fumes! The bottom of crucible was red hot on outside.. Here I ceased heating.

It cooled after about 1 min. Though while stuff was still dark red glowing, about half of the mix have separated as solids, while rest was liquid. I'd guess that solids was sodium oxides because of high mp and because extraction liquid was very basic. I'm currently extracting it..

So, is this glowing normal? I'd guess that this was a runaway reaction, but how could one prevent it using only a gas burner..? Maby it ignited because of friction between stirrer and crucible..

In earlier test I used same ammount of carbattery lead (~6% antimony). After ~40 min reacting (with same setup) this gave 17% conversion with regard to recovered lead.

EDIT: Oki, I've got about 55g KNO2, most likely not pure.. Maby that "runaway" wasn't that bad, though there are probably lots of KOH in the product..

[Edited on 21-9-2004 by frogfot]

[Edited on 22-9-2004 by frogfot]

Theoretic - 25-9-2004 at 05:24

The solids were probably a mix of KPbO2 and K2PbO3, they hydrolysed to give KOH.

I suggest that nitrate is reduced in molten state by Na2SO3. I don't see any probable side reactions, to separate the product from Na2SO4 let it cool, I think they'll form two separate layers. To fully use the reducing capacity of NaHSO3, convert it to Na2SO3 with Na2CO3, then SO2 won't be lost during conversion to CaSO3, and CaSO3 will all precipitate, you'll also get twice as much. Reduction with Na2S in a molten state should also work.
A combined reduction with sulfur should be interesting. Sulfur powder can be melted down and solidified to form a gob (to prevent runaway and over-reduction), then put into molten nitrate. It should be reduced like so: 2NaNO3 + S => 2NaNO2 + SO2. SO2 should be led into Na2CO3 solution (not, NaOH, for reasons of greater availability and lesser alkalinity, which could otherwise induce suckback). Na2SO3 thus formed could be reused for reduction.
To prevent splashing during carbon reduction, first metl down nitrate and then throw in C powder, bit by bit.

making KNO2

crushpack - 29-10-2004 at 01:23

Here's my contribution to problem of KNO2(potassium nitrite) making. Here in Czech rep nobody can buy this substance legally and the classic way for its preparation (heating of melted mixture of lead and potassium nitrate) is at least somewhat complicated and no ever successfull.

so what's this method I use: the yield is practically kvantitative and it's simple and fun. I can't remember no disadvantage of that.

You must have some lead tartrate by hand before you start, if not then let it precipitate from the solution of sodium potassium tartrate(cream of Tartar) to which you pour the soln of lead nitrate- nothing can't be simpler.
Once you have this "organic" lead add it to the melted Potassium nitrate, but only in small portions, since the reaction that occur is really violent. the mixture of KNO3 and lead tartrate is pyrotechnic in fact.
When you've done this you've done all if you don't consider purification of KNO2 from PbO2 a problem. (If you do, then extract it by hot water and crystallize).

For those interested in reaction mechanism here is my own explanation(which means that it can be wrong):
when you heat lead tartrate you'll get a pure lead but in so a very fine form that its pyrophoric, which means that it will burn in air (when aerosolized) or with some oxidizer as f.e. KNO3:

Pb + KNO3 = PbO + KNO2

So it seems that only differrence between this method and the classic one is only in the form of lead. this method takes an advantage that the larger the surface of the reactants is the bigger is the yield.
Its simplificated but I can't give here more detailed explanation,for I must go now. So I'll give here only the total reaction scheme (it can be wrong I repeat and for some other suggestions i'll be thankfull)

PbH2C4O6 + 4KNO3 = PbO + 4KNO2 + 4CO2 + H2O

Edit by Chemoleo: Çrushpack, there was an existing thread on nitrites as you can see. Although your contribution is appreciated, a quick search does help too :)

[Edited on 29-10-2004 by chemoleo]

crushpack - 1-11-2004 at 00:31

Apologize: the formula for potassium tartrate is PbH4C4O6, I've missed two hydrogens attached directly to carbons...So the equation of reaction should be corrected:

PbH4C4O6 + 5KNO3 = PbO + 5KNO2 + 4CO2 + 2H2O

Anyway, since I really can't be sure that the reaction run exactly according to this equation maybe that was not so a big mistake. But the truth is that the KNO2 really originate by this (when I dropped a little of diluted H2SO4 on isolated product, voluminous red foul smelling fumes of NO2 appeared). I've also in my hurry forgot to post something about advantages of this method (compared to classic reduction of PN with lead):

1/the reaction does not need stirring

2/lead tartrate regeneration from PbO byproduct is possible. Just transfer PbO to Pb(II) salt by some acid (acetic is preferable) and then to precipitate lead tartrate by some soluble tartrate salt (Cream of Tartar f.e.)...let's imagine how would you try to get pure Pb from PbO instead)

3/ Normally you need to bubble CO2 through the aqueous soln of prepared KNO2 to get rid of some soluble lead salts by precipitation. I suppose that the CO2 originated by this "improved" method should be sufficient to solve this problem completely.

Well, one disadvantage is here as well:
when you put lead tartrate into the molten KNO3, you must do it carefully in reasonable small batches. When the "organic" lead is added to molten PN at once or when mixed with it before and then heated, it has almost the same effect as black powder under same conditions.

I'd be thankfull if someone would find out the real yield of KNO2, since I have now lack of time for experimentation. :(

JohnWW - 1-11-2004 at 01:13

"Apologize: the formula for potassium tartrate is PbH4C4O6"? You will need to give another apology: that is the formula of lead(II) tartrate, not potassium tartrate.

John W.

crushpack - 3-11-2004 at 01:41

thank you, it was only misprint from me, I wanted to give the formula for lead tartrate of course

potassium nitrite - qualitative success

skippy - 10-1-2005 at 20:43

I've given the fabled KNO3 + Pb -> KNO2 + PbO a couple of tries and have a report to make.

Basically it comes down to stirring, from what I've seen. I tried three attempts and each time I stirred more than the last and each time had better results. On my third attempt I bent a fork and lashed it to a dowel so I could better mix the pot. The prongs were bent so they would both scrape the bottom of the pot and comb through the lead and nitrate. This seemed to really help. The lead gets broken up during the stirring into little dark grey lead flecks that react alot better than a big pool of lead.

The melt was reacted for 1 1/2 hours, with constant stirring, and then allowed to cool. The salt was dissolved in warm water and a cake of tan orange lead oxide intersperced with lead flecks was left. A small fused pool of lead was underneath this cake. The amount of lead left suggests that well over 50% of the nitrate was converted. The salt extracted also reacted vigorously with HCl, producing copious deadly brown fumes. :D

Anybody have a good idea on how to test how pure the nitrite is?


Stir, stir, stir!

When/if I get time to do this again, I think I'll try a mechanical stirrer, as it is not much fun stirring this stuff for so long!

Edit by chemoleo: The title is NitrIte, not NitrAte. Merged thread with existing thread.

[Edited on 16-1-2005 by chemoleo]

PainKilla - 16-1-2005 at 09:47

Why bother with that reaction... Im pretty sure that KNO3 decomposes to KNO2 via heating.

2KNO3 ---(heat)--->2KNO2 + O2

Or am I mistaken?

frogfot - 16-1-2005 at 12:28

Skippy, what kind of lead did you use?
With car battery lead, reaction seems to go very slow like you've described.

S.C. Wack - 17-1-2005 at 14:10

Many find the simple heating of nitrates unsatisfactory. I made some this way, but then I found a better method - burning some cash instead of some nitrate. The loss of O is slow below 650C, but the loss of N (Nasty oxides) is fast above 850C. One ends up with an equilibrium mixture of nitrate and nitrite up to 800 or so, where oxide formation becomes a factor.

This is because the nitrite, despite the usual temp of dec given, loses O quite slowly to a point, but gains O from the air at a higher temp in a narrow range. There isn't a whole lot of room between this range and the loss of N temp. 780 looks favorable, though I've never seen a ref explicitly say so or give exact yield.

Experimenters may find it best to use a scale and watch out for darkening.

Mellor's N supplement gives refs for alternatives to Pb, etc - but they are all in German.

skippy - 17-1-2005 at 18:07

The lead was from some old lead piping I had kicking around. I don't know how it compares with battery lead.

I guess the whole deal is a bit of a hassle if all you want is the nitrite, but I wanted some PbO too, so getting both seemed pretty cool.

frogfot - 18-1-2005 at 05:14

About nitrite content analysis...

Do anyone know if reaction:
NH4Cl + NaNO2 --> NaCl + 2H2O + N2

goes to completeon? Maby some very kind person can even test this with pure nitrite (rolleyes).

If reaction is compleate, this would be a very simple method to determine the content of nitrite. One would only collect evolved nitrogen in an upside down measuring cylinder (filled with water) and then substracting moisture content. This could give very precise results.

In inorganic book they use following procedure:
Saturated NH4Cl solution is heated nearly to boiling and a saturated solution of NaNO2 is added dropwise.

Since I recently got a mechanical scale (+/-0,01g) I'm quite eager to analyse my nitrite from the runaway Pb/KNO3 reaction I mentioned above..

[Edited on 18-1-2005 by frogfot]

S.C. Wack - 18-1-2005 at 09:59

Vogels Inorganic Analysis says that of trace of nitrate is formed. How much a trace is, they don't say.

frogfot - 22-1-2005 at 13:44

Alrighty, I've just conducted abovementioned nitrite determination experiment.. Set up a twonecked flask with reflux condencer and a rubberstopper (with syringe). The gas outlet was connected to a washing bottle with dilute H2SO4 and then to a reversed measuring cylinder. Washer with acid was to remove formed ammonia (my nitrite is very basic).

I refluxed NH4Cl soln, and when pressure stabilised, I added a water soln of 1,00 g of my impure KNO2 (from runaway exp) with a syringe, dropwise. This gave 215 ml gas, I calculated content of KNO2 to 0,705 g (taking into account added volume of nitrite, the temperature of gas and partial pressure of water).
This means my KNO2 is of about 71% purity.. Quite simple test, though I'm not sure about it's accuracy.

Btw, "trace" of nitrate wouldn't hurt :)

[Edited on 22-1-2005 by frogfot]

[Edited on 22-1-2005 by frogfot]

Sodium or potasium nitrite

CycloKnight - 25-1-2005 at 07:54

Is there a reliable process for converting the nitrate into the nitrite using heat?
I recall using heat to decompose KNO3 into oxygen and nitrite. Does anyone have any experience doing this on a scale larger than test tube proportions?
I did use the search engine, and I couldn't find the info I'm looking for.

I read that sodium nitrate can explode if heated beyond a certain temperature. Is that only if contaminants are present or is that a genuine risk once takes by heating sodium nitrate to decomposition?
I assume the latter to be the case.

Polverone - 25-1-2005 at 13:36

Sodium nitrate alone will not explode from the application of heat. People here have had some success making nitrites on a somewhat-larger-than-test tube scale by making fuel-poor pyrotechnic mixtures of nitrates with sugar or charcoal and igniting them. Muspratt suggests heating a fuel-deficient mixture of KNO3 and lampblack to produce crude KNO2. The classic method of stirring molten nitrate with molten lead has also been used on a modest scale, though it is tedious. I have tried heating calcium sulfite with NaNO3 in just sufficient quantity to reduce it to NaNO2, as per a patent, and it seemed to work though I did not scale up beyond test-tube quantities.

A zinc-copper couple in nitrate solution or amphoteric metals in alkaline nitrate solution are supposed to reduce nitrate to nitrite. Unfortunately, they will keep on reducing to produce ammonia. I don't know if the process can be controlled to give reasonable yields of nitrite if the quantities are calculated and vigorous stirring is used.

I found a one-line reference in an old book saying that zinc dust would reduce a neutral aqueous solution of KNO3 to KNO2, but I never saw any reaction myself. Perhaps my zinc was too pure, or there were unnamed necessary conditions (like the reaction taking place in an autoclave at 150 C).

I know that heating aqueous ammonium nitrate solution with lead dust will produce some lead nitrite, though that's not exactly what you're after.

Simply heating NaNO3 or KNO3 strongly can give a fair amount of nitrite. It's good enough to make alkyl nitrites without further purification, but it wouldn't be suitable for all purposes. I see no reason why it shouldn't work on larger scales, but you need a suitably durable vessel and an intense heat source. Finely divided manganese dioxide is supposed to accelerate the thermal decomposition of nitrates.


CycloKnight - 26-1-2005 at 03:20

That is very informative, thanks Polverone. That is indeed the information I'm looking for. In the near future, I will be conducting some experiments so I can establish a working, reliable and replicable procedure for the conversion of sodium nitrate to nitrite. At that time, I will post the results on this thread.

runaway lead + nitrate

skippy - 2-5-2005 at 08:36

I've found that Frogfot's above mentioned runaway reaction is actually a good thing. Without the runaway, the reaction procedes too slowly. The runaway can be started consistantly by strong heating while beating the molten mix of nitrate and lead with an old whisk. (protective clothing!). The stirring seems to whip the lead into fine granules which can be burned quickly if the nitrate is sufficiently hot. The 2l stainless steel pail I was using turned a bright cherry red from the exothermic heat. It was a pretty cool sight, a glowing 1000 deg C pot of seething molten goo!

Here's some good stuff from the hive archive (

After removing the small dome from the inner file-clay mantle of a Rossler gas-furnace, place upon the mantle a strong iron-wire triangle, and set upon the triangle a shallow iron dish (2.5cm high, 12cm upper diameter) having a smooth bottom. Place 85 grams of Sodium Nitrate (NaNO3) in the dish and close the furnace. As soon as the dish has become faintly incandescent and the molten nitrate just begins to give off bubbles of oxygen, gradually add 206 grams of lead in the form of old pieces of sheet lead or lead tubing. The lead is at once vigorously oxidized, and, if sturred continually with an iron spatula, becomes almost completely converted into oxide of lead in half an hour. Empty the contents of the small iron dish into a large deep iron one, and repeat the operation several times, using the same amounts of Sodium Nitrate and lead. Place the various products in the large iron dish, extract once with boiling water, and decant upon a creased filter. Dry the reside of lead oxide and set aside for the experiment on page 52. Pass a strong current of CO2 into the still boiling-hot filtrate,for a few minutes only, filter off the lead carbonate which separates, and neutralize the solution while stirring it, by carefully adding nitric acid from a pipette or burette. Evaporate the solution to crystallization. The crystals which separate first, consist partly of nitrate and may be used again for remelting with lead; the mother-liquor gives pure nitrite. A normal solution of the nitrite is prepared by dissolving 69g of it in water and diluting to one liter.

frogfot - 2-5-2005 at 23:05

Great!! Skippy, I bow to you.
This means I have a reliable supply of nitrite for this summer (and further). One only needs to come up which nitrate is better to use.. sodium or potassium..

Seems like syntheses with both NaNO3 and KNO3 has a possibility of purification. That is, excess of the first can be precipitated from aqueous solution by alohol and the second can be removed by cooling out. So, using KNO3 seems to be cheaper..
I only have KNO3 on hand, and NaNO3 has to be made (from Ca(NO3)2 and Na2CO3).

garage chemist - 21-5-2005 at 08:54

I had some sucess with the preparation of NaNO2 by simple heating of NaNO3.

I used a 50ml quartz crucible, which I find to be a very useful piece of equipment. They are thermally indestructible, you can heat them to red heat and dunk them in cold water without any risk of cracking.

I melted some NaNO3 and heated it until it was bubbling at a slow but steady rate. It was held at this temperature for 20 minutes.
From time to time, carefully smell the contents of the crucible and check if it begins to smell of NO2. If it does, it's too hot and the nitrite starts decomposing to nitrogen oxides and Na2O.
Only Oxygen should be given off.

After cooling, I dissolved the residue in a small amount of water, and on addition of a few drops of HCl, it fizzed vigorously and lots of brown fumes were given off. I was surprised at how much gas was produced.

This method only produces a nitrite/nitrate mix, but the abovementioned separation procedures can of course be applied.
This seems like the "cleanest" preparation of nitrites to me.

Drying Sodium Nitrite

bio2 - 30-7-2005 at 16:40

As this seems to be the primary Sodium Nitrite thread I put the simple question here which I couldn't find searching.

What is the preferred method of drying NaNO2?

Just got some 99.1% recently and it seems a little damp! Also just how hygroscopic is this stuff?

How to dry NaNO2?

bio2 - 1-8-2005 at 15:50

What is the preferred method of drying NaNO2?

Will one of the inorganic guru's here please take a moment and help me with this question. I would just as soon not re-invent the wheel.

Also would fusing the nitrite at the 271mp be OK being careful not to approach the 320deg decomposition temp??

Alternatively how about vacuum dessicator at 90deg without (with) dessicant?

Thank you very much.

oldskool manufacture of sodium nitrite

komodo13 - 12-9-2005 at 18:10

J. Soc. Chem Ind., 27, 483-5 (may30)-The author reviews various methods for the production of nitrite from nitrate by reduction with metals (see C.A. 1908, 1330). the non-metal , sulphur, has been used with sucess by Messers, Read, Holiday & Sons. The mixture of nitrate, sulphur, and caustic alkali is fused in open pans, fitted with stirring gear. the following reaction takes place: NaNO3 + S +NaOH= NaSO4 +NaNO2 + H2O. The hot, fused product is added to sufficient warm water to dissolve the whole of the nitrite, leaving most of the sulphate behind in granular condition. The nitrite is filtered through vacuum filter and crystallized in fractions from the small amount of dissolved sulphate. Product is good nitrite, and troublesome by-products are avoided. source Gilbert T. Morgan 1909;)


jimmyboy - 12-9-2005 at 23:18

anyone try this already? what is the yield?

Magpie - 27-12-2005 at 11:08

This may not seem very exciting but we haven't had many experimental results lately so I though it worth a post or two.

My ultimate goal is to prepare some sodium nitrite using the lead reduction of sodium nitrate according to a procedure from "Laboratory Studies in Chemistry" (1923) by Robert H. Bradbury.

First off I needed some NaNO3 so have prepared about 30 grams of this by the neutralization reaction of my homemade HNO3 and Red Devil lye (NaOH). This reaction went well as I first diluted my 65% acid about 2:1 with water. Next I placed this in an ice bath with magnetic stirring. I then added the lye very slowly as the reaction is highly exothermic. I trimmed to neutral via pH paper then dried this brine at 110C for about 8 hours. Attached is a photo of the product.

When I get my 6" iron frying pan I will be ready to make the NaNO2. ;)

Edit: Will try again with the photo.

[Edited on 27-12-2005 by Magpie]

Magpie - 27-12-2005 at 11:49

The NaNO3:

[Edited on 30-1-2007 by chemoleo]

NaNO3.jpg - 74kB


chloric1 - 27-12-2005 at 15:16

Oh magpie!:D

You are not shy about showing off your love of chemistry. The holiday decor really brings out the snowy white hues of your nitrate.:D

Just so you know, since you are here in the USA you can buy nitrate of soda from garden supply centers. That is where I buy mine and I simply dissolve in boiling water, filter then cool. It is easy to get a pure white product at about $1 per pound. Save your nitric acid for more dignified uses. Like making selenium dioxide.

[Edited on 12/27/2005 by chloric1]

Magpie - 27-12-2005 at 16:41

I wasn't aware that I could buy sodium nitrate at garden suppliers, at least not in my area. I'll have another look around. Potassium nitrate, however, is readily available as a stump remover.

Magpie - 2-1-2006 at 22:03

My attempt to reduce NaNO3 to NaNO2 using the classical Pb reductant was a failure. I was suspicious when I noted all the unreacted lead in my iron frying pan even though Pb was not in stoichiometric excess.

I did end up with a fair amount of white salt crystals - NaNO3 I think. I tried several confirmatory tests all of which were negative.

So I thought I should read this whole thread. What an education! It's certainly not as easy as indicated in my 1923 lab procedure. :o

I will try again, if not with Pb, perhaps using the S + NaOH procedure. (I used fishing lead by the way - I suppose this is as good as any. I cut it into thin slices with a hacksaw.)

Microtek - 3-1-2006 at 07:48

If you do the NaNO3 + NaOH + S reaction, you need to find a suitable way to mix the reactants. If you just heat the mixture, it will explode and scatter hot NaOH.
In my opinion the mixing represents a serious problem because if you melt the NaNO3 and NaOH first and then add the sulfur, it tends to remain on the surface and burn. This may be allright, but surely some of the oxidation must be from atmospheric O2.
If you melt just the nitrate and then add a mix of the NaOH and S, the latter two will begin reacting to produce polysulfide which may also be OK but is quite messy.

Magpie - 3-1-2006 at 15:04

Thanks, Microtek, for the warning. I really want to make the lead reduction work anyhow - it's classic and would thus be more satisfying. ;)

I tried it again today using about 12 g of NaNO3 and 24 g of Pb. This time I put as much heat to it as I could using MAPP gas instead of propane. The frying pan did not get red but was close, I believe.

This time I tried to be patient although I was still done in what I estimate as 15-20 minutes. There was much more red mud this time which I take as PbO and a good sign. I could only see a few tiny BB's of lead at the end of the cook. I extracted the NaNO3 with boiling water and have it drying in my oven now. The extract had a slight yellow color which I also take as a good sign. When I washed out the frying pan there was some solid particles other than the fine red mud. I suppose that was lead as I don't know what else it coud be. But not nearly as much as last time.

The proof will be tommorrow when I do the NO2- identification tests. I invented a test of my own based on the different molar heats of solution (very endothermic) of the NaNO3 and NaNO2. That for NaNO2 is about 2/3's that of NaNO3.

I took some pictures, of course. I'll attach a couple in the next post.

Magpie - 3-1-2006 at 15:29

NaNO2 preparation pictures:

Note: there are 2 homemade laboratory aids (fabricated, unpatented) in these pictures (1 ea). I wonder if board members can identify them? Darkblade48 this is your specialty. :D

[Edited on 3-1-2006 by Magpie]

[Edited on 30-1-2007 by chemoleo]

end.jpg - 56kB

Magpie - 3-1-2006 at 15:34

2nd picture for NaNO2 prep:

[Edited on 30-1-2007 by chemoleo]

2nd.jpg - 42kB

chemoleo - 3-1-2006 at 16:27

Magpie, I am a little confused. Wouldn't you want a strong excess of Pb?

According to my calculation, it is 207 g Pb per 85 g of NaNO3 (Pb + NaNo3--> PbO + NaNO2), or 29.2 g lead per 12 g NaNO3. So I'd use probably 50 grams of Pb, but rather even more. That way you also end up with NaNO2 only, no mixture with NaNO2.

Also, isn't the formation of insoluble product a DIRECT indicator of how well reaction worked?
I.e. you should get theoretically 31.5 g of PbO if you used 29.2 g of lead. This is quite a noticeable yield, and should give a decent precipitate in solution. The weight of this dried precipitate alone will determine how well the reaction worked! (i.e. weigh filter paper, then filter extract, wash with h2O, and dry again, then weigh, to find out the total amount of PbO). From your filter paper, it didnt look like there was 25 g worth of PbO!

AT last, another suggestion would be to heat Pb in a can/cradle/crucible until it melts (not teflon gear, as when it decomposes, it is at 322 deg C. If it didnt compose your temperature is too low anyway). Then add, gradually, NaNO3. Wait until it has reacted. Add more NaNO3. EVentually scoup out the product, and continue this until all your NaNO3 is used up. This should facilitate a much greater efficiency of reaction, as the excess of lead is always enormous. Run this through a sieve to remove any Pb particles prior to water extraction.

Magpie - 3-1-2006 at 19:11

Chemoleo I'll try to remember your 3 questions and answer in order:

1. I agree that it is short on Pb. I noticed this when checking the stoichiometry before doing any experimentation. The procedure called for 10 g NaNO3 and 20 g of Pb. I'm guessing that the tremendous heating drives off some of the oxygen and a full complement of Pb is not required.

2. Yes, I take it that the red/orange mud is PbO and should be a direct indicator of the progress of the reaction. There was quite a bit of it. After adding the boiling water to extract the NaNO2 I just drained it off the mud, leaving it in the frying pan. Besides, the picture I show is of the 2nd filtration. There was more in the first filtration. I also wonder if some of this is Fe2O3 and if some of the oxygen came from the air. But, really, I think it mostly PbO and the oxygen came from the NaNO3.

3. These are not really homogeneous reaction conditions I don't think. The Pb sort of goes into very small droplets. I don't know if flooding it with more Pb would actually drive the reaction by LeChatlier effect.

The proof of the (hasty) pudding will be in the testing tommorrow. First off I should get some pale blue nitrous acid when I mix the salt with water and a little H2SO4.
I should also get some NO generated. When a little FeSO4 is added I should see the brown complex [Fe(NO)]+2.

Right now I am patiently evaporating the brine under gentle heat. :)

My procedure is all of about 5 sentences long. I'm thinking that every word is important. If it works I'll conclude that it is a masterpiece of brevity.

Magpie - 4-1-2006 at 09:44

Chemoleo it just dawned on me why you mentioned teflon - my frying pan? No, my 16cm frying pan is pure iron. You know, the kind granny used to cook biscuits and gravy on, only much smaller. :D

My salt was almost dry this morning so I tested it. Results:

1. color: white
2. acidified dilute solution color: clear
3. very small bubles generated when cold dilute solution is slightly acidified with 5N H2SO4 - assume nitric oxide, NO
4. FeSO4 test with solution of "4" above: brown color indicating [Fe(NO)]++

So the first two tests are negative and the last two are positive. I would rank the last test as most important.

When completely dry I will calculate a yield. I am now wonderng what test I could use to determine purity. Any suggestions?

Microtek - 4-1-2006 at 16:39

I'd suggest a redox titration with acidified KMnO4 which goes all the way to colorless Mn(II) so the end point is very clear. Of course, you will need a standard to calibrate your KMnO4 soln.

Magpie - 4-1-2006 at 19:07

Thanks Microtek! I have the KMnO4 but will have to find some nitrite standard. Then I'll do it for sure. I actually like analytical chemistry, at least the wet methods.

My NaNO2 finally dried today. I never let my drying oven get over 150F (66C). I did the final drying in a dessicator over CaCl2. Measured yield was 64%. This was after I had taken some out for analysis. So I'm fairly pleased. Perhaps my yield would have been better if, like Chemoleo suggested, I used more Pb, and if I had cooked it longer. But I just tried to follow my 1923 procedure to the letter. I did notice that my NaNO2 handled and dried differently than my NaNO3. When I placed it in my small bottle where it is more in bulk I could see that it has a very faint yellow cast. ;)

I thought that I could also use it for a diazotization. If that is successful then it is should be good enough for my mad science.

Edit: I think I could have increased yield by washing my PbO mud also, which I didn't do.

I checked the MSDSs for NaNO3 and NaNO2 after doing my NaNO2 preparation. NaNO2 is highly hygroscopic whereas NaNO3 is not. This is also confirmation that I do indeed have NaNO2.

Speaking of the MSDSs. It is a good thing I didn't read them before doing the NaNO2 preparation. All the WARNINGS ABOUT KEEPING AWAY FROM HIGH HEAT would have stopped me. :o

[Edited on 5-1-2006 by Magpie]

woelen - 5-1-2006 at 10:56

Magpie, that is a very nice result. Sodium nitrite indeed is not white. I have some quite pure sodium nitrite (it is lab grade, commercial sample from a lot of 1 kg) and it is light yellow. I also have lab grade KNO2 and that also is light yellow, exactly the same color, but the crystals are smaller.

Here follows a picture of the NaNO2:

I was so lucky to be able to buy this for just $12.50 per kg.

So, if yours also is light yellow, then I think that is a good sign, if you compare it to commercial samples.

Magpie - 5-1-2006 at 17:07

Woelen thanks for the nice picture of NaNO2. Mine doesn't appear to be quite that yellow. But it is hard to tell as I only have about 5 grams.

Food grade NaNO2 is available on the internet for about $88/kg. You did get a good bargain.

I asked a pharmacist at Wal-Mart if he would sell me some sodium nitrite that I could use as a standard. I was surprised and elated that he went right back and looked for some. But US pharmacies don't compound much anymore so didn't have any. Emboldened I tried this again at a Rite-Aid pharmacy. Same result: willing but not able. I then asked for some chloroform - he would have sold me this also if he would have had any. I must indeed look harmless. ;)

neutrino - 5-1-2006 at 19:51

How much nitrite is in that picture, woelen? It's a little hard to get an idea of how yellow the crystals without knowing how much crystal matter there is to scatter the light.

woelen - 6-1-2006 at 05:35

The vial, containing the NaNO2 has a size of approximately 5 cm (2 inch) from bottom to top of the black cap and the diameter of the vial is approximately 2.5 cm (1 inch). This should give you a sufficient indication of the amount of NaNO2 in the vial.
The glass of the vial is very thin, less than 1 mm.

Magpie - 6-1-2006 at 14:52

I performed another confirmatory test today, i.e., pH. The pH of my NaNO2 was 9; that of my NaNO3 was 7. This agrees with the MSDSs.

I tried again to get the faint blue color of nitrous acid by acidifying some dissolved NaNO2 with HCl. But it was still clear. I wonder why? I have seen nitrous acid at work: when dissolving uranium in nitric acid we absorbed the generated NOx in water. This water had a beautiful faint blue color which we attributed to the presence of nitrous acid.

woelen - 8-1-2006 at 12:32

Try with cool dilute H2SO4 or cool dilute HNO3 instead of HCl. HCl is reduced quickly and Cl2 is formed in small amounts. With dilute H2SO4 and HNO3 you don't have that problem. Also, the liquid should be cool, but not freezingly cold.

Try another nice experiment, forming HNO3, but not nitric acid, but one of its isomers, HOONO (peroxynitrous acid). That acid is red.

You can make it by first making a solution of HNO2 by mixing some dilute HNO3 with NaNO2. This must be ice cold. Then add a few drops of ice cold 10% H2O2. The liquid turns orange, but this orange color quickly fades. The HOONO quickly rearranges to HONO2.

garage chemist - 8-1-2006 at 12:47

Just add a few drops of 20% HCl to a spatula of your solid NaNO2 product. If there are appreciable amounts of NaNO2 in it, there will be a vigorous reaction, liberating lots of brown gas (NO2, from initially produced NO by reaction with atmospherical oxygen).
This is the easiest and most reliable test for nitrite. Nitrate gives absolutely no reaction here.

An equally good test for nitrite is to add a solution of an ammonium salt, like NH4Cl or NH4NO3 and warm gently.
Nitrite ions readily oxidise ammonium ions to elemental nitrogen:

NH4+ + NO2- -----> N2 + 2 H2O

The reaction takes place in dilute aqueous solution on gentle warming.
The reaction goes to completion. If an excess of ammonium salt is added to the test substance, all of the resulting gas is collected and the volume measured and the reaction allowed to get complete by warming to near the boiling point (cool down before measuring the amount of gas! Otherwise your measurement will be way too high, due to expansion of the gas and water vapor), a quantitative measurement of the percentage of NaNO2 in your product is possible.

Esplosivo - 8-1-2006 at 13:44

Originally posted by garage chemist
Just add a few drops of 20% HCl to a spatula of your solid NaNO2 product. If there are appreciable amounts of NaNO2 in it, there will be a vigorous reaction, liberating lots of brown gas (NO2, from initially produced NO by reaction with atmospherical oxygen).
This is the easiest and most reliable test for nitrite. Nitrate gives absolutely no reaction here.

Not to miss is also the very evident blue colour of the HNO2 (during addition of HCl to NaNO2) if the NaNO2 solution is concentrated enough. The 'strength' of the blue colour can be used to approximate conc. of NaNO2.

garage chemist - 8-1-2006 at 14:08

I meant solid NaNO2 with 20% HCl dropped on it, this way the HNO2 decomposes to NO and makes a brown cloud.
I have pure NaNO2 at hand, I'll test if I can make a blue solution of HNO2, because I think that a rather strong solution of NaNO2 is required for this!

Another thing:
Magpie, it is possible that some soluble lead compounds have formed during the Pb reduction of the NaNO2. I have a preparation where this is mentioned and the lead is removed with a special process.
To a solution of your substance add a solution of sodium sulfate, white insoluble PbSO4 will precipitate if there are lead compounds in it.

And the pH value says NOTHING about the identity and purity of your product! Every tiny contamination, like decomposition of the NaNO3 or NaNO2 with loss of nitrogen (a bit of this always happens) can strongly raise the pH.

The color is also worthless as an indicator of purity. I have fertilizer grade NaNO3 that is yellowish when powdered and might be mistaken for NaNO2.

During the reduction of NaNO3 with lead, efficient stirring is of utmost imortance, as is strong heating. Some preparations call for lead powder, as this gives a better reaction rate.

Purification of NaNO2 can be achieved by making some isopropyl nitrite with it (nitrate doesn't react here), isolating and purifying the ester and saponifying it by refluxing with alcoholic NaOH solution. NaNO2 separates right out, as its solubility in alcohol is small (3g/100ml in anhydrous ethanol).

chemoleo - 8-1-2006 at 14:43

Yes, putting some NaNO2 into 20%HCl shold work best, because of NOx evolution.
Having made some organic nitrite recently, the blue colour is evident by dripping HCl into NaNO2(aq), but at least in the presence alcohol it disappears wihtin seconds. A faint blue hue should be seen right away, even at lower concentrations, and the smell is very NOx like.

Anyway, Magpie, I'd probably repeat the whole thing, but this time with a large excess of Pb!

GC- I like your alkylnitrite-based purification!

Magpie - 8-1-2006 at 14:54

Garage chemist I placed a spatula of NaNO2 in a small beaker and dropped one drop of 6N (20%) HCl on it. There was an immediate and strong reaction giving a brown gas. A control test on NaNO3 gave no reaction. So I do indeed have NaNO2. The only question is of what purity.

I will test for Pb. I will also do more testing for HNO2 blue color and report back. Thanks to all for all the help. It really is an interesting compound. :D

chochu3 - 11-1-2006 at 23:45

I know that every one should know the following reaction:

Pb + KNO3 -> KNO2 + PbO

the reaction was done using 100 ml crucible, all reagents put inside and melted . Once it started to melt it was stirred then allowed to react one more minute and after stirring it was covered with a lid. It was allowed to raect for five minutes until reaction was stopped. Some of the substance was put into HCl acid and fumes of brown NO and NO2 was given off.

The next method I will try will be ammonium nitrate with spongy coppper to form ammoinuim nitrite.

heres a pic of the substance it was moslty yellow from lead oxide and some of the lead turned a redish color from over oxidation with air.

[Edited on 14-1-2006 by chochu3]

HPIM0522.JPG - 32kB

woelen - 12-1-2006 at 04:38

Originally posted by chochu3
The next method I will try will be ammonium nitrate with spongy coppper to form ammoinuim nitrite.

Forget about that, it will not work. Ammonium nitrite is very unstable and in the solid state it quickly decomposes to water and nitrogen gas. When heated, I can even imagine that it decomposes very violently. In aqueous solution, NH4NO2 also is quite unstable. In previous posts it was pointed out already that ammonium and nitrite ions react quantitatively to form N2 gas.

It might be that mixing KNO3 and Cu metal gives KNO2 and CuO, but that should be a matter of trying. Of course, there also will be side reactions and you get Cu(2+) ions in your final product. Cu(2+) ions form a green complex with nitrite and I can imagine that it is hard to separate this, harder than lead.

chochu3 - 12-1-2006 at 05:16

yes woelen but if the nitrite is formed add sodium hydroxide to precipitate sodium nitrite.

Basset, H. and R. G. Durrant. 1922. J. Chem. Soc., 121, 2631

This is the original article where ammonium nitrate was reduced by spongy copper to ammonium nitrite. It should be done in the absence of air and spongy copper as copper wire will not reduce the substance to the nitrite ion.

unionised - 26-1-2006 at 10:26

I saw a reference today to spongy cadmium as a reductant for NO3- to NO2-

Given the toxicity of Cd I'm not sure this is useful but it sugests that there may be other selective reductants that work well in solution.

woelen - 26-1-2006 at 10:43

Originally posted by chochu3
yes woelen but if the nitrite is formed add sodium hydroxide to precipitate sodium nitrite.

Basset, H. and R. G. Durrant. 1922. J. Chem. Soc., 121, 2631

This is the original article where ammonium nitrate was reduced by spongy copper to ammonium nitrite. It should be done in the absence of air and spongy copper as copper wire will not reduce the substance to the nitrite ion.

Ah, that sounds interesting. I did not know this. I once tried the combination ammonium/nitrite and that does decompose quickly on heating, but under the special conditions you mention, of course things may turn out different.
Even then, however, I think it will be quite hard to do this in practice, with materials, available for the general public. If you obtain any results, I would be pleased if you post them here, just out of curiousity.

unionised - 29-5-2006 at 09:54

It seems that there is an OTC way of reducing nitrate in solution
Not sure it's a lot of use, perhaps best for purification of nitrite made by a more conventianal pathway.

garage chemist - 29-5-2006 at 16:39

Sounds interesting, but I doubt that the reduction to nitrite will be clean at higher concentrations. Reduction of nitrates mostly gives several different products.
Also, the method introduces Zn ions into the solution. In neutral pH, Zn(OH)2 will be produced and stay in the reducing column, clogging it.

guy - 1-1-2007 at 17:45

Can KNO3 decomposotion be catalysed by MnO2 similarly to KClO3 decomposition?

Aqua_Fortis_100% - 1-1-2007 at 20:36

I used to prepare some DDNP and works, but still never tried to CTMTNA or other nitrosamine synthesis...maybe in next holidays!
My nitrite is very impure (theses days i no have ..i need make more) , because is made from "purified" double nitrates(50% NaNO3 ,50% KNO3 with sulfates ,carbonates and others impurities as coating,leaving the fertilizer looks as small pinkish "prills" from fertilizer called here "salitre do chile , sal duplo de sódio e potássio") which i only "purified", withdrawning the coating by dissolution/filtration (the coating isn't soluble or then weak soluble), and evaporation of filtrated liquid, leaving a white powder composed of KNO3 , NaNO3 and some impurities... but i no realized the fractional crystalization because for my homemade chemistry, i think, is sufficient ... :D

well,aprox. 2 months ago i tried 4 methods :
conventional nitrate/Pb
heating nitrate alone

i'm away from heating nitrate alone and Ca(OH)2 methods because all times which i tried ,i failled...

first the Pb/nitrate.. my favorite method for make nitrite at the moment..
using heat resistent gloves, put 25g of double nitrates in improvised crucible (a tomato can with wire support), then added 65g of lead (excess caulated of aprox. 15g.. and this Pb is from fishing weights). heat on improvised alcohol burner and constantly stired with a old and useless spoon by 1 hour... tired for the heat, wait the melt mix turns a "cake" and breaking this cake in pieces first with own spoon after with a clear hammer.. dissolved and washed this yellow/brown granules/powder in 2 portions of 150ml of hot water each , and then filtered this.. remains in filter white powder of lead oxides and unreacted Pb in "drops" or balls formates.. evaporated the solution in a enameled iron pan, (when almost all of the liquid evaporated the solution turned yellow INTENSE and i haved problems when drying the nitrites ...)... ready! my calculated theoretical yield was 20,1g of "double nitrites" , but i got aprox. 15g of crude product.. for me isn't to bad...
heating alone the nitrate is a very BAD method, because after which i just looked in some books and ,of course, web, the temperature needed for decomposing nitrate into nitrite and O2, is almost the temp. to decomposing also the nitrite..but i'm not totally sure of this..maybe has any secret in procedures.. :P
i put 100g of pure expensive NaNO3 (from food store), in same improvised crucible and heating on same improvised alcohol burner by 90 minutes..BAH! only some small bubbles in melt NaNO3 and no visible change of colour(stays white), but when i keep the white cake becomes solid and crushing small pieces and put in HCl nothing happens!!!! what about?
the sugar/nitrate ,firts i observed which when i burned unmeasured nitrates/sucrose (for some smoke ..)in a metal can , after burns , stays in can small amounts of yellow cake. i always asked me what is this.. then when i searched more for nitrites i looked in E&W ,some things and looked the following comment in a topic created by kingspaz:

I've always mixed a small amount of sugar with KNO3, without melting or even powdering them. The amount of sugar must be more then what is used for cheap rocket fuel...

Then ignite. It won't exactly burn, it'll smoke alot, boil, and produce sparks and heat.
If it fails to sustain itself, use more sugar. If it ends up with alot of black crud, use less sugar.

It's crude, but it works, it's always done the job for me.

meselfs, E&W member... topic there

so i to find the reason which turns yellow! the nitrite! actually i find which several reductible substances can forms nitrites if oxodized by nitrates.. this info procced?
well, the fact was which i put some piece of the yellow cake after burn in a glass cup with common muriatic acid, then the piece instantly starts to bubble and give brown/red gas! i think which is good method for amateurs as i, but should has a lot of impurities..(maybe oxalic acid, etc)..

finally, the Ca(OH)2/nitrate/graphite method...( also discussed in same topic by Polverone and kingspaz which tried also and have success with the synthesis..)
from the equation:
2NaNO3 + Ca(OH)2 + C ---> 2NaNO2 + CaCO3 + H2O
2 KNO3 + Ca(OH)2 + C ---> 2 KNO2 + CaCO3 + H2O ..(CaCO3 insoluble)
i tried 3 times, in similar attempts..
in all attemps i used same amounts:
60g of (Na,K)NO3 (excess) , 5g of graphite powder (bought in hardware store as lock lubrificant) (excess) and 20g of Ca(OH)2 (also from hardware store in cheap 7Kg bags containig Mg(OH)2 and others impurities)..
i started putting the double nitrate in improvised crucible on burner and wait a while for melt the nitrates, then i added slowly the Ca(OH)2 (which is difficult to handle in synthesis due the big volume of small amounts ,compared with the others substances used (C, (Na,K)NO3) ), then i put also slowly the graphite and aparentlly released SMALL amounts of water vapour.... the BIG problem mine is with Ca(OH)2: when i add the Ca(OH)2 ,it make the melted nitrates more hard to maintain in this state(more Ca(OH)2, more difficult to melt the mixture), and even on the burner it's becomes solid!!
what i must to modify??? maybe mixing first the Ca(OH)2 and graphite, but my afraid is which the nitrate oxodizes the graphite before reacting with Ca(OH)2!!!(??) adding water!?!?
well, if i to get found a place which sell me diluted HNO3 , i will try the HNO3/starch/NaOH method... others interestings things which i wish try: nitrate/NaOH/C method (the advantage which not has the problems of calcium hydroxide, but the disadvantage is which formed Na2SO4 should be quite hard to remove by crystalization and reduces the nitrite yield ..)
frogfrot , you tried to remove KNO2 from your Potassium dichromate synthesis ???
( just: Cr2O3 + 3KNO3 + 4KOH --> 2K2CrO4 + 3KNO2 + 2H2O ) is much hard separate acetates and the "other compounds in negligible concentrations"(extracted from link) from KNO3 ???

[Editado em 2-1-2007 por Aqua_Fortis_100%]

[Editado em 2-1-2007 por Aqua_Fortis_100%]

guy - 7-3-2007 at 23:17

36g KNO3 mixed with 9.8g of sucrose and roughly mixed togther. This was put in an steel bowl and heated, and lots of H2O vapor came off and all carbon reacted and gave a yellowish-greenish liquid which forms a yellow solid. About a gram of the solid gives a medium brown color but had a lot of CO2 bubbling.

Conclusion: not a good way to make very good yields of nitrate. However the reaction is very fast (about 1 minute).

obsessed_chemist - 23-3-2007 at 06:54

Originally posted by chochu3The next method I will try will be ammonium nitrate with spongy coppper to form ammoinuim nitrite.

I have read that ammonium nitrite can be acheived by reacting hydrogen peroxide with ammonia. I will try to find the reference.

If the solutions were pre-chilled, and added slowly with much stirring, and tempature closely monitored, perhaps the yield would be decent. This could then be reacted with sodium hydroxide to yield sodium nitrite and ammonium-water, which could be mostly boiled off, and then the sodium nitrite solution chilled and mixed with alchol to precipitate it. Any thoughts?

[Edited on 3/23/2007 by obsessed_chemist]

obsessed_chemist - 23-3-2007 at 07:33

quick update; just as an experiement I tried slowly mixing portions of both room temperature household ammonia (3%) and hydrogen peroxide (3%). I figured that the reaction stoichiometry is approximately a 1:2 ratio (for the solutions) of ammonia:peroxide, assuming that a reaction would proceed as follows:

4NH4OH + 7H2O2 ---> 2NH3NO2 + 14H2O

No heat nor bubbling evolved, and it still smelled a bit like ammonia. Then again, most ammonium salts smell like ammonia, so a reaction could be deceiving. Still, the ammonia scent lingered. My guess is that some gentle heating might be worth a try in the future. Or maybe bubbling gaseous ammonia through an ice-bath chilled 30% peroxide solution would yield good results. But at that point, it would no longer be OTC.

On a seperate note, has anyone tried mixing NaNO3 and Pb powder in a ball mill, than shielding the reaction mixture with an inert gas like argon?

[Edited on 3/23/2007 by obsessed_chemist]

Levi - 24-3-2007 at 05:54

This patent claims to use peroxide to oxidize ammonia to hydroxylamine.

mick - 1-4-2007 at 14:00

The closest I got to ammonium nitrite was attempting to recycle some ethanol / water containing ammonium nitrite as a by-product. This was at work and the idea was to recover the solvent. On my first run I stopped the 1 L flask distillation with about 150 ml in the flask. When a bit of air got into the flask the reaction started and the residue went like rocket fuel. I found out afterwards ammonium nitrite does not exist as a (stable) solid. The couple of 45 gal drums containing the stuff was still stored at the back of the place when I left.
Closest I ever got to blowing my self up.

Filemon - 14-8-2007 at 17:10

Do you know some nitrite salt that is insoluble in H2O? I know the Co(NO2)2. The Fe(NO2)3, Fe(NO2)2, Al(NO2)3, are they insoluble? Can somebody check it?

not_important - 14-8-2007 at 17:55

Aluminium nitrite may not exist. Most simple nitrites are fairly soluble, the low solubilities seem to come with complex salts.

Cu(II) forms a low solubility complex nitrite, given variously as 2CuO . N2O3, and Cu(NO2)2 . 3Cu(OH)2

Mixed salts listed as slightly soluble to very slightly soluble include Ca(NO2)2.Ni(NO2)2.2KNO2 ; CaCuK2(NO2)6 ; CaNiK2(NO2)6 ; FePbK2(NO2)6, NiPbK2(NO2)6.

Filemon - 16-8-2007 at 15:00

Originally posted by Theoretic
nitrate. It should be reduced like so: 2NaNO3 + S => 2NaNO2 + SO2. SO2 should be led into Na2CO3 solution (not, NaOH, for reasons of greater availability and lesser alkalinity, which could otherwise induce suckback). Na2SO3 thus formed could be reused for reduction.

I have proven mixing Na2CO3 and KNO3 in H2O ammonia it has taken place, if it is not ammonia it looks like much. I toss sulfur and it didn't react, it continued taking place ammonia. What reaction does it take place? Then, I proved with KNO3 and NaOH when I toss sulfur it reacted being dissolved in a product red-orange. What substance is it? It is not K2SO4 neither K2SO3.

guy - 16-8-2007 at 15:31

Originally posted by Filemon
Originally posted by Theoretic
nitrate. It should be reduced like so: 2NaNO3 + S => 2NaNO2 + SO2. SO2 should be led into Na2CO3 solution (not, NaOH, for reasons of greater availability and lesser alkalinity, which could otherwise induce suckback). Na2SO3 thus formed could be reused for reduction.

I have proven mixing Na2CO3 and KNO3 in H2O ammonia it has taken place, if it is not ammonia it looks like much. I toss sulfur and it didn't react, it continued taking place ammonia. What reaction does it take place? Then, I proved with KNO3 and NaOH when I toss sulfur it reacted being dissolved in a product red-orange. What substance is it? It is not K2SO4 neither K2SO3.

Na2CO3 and KNO3 does not make NH3.

Sulfur dissolves in NaOH to make polysulfides which are usually red or yellow.

Filemon - 16-8-2007 at 16:22

Originally posted by guy
Na2CO3 and KNO3 does not make NH3.

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

Polverone - 16-8-2007 at 17:23

Originally posted by Filemon
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?

Aqua_Fortis_100% - 17-8-2007 at 04:26

chochu3 , i've tried similair procedure , but in a new batch...

600g of double saltpeter ("50% NaNO3 - 50% KNO3) plus about 1,3 Kg of lead fishing weights , heated in a paint can for several hours (heated by wood :) ) and well stirred over the time.
this gave very similair hard pieces that i havent still stracted.. many lead seemed to be unreacted, so as soon as possible i will add more lead and remelt to see improvements in yields... i expect that this amount of nitrite last for many time...

these pieces seems to be extremely hard.. the clean hammer was my choice in early batchs..this allow better extraction with hot water

(the earlier batch , after purification, also gave a reddish cloud when added to muriatic acid , and also reacted with urea to gave , thought, some nitrogen )

[Edited on 17-8-2007 by Aqua_Fortis_100%]

kmno4 - 17-8-2007 at 11:05

I used to make about 500g pure NaNO2 by NaNO3/Pb reduction (about 10 runs, heh). It is "nice and easy" reaction but....
1. Always remains some unreacted Pb even if reaction lasts several hours and more than 1 mole of Pb per 1 mole of NaNO3 is used
2. Molten mixture has to be very often stirred.
3. Always remains some soluble Pb salts in extracted melt [Pb(NO2)2 or some NO2(-) complexes of Pb - I do not know - is intensively yellow-orange. Try reaction, for example, Pb(NO3)2 and NaNO2 in water] and they should be removed.
4. Reaction yelds some Na2O (?) and extract has to be neutralised with dilute HNO3 [or CO2 gas] drop by drop. During carful neutralization, all Pb precypitates as some bacis insoluble salts [or just Pb(OH)2] and in one step solution is neutralised and freed from Pb.
5. To prevent plaing with all mass of solidificated melt, it is much better to transfer it ,when still molten, on sheet of steel (or something like that) by spoon, in a few portions and let to cool down.
6. Playing with Pb - it is not good and reaction leaves a lot of useless (or not, heh) Pb/PbO/Pb34O mixture. I still have it - few kilograms in a glass jar.
Besides, there are a lot of books describing these procedures.
NaNO2 content is easy to determinate using KMnO4/H2SO4.
(add slowly [with stirring], NaNO2 solution to KMnO4/H2SO4 sol. untill it becomes colourless)
... but much better is to buy it (as I did later ;))

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