Sciencemadness Discussion Board

NaNO2 failure

Wolfram - 22-11-2003 at 06:11

Airgun pellets were melted. (I assumed that they consisted of lead.)
The metal block formed was filed down with a file.
50g of the metal powder was slowly added in smal portions to melted 20g melted NaNO3. To my surprice the "lead powder" was actually floting on the melted NaNO3. When the termprature raised even higher some brown salt looking like common rust began to form.
Could it be SnO? What did the pellets consist of, they were heavy and the metal was soft and melting at about 200-300.

(The reaction I was trying to perform was:
Pb + NaNO3 + heat -----> PbO + NaNO2)

MAke sure you have good lead

chloric1 - 22-11-2003 at 08:44

Well that lead in those pellets might have been alloyed with antimony, tin, or ??? First I would try the lead sinker weights for fishing lines and if you still can't get respectible yields go to McMaster Carr on the net and they have lead in any form you want.

Incidently, a few years back I feel in love with the potassium nitrate granulated sugar mix. I mixed 6parts KNO3 and 4 parts sugar and the smallest glowing ember would send the comp instantly into a whitish purple flame and a hissing sound with grey smoke. LOVED IT! Anyways when it was all said and done the salt had melted into a yellow mass with readily yielded Nitrogen Oxides with HCL and H2SO4. Problem is the purity may not be too great but it would separate from the burnt carbon junk by dissolving in cold water.

Another way to nitrites / Just a thought

Mephisto - 22-11-2003 at 15:05

The reduction of an alkali nitrate with lead has the disadvantage, that the product is polluted with nitrate. And the separation of nitrate and nitrite is (relatively) difficult. Although most chemists here haven’t a problem to buy pure nitrite, the unavailability of an OTC-method to synthesise pure alkali nitrites is annoying.

My idea is to reduce a strong alkali nitrate solution with zinc (in presence of acetic acid). Than the zinc-ions can be removed with ammonia by forming insoluble zinc hydroxide, which can be filtered of. The remaining acetate and ammonium-ions can be removed by heating (the alkali nitrite shouldn’t be destroyed by this).

The product should be quite pure, but the zinc hydroxide precipitation might be difficult, because it is dependent to the pH of the solution (so too much ammonia will dissolve the precipitation again).

Microtek - 23-11-2003 at 08:00

In my opinion, the most convenient method, at least in theory, is reduction with charcoal in the presence of Ca(OH)2. The only products are NaNO2, NaNO3, CaCO3 and Ca(OH)2 so you can cycle your product through the process a few times and should then be left with only NaNO2.
I think mr anonymous posted about this method a while ago, and I have also tried it myself. I found that the molten nitrate/nitrite becomes so thick when adding the Ca(OH)2/C mix that the melt puffs up and has a tendency to burn partially. The process can still be used, but I think it works better if you don't add quite so much reducer in each cycle.

Could it possibly be..

Wolfram - 23-11-2003 at 11:20

Could it possibly be so that PbO2 or Pb3O4 formed? It was a reddish-brown unsoluble heavy junk.

unionised - 23-11-2003 at 11:57

That's an expensive synthesis of washing soda. Sodium nitrite oxidises carbon (I just tried it).

Real powdered Pb still ..

Wolfram - 27-11-2003 at 03:12

Real powdered Pb still gives me redbrown-color salt instead of pale yellow wtf is this? :mad:

Microtek - 27-11-2003 at 09:12

Unionized: I don't think so. The procedure comes from an old patent that is an improvement of an even older common method for producing nitrite. The improvement consisted of using Ca(OH)2 instead of NaOH which means that the only solubles are the product and the unreacted reactants.
It is essential to use only enough carbon to reduce the nitrate to nitrite; that is one half mole of carbon per mole nitrate. The carbon is added slowly, and diluted with Ca(OH)2 to the melted NaNO3.

unionised - 28-11-2003 at 12:30

Cycling the product through the reduction stage repeatedly will reduce all of it.
With a mixture of nitrate and nitrite the nitrite will melt at a lower temp (OK I realise there will be a eutectic somewhere) and will diffuse through the material and react with the charcoal before the nitrate.
I realise the method has been patented. The patent office don't check if a method works before they issue a patent.

The colour of lead's oxides will vary with the particle size and the crystal form as well as the oxidation state.
generally PbO exists as litharge (red), stable at room temp. On heating you can get massicot (yellow) and stable above 488C
Pb3O4 (red lead)
PbO2 (brown whenever I have seen it, but sometimes decribed as maroon) and also a high pressure form that is black.
There are also other oxides like Pb12O19, but that's just getting silly.

[Edited on 28-11-2003 by unionised]

Microtek - 29-11-2003 at 08:05

Why would the nitrite react before the nitrate ? As I see it, the method using carbon is no different from using lead or aluminum to reduce the nitrate, except that much less carbon is needed.
Besides, as I said, the patent is just a variation on an established theme ( heating NaNO3 with alkali and carbon ).
In fact I think the Muspratt book which Polverone has on his website used this process for nitrite production.

Look at this:

Wolfram - 29-11-2003 at 14:14

Look at this:

, it states that:

KNO3 + 2C ---> KNO2 + 2CO

So maybee one could use it for NaNO3 also..?

unionised - 30-11-2003 at 03:48

Let me now when anyone gainsays the point that repeated reduction with an excess of carbon will reduce the nitrite too.

Generally, solids don't react well with one another because they can only mix at the points of contact. Liquids will react wth solids rather better. First one to melt is the first one to react.
Sodium nitrite starts to decompose, even when it is on its own, when it is about 10 C above the mpt of NaNO3. (Merck)
That means you need good temperature controll. That's a bit hard to achieve given that the reactions are exothermic.
I'm not saying you can't get any NaNO2 this way. I'm saying that repeated reduction will ruin the yield unless you can separate out the nitrite betwen runs.

Microtek - 30-11-2003 at 04:12

That is the point I've been making all along; you need a stochiometric amount of carbon, but since the addition of all the C+Ca(OH)2 in one batch makes the melt unstirrable, I recommend adding perhaps one third of that ( in little portions ) to the melted NaNO3, then cooling, dissolving in water, filtering off the CaCO3 and any Ca(OH)2 and carbon, evaporating the water and then do the same with the rest of the C+Ca(OH)2.

unionised - 30-11-2003 at 13:37

And the thermolysis of the nitrite?
What guarantee is there that, having been freshly generated in the presence of the hot charcoal, the nitrite will (because it has read a patent) defer to the (possibly still solid) nitrate, get out of the way, and let it take its turn?
(Merck gives the decomp temps for the K salts and the nitrite is easier to decompose this tends to indicate that it will react quicker too)

Microtek - 1-12-2003 at 07:17

If that was the case, the nitrite would be decomposed by all of the reactions which involve the melting of a nitrate. I think the rate of thermal decomposition is actually quite low for both KNO3 and KNO2. This is based on my attempts at producing nitrite by the supposedly common laboratory practice of heating a nitrate. I heated it strongly; way above the melting point but gas evolution was very slow.

So go and try.

Wolfram - 1-12-2003 at 09:38

So please someone go and try:

NaNO3 + 2C ----> NaNO2 + 2 CO

use active carbon not barbecue cole. ;)

..come back and report the results.

What is the meaning with involving

If anyone doesn´t try it I will, but it will take some time becouse I have much to do and have no activated carbon.

[Edited on 1-12-2003 by Wolfram]

unionised - 1-12-2003 at 13:41

I know it' not the same thing but I did the experiment with charcoal (activated as it happens) and NaNO2.
It flashed in much the same way as C and NaNO3.
That's why one of my earlier posts says "I just tried it"
The question people now need to address is why would this reaction not go in the presence of hot NaNO3.
Please let me know.

Microtek - 2-12-2003 at 03:37

I have tried it. With NaNO3, ordinary charcoal from wood, and Ca(OH)2. I think the reason that alkali is inkluded is to moderate the reaction so it doesn't flash, but decomposes in a controlled manner. The product evolves nitrous fumes when reacted with acid, but I haven't done any quantitative experiments to determine the amount of nitrite in it.


Wolfram - 2-12-2003 at 05:27

Wtf in chemistry you become surpriced almost every day. I have thought that PbO is yellow. Now I read that there are two forms of PbO; yellow AND red!
This explains why the Pb powder got red-brown in my experiment.

[Edited on 2-12-2003 by Wolfram]

JDP - 2-12-2003 at 14:53

I believe the red vertion is Lead Tetroxide(Pb3O4). I've done the reaction of NaNO3 + Pb + heat be for and I got a fine tan to white pewdered PbO.

unionised - 2-12-2003 at 15:35

How does the NO3- or NO2- know that the Ca(OH)2 is present? (it's barely soluble in the melt.)
If it is un-aware of this it will react as I observed.
Would anyone who wouldst reply to this please address this and my other questions rather than saying "its in a patent" or "it makes a difference", without explaining why.

Microtek - 4-12-2003 at 03:22

One of the essential principles in science, is that no amount of theorizing can undo the findings of an experiment.
So, if an experiment has been conducted ( and I have conducted one ) that shows that it works, then that isn't changed by whether you or I can understand how. Besides you haven't offered any solid evidence that the nitrite "should" be reduced before the nitrate either.
Saying that nitrite has a lower decomposition temperature than nitrate is not sufficient to conclude that it must be reduced faster. As I have said before, if nitrite decomposed faster than nitrate, it wouldn't be possible to produce it by heating NaNO3.
As for the mechanism of moderation by Ca(OH)2, I can offer one which is nothing more than a plausible guess:
I you have a melt of NaNO3 and you pour powdered charcoal into it, the C will be oxidized by the nitrate just as in blackpowder. And just as in blackpowder, the heat of reaction will accelerate the oxidation beyond control.
Now, if you add something like Ca(OH)2, the released energy will be absorbed by the inert ( in this context ) matter and so reduce the reaction rate to a controllable level.
Whether this is the actual mechanism or not ( and I doubt that it is ) it demonstrates that additives doesn't have to be in solution to affect the reaction.

TrollchEmist - 4-12-2003 at 05:29

Did this reaction:

NaNO3 + C ----> NaNO2 + CO

Powdered NaNO3 and stoichiometric amount of active charcoal dust was mixed and then melted. Mixture was black-grey coloured. When melted some bubbles evolved and then one spark(then I moved bit further and watched the fun part.) Huge blaze evolved from mixture. After reaction mixture colour was changed to pale yellow-white. There was only minimal amount of charcoal left.

Seems a good produce for NaNO2:cool:


Wolfram - 4-12-2003 at 07:15

Congratulations!!! :D
Did you really set it on fire? How long did the fire last?
If you would have a good scale you could messure if the corect mass has become gas...
But you can instead buy a rat and check the effect. The LD dose for a human (75Kg) is about 3g. So for a 500 g rat it should be 3/150=0.02 g. im joking but if would be nice to know if its the nitrite you have.

[Edited on 4-12-2003 by Wolfram] :)

[Edited on 4-12-2003 by Wolfram]

TrollchEmist - 4-12-2003 at 10:17

Yes it blazed really furiously about 10s or something(it was beautiful flame(height about 30cm and bright).

here the calculations:

58g NaNO3 and about 8g C
that would make a 0,68mole reaction

then (after filtering) getting NaNO2 powder(very white with pale yellow) about 31g. that's little bit less than expected but some of the blazing mixture propably splashed away( there was amounts of that around heating place.

Then tested the pH of my 'product' in water and that was definitely base(pH paper)

The NaNO2 solution in water is weak base

Then I did little sulfuric acid test with my 'product' and NaNO3 side by side. (dropped little both of them into sulfuric acid) The 'product' evolved some gases(nitrogen oxides) while NaNO3 was still. After couple minutes 'product' was dissolved and NaNO3 was still.

Broked my thermometer so i haven't done melting point test.

'Product' is definitely NaNO2

This is unbelievable easy synthesis of NaNO2:D

[Edited on 4-12-2003 by TrollchEmist]

Very nice

Wolfram - 4-12-2003 at 18:24

Very nice but I assume that you better have a god ventilation doesn´t there come a huge cloud of smoke that fills your kitchen from the reaktion?:)

[Edited on 5-12-2003 by Wolfram]

I found this..

Wolfram - 4-12-2003 at 20:24

1 Cr2O3 + 2 Na2CO3 + 3 NaNO3 = =>
2 Na2CrO4 + 3 NaNO2 + 2 CO2

[Edited on 5-12-2003 by Wolfram]

TrollchEmist - 4-12-2003 at 21:35

Yes you're right about huge smoke and it's better to do this reaction outside or have a very good ventilation. CO is toxic gas (takes oxygens place in hemoglobin) so be careful.

Microtek - 5-12-2003 at 02:43

The problem with this violent reaction is that, while some nitrite probably has been produced, a lot of the NaNO3 has probably reacted something like this:
4NaNO3 + 5C --> 5CO2 + 2Na2O + 2N2
and then Na2O + CO2 --> Na2CO3
which is alkaline and gives bubbles when acid is added.
One way to check for carbonate is to add a little CaCl2 or Ca(NO3)2 in water to some of your product in water solution. If carbonate is present, you will see a precipitate ( CaCO3 ) after some stirring.

TrollchEmist - 12-12-2003 at 11:22

I agree with you. Did one synthese with 'propable' NaNO2 and yield was very low.

but there were NaNO2 present(tiny amounts) reaction is too violent with only charcoal.
sorry :(

Have to try different methods.

guaguanco - 12-12-2003 at 15:14

If you have a source of NO2...

TrollchEmist - 13-12-2003 at 08:16

Yeh. This is what I'll try:

One need Pure nitric acid and Cu powder and NaOH or Na2CO3 Reaction is this:

Cu + 4HNO3 --> Cu(NO3)2 +2NO2 + 2H2O

if used concentrated nitric acid. reaction would be:
3Cu + 8HNO3 --> 3Cu(NO3)2 + 2NO + 4H2O

So drop HNO3 on Cu powder and bubble formed NO2 gas trough NaOH or Na2CO3 water solution. I assume that NaNO2 would form.

This looks promising, I'll try it soon and report.( maybe after christmas)

[Edited on 13-12-2003 by TrollchEmist]

unionised - 13-12-2003 at 13:57

Would someone care to balance the equation
NO2 + Na2CO3 --> CO2 + Na NO2

And do be careful disolving the gas, if it sucks carbonate solution back into the acid you will have problems.

guaguanco - 13-12-2003 at 18:28

Originally posted by unionised

Would someone care to balance the equation

NO2 + Na2CO3 --> CO2 + Na NO2

And do be careful disolving the gas, if it sucks carbonate solution back into the acid you will have problems.

Good point, it's probably NO that is the required oxide.

unionised - 14-12-2003 at 04:20

Would someone care to balance the equation for NO please?

Microtek - 14-12-2003 at 05:48

Na2CO3 + NO + NO2 --> 2NaNO2 + CO2

I have tried this using 50% HNO3 and starch to produce the N2O3 gas ( the equimolar mix of NO and NO2 acts as if it is N2O3 ). As HNO3 is consumed and the concentration drops, a greater proportion of NO is produced. If only NO2 is used, there will be produced some nitrate ( but using K2CO3 should make that easy to remove )

pyroscikim - 29-12-2003 at 00:42

Try sulphur as reducing agent for nitrate salt, as i doubt it reduces nitrites.

just add it slowly to molten nitrate to form nitrite and sulphur dioxide which vents off. the excess sulphur will melt and form a lump which is easily removed upon cooling. It also ensures that all nitrate reacts.

and isn't this a discussion to make *pure* nitrite?

KABOOOM(pyrojustforfun) - 30-12-2003 at 20:20

Inert gas is needed as autoignition T for S is 232°C
some polynitroaromates hydrolyze in alkali conditions forming nitite salt and the nitrogroup replaces with a -OH or with a -NH<sub>2</sub> if ammonia is used (plus ammonium nitrite)
of course it is not worthy as a way to produce nitrites but as a byproduct of preparation of some explosives
eg preparation of TATB from pentanitroaniline and ammonia should produce 2 mol ammonium nitrite per PNA. and preparation of styphnic acid from tetranitroaniline produces 1 mol nitrite per mol of TNA
H<sub>2</sub>NC<sub>6</sub>(NO<sub>2</sub>;)<sub>5</sub> + 4NH<sub>3</sub> <s>&nbsp;&nbsp;&nbsp;></s> TATB + 2NH<sub>4</sub>NO<sub>2</sub>
H<sub>2</sub>NC<sub>6</sub>(NO<sub>2</sub>;)<sub>4</sub>H + 3NaOH <s>&nbsp;&nbsp;&nbsp;></s> Na<sub>2</sub>(O-)<sub>2</sub>C<sub>6</sub>(NO<sub>2</sub>;)<sub>3</sub>H + NaNO<sub>2</sub> + NH<sub>3</sub> + H<sub>2</sub>O
the nitrogroup which is replaced is the one which doesn't agree with other groups about the position it activates (it's the way I construed for myself)

Ramiel - 1-1-2004 at 02:25

pyroscikim, I tried your method. It is pretty labor and time intensive.
I don't know if I'm doing it improperly, but as I add the Sulphur, The molten Pot. Nitrate solidifies and cakes up in weird ways - preventing any furthur reaction.

In addition, I mixed an excess of sulphur to pot nitrate and then ignited it (by reaching it's flashpoint and, in a seperate reaction, by adding burning coke) with similar complications.

From experimental experience, the low yeilds of NaNO<sub>2</sub> from the carbon-type reaction are fairly acceptable once you take into account the ease of the reaction.

perhaps some investigation into a wet reaction would pay dividends?

p.s. woof, I'm still a bit light headed from all that SO<sub>2</sub>! it's a bit like being drunk so I apologise if the post is hard to follow in places. :cool:

In fact, my brain insisted that I take a lie down mid-experiment. Is this bad? :D

unionised - 3-1-2004 at 12:50

Sulphur reduces sodium nitrite on heating.

Nick F - 7-1-2004 at 10:07

What about :

Getting hold of some CaO - either buy it, or make it from limestone or whatever. There are many options for getting this.
Get some sulphur, and place in a two-holed container. The sulphur is ignited, and dry air is pumped into one hole. The gas coming out of the other hole is passed over your CaO, forming CaSO3.
This is mixed with your nitrate, and the mixture is fused. After the reaction is over, cool the mixture, dissolve in water, filter off CaSO4, crystalise, recrystalise, dry. You should have quite pure XNO2.
It's a slight modification of one proccess that passes SO2 over a fused mixture of CaO and XNO3. I just think it'd be easier if you didn't have to handle high temperatures and toxic gases at the same time.
Actually, you might be better off making the SO2 from sodium metabisulphite (wine/brewing suppliers...) and a strong acid, preferably H2SO4 (conc.) since this would produce dry SO2. Then you don't need to pump any air. Or you could get it by reducing H2SO4 with zinc. Again, there are many possibilities. I think I'd prefer the sodium metabisulphite route actually.
SO2 absorption by dry solid CaO might be quite slow, in which case you could bubble the SO2 through a suspension of Ca(OH)2 in water, and dry the CaSO3 afterwards.
I think I'd use a method like that if I needed to make some.

Edit: if you were feeling adventurous and had a good freezing mixture to hand, you could pass your SO2 into a container of CaO that had been chilled to about -20*C. The SO2 would condense and soak into your CaO, which might give you a better reaction.

[Edited on 7-1-2004 by Nick F]

another rejected method

Polverone - 7-1-2004 at 14:04

I buy NaNO2, but it's something that I'd like to be able to make should I ever need to. I wondered if tin metal (somewhat expensive, I know) might be able to reduce NaNO3 to NaNO2. I used lead-free solder that was mostly tin, along with a little silver and other undisclosed components. Conclusion: it is an effective reducing agent, but sometimes too effective. I overheated the test tube that had my molten NaNO3 and solder in it, and there was a flashing/glowing ignition in the tube that spewed smoky NO2 out of the tube and melted a hole in the bottom. On another attempt I heated the two together without an ignition, and the strong evolution of NO2 on acidification showed that nitrite was definitely present. But this method seems about as tricky as using molten lead to effect the reduction, and tin's more expensive (and easier to start fires with).

I found a very brief reference in some reaction reference manual that heating a neutral solution of KNO3 with zinc dust would reduce it to KNO2. No details were offered, and my couple of attempts were unsuccessful. However, when I boiled NH4NO3 solution with lead dust some time ago, I did form some lead nitrite. I could tell from the coloration of the solution and the evolution of NO2 on slight acidification.

Edit: reading a post by Organikum elsewhere makes me wonder if my zinc dust is too pure to effect the reduction of MNO3 to MNO2 in neutral aqueous solution. The original reference was from the 1880s.

[Edited on 1-7-2004 by Polverone]

jimwig - 9-2-2004 at 14:56

Just for the record.

Melted lead (unknown purity) over wood fire in large kitchen spoon. Added some agricultural grade NaNO3 to molten lead.
Presently a reaction happened- bubbles etc. The spoon contained a water soluble yellowish powder/crystals along with other gunk.

Pb + NaNO3 = NaNO2 + PbO (?)

I have not tried this again but will soon.

I want this to work with charcoal but just agricultural grade charcoal. Activiated is available but precious for other non-substitute type reactions.

jimwig - 9-2-2004 at 17:37

the history of nitrite production is summarized in <A HREF="">US4312722</A&g t; (patent)

[Edited on 2-10-2004 by Polverone]