Sciencemadness Discussion Board

Very OTC Sodium NItrite

Σldritch - 29-9-2017 at 11:40

I found a really nice nitrite synth on atomistry:
The Sodium nitrite, NaNO2, is obtained by reducing sodium nitrate with metals such as lead or iron, with sulphur or carbon, or with material containing these substances. In Dittrich's process the nitrate is heated with slaked lime and sawdust, the yield being almost quantitative, where as the action of coal and charcoal is too energetic:


Im very suprised this has not been posted before. I have read about people here attempt something similar with charcoal but getting bad results. It seems any reducing agent with a base will reduce nitrate to nitrite.

I did not have any sawdust nor calcium hydroxide on hand so i tried it with flour and sodium hydroxide drain cleaner:

12 NaNO3 + 12 NaOH + C6H10O5 = 12 NaNO2 + 6 Na2CO3 + 11 H2O

I heated 115g of Sodium Nitrate with the stochiometric ammounts of the other reagents, mixed sloppily together in a steel can with plenty of head room, on a burner. A lid was loose placed over it. As it heated up it started smoking and foaming. The reaction seemed to stop after the smoking did and i was left with a light yellow mixture at the bottom of the can.

Afterthe cake had cooled it was crushed up and dissolved in boiland water and cooled in a fridge. The mixture froze to a slush. It then filtered it on a vacuum pump until mostr of the slush had melted. (The slush was probely a mixture of sodium carbonate and some sodium nitrite hydrate). The mixture was then boiled down to the theoretical volume of a satured solution containing the theoretical amount of sodium nitrite formed in the reaction and chilled and filtered again.

Then i thought i would neutralize the residual sodium hydroxide so i added sodium bicarbonate repeated the filtering again. I really should have thought about that earlier but oh well.

I evaporated the final solution on a steam bath and tested with hydrochloric acid. It seems i got a relativly high yield even with extra steps though it would probely be pointless to weigh.



This seems WAY better than using lead or sulfur or some other obscure ways to reduce nitrite. And with calcium hydroxide you could make it even simpler. Nitrite can be made dirt cheap with this method im sure.

I really recommend this, im going to make some isopropyl nitrite now aga ;)

[Edited on 29-9-2017 by Σldritch]

UC235 - 29-9-2017 at 15:58

http://www.sciencemadness.org/talk/viewthread.php?tid=52

Separation of unreacted nitrate is very difficult, and some brown gas on acid addition is hardly a good way to quantify such a mixture. Converting the crude mix to a water-immiscible nitrite ester is probably among the best approaches.

clearly_not_atara - 29-9-2017 at 17:30

This is a pretty cool idea. Not sure why lime would be simpler... the insolubility of resulting CaCO3 might make it difficult to extract the product.

Solubility at 0 C:

KNO2: 280 g / 100 mL
KNO3: 13 g / 100 mL

This is large enough that a saturated solution prepared from the potassium salts at 0 C should have a sufficiently high proportion of nitrite for practical purposes.

If nitrite free of nitrate is desired, such as for the preparation of N2O3, this reaction can be used:

NiCl2 (aq) + 6KNO2 (aq) >> K4Ni(NO2)6*H2O (s) + 2KCl (aq)

"Potassium nitrite (80 g in 25 ml water) was addedwith brisk stirring to NiCl2*6H20 (20 g in 20 ml water). The crystalline precipitate was filtered, washed with cold methanol, and dried in the air (yield 88%)"
http://www.publish.csiro.au/CH/CH9731663

Potassium hexanitronickelate monohydrate precipitates as orange-brown crystals which dry to a violet solid when heated under vacuum. However, I am not sure if this solid can be used to generate N2O3 by rxn with acids (although I suspect the answer is "yes" so long as the acid is strong enough, eg H2SO4).

Σldritch - 30-9-2017 at 01:28

Quote: Originally posted by UC235  
http://www.sciencemadness.org/talk/viewthread.php?tid=52

Separation of unreacted nitrate is very difficult, and some brown gas on acid addition is hardly a good way to quantify such a mixture. Converting the crude mix to a water-immiscible nitrite ester is probably among the best approaches.


Misremembered the solubility of sodium nitrate... Anyway if the reaction was not complete i would not expect a white cake after the reaction but a grey one. Of course the nitrite might react faster than the nitrate with the carbon so it does not guarantee purity but nitrate as the stronger oxidizer should react first.

Ill try making a alkyl nitrite soon but i have guests now so it will have to wait.

brubei - 30-9-2017 at 03:21

Sodium Nitrate is commonly sold for ceramic making.

Sulaiman - 30-9-2017 at 04:23

OTC sodium nitrite is legal and easy in UK http://www.ebay.co.uk/itm/1kg-Sodium-nitrite-high-quality-/3...

do I get a prize ? :P

XeonTheMGPony - 30-9-2017 at 05:09

Only if he'll ship to Canada!

Pulverulescent - 30-9-2017 at 08:49

Quote:
I found a really nice nitrite synth on atomistry:
The Sodium nitrite, NaNO2, is obtained by reducing sodium nitrate with metals such as lead or iron, with sulphur or carbon, or with material containing these substances

Talking about ebay, you could exploit your easy prep. there ─ and, er, cash-in? :cool:

symboom - 30-9-2017 at 08:55

NiCl2 (aq) + 6KNO2 (aq) >> K4Ni(NO2)6*H2O (s) + 2KCl (aq)
Potassium hexanitronickelate
Interesting I wonder if using sodium nitrate with kcl would work

clearly_not_atara - 30-9-2017 at 09:11

Quote: Originally posted by Σldritch  
Ill try making a alkyl nitrite soon but i have guests now so it will have to wait.
Just use the potassium salts, it's much easier.

In particular, K2HPO4 is much more soluble (140% w/w) than Na2HPO4 (7% w/w), so a saturated solution of the former should precipitate the latter. KHCO3 has this property to a lesser extent.

Salt metathesis is highly underrated, I see.

Σldritch - 30-9-2017 at 09:54

I have a hard time getting potassium salts. The best i can get is 40%KCl 50%NaCl 10%MgSO4 mineral salt. I have had no sucess separating the KCl. I have a little bit of KNO3 left bought from a now closed down store though. Id rather use a small excess of flour.

Im pretty confident that there is not much nitrate in it because if there was the nitrite would have to react faster than the nitrite in the molten mixture. If that was the case almost no nitrite would be formed at all. It makes sense too since nitrate is a stronger oxidizer in these conditions. (Ex. permanganate and ferrate)

clearly_not_atara - 30-9-2017 at 12:46

They don't sell cream of tartar where you live? Burning this salt gives potash.

The Volatile Chemist - 30-9-2017 at 16:13

Quote: Originally posted by clearly_not_atara  
They don't sell cream of tartar where you live? Burning this salt gives potash.

cream of tartar is expensive OTC and burning it would reduce it to even less mass. Not an economical method.

clearly_not_atara - 1-10-2017 at 14:34

By this method you would be paying about $35/kg for potash with Amazon prices. It's not cheap but considering the costs associated with performing any kind of amateur chemistry it doesn't sound expensive unless you're making silly amounts of nitrites (and for what?). OTOH K2SO4 is available on Amazon for less than $10/kg but maybe he can't buy that.

I'm a bit surprised you couldn't just buy some kind of potassium fertilizer, given it's one of the big three plant minerals. Or you could always do this:

NaNO3 (aq) + KCl (aq) [0 C or lower] >> KNO3 (s) + NaCl (aq)

OP apparently has access to KCl salt replacement so this should work, although it requires redoing the whole process.

[Edited on 1-10-2017 by clearly_not_atara]

[Edited on 1-10-2017 by clearly_not_atara]

j_sum1 - 1-10-2017 at 20:18

I'm with Sulaiman on this one. In my world, sodium nitrite is OTC. And easier to obtain than nitrates.

http://www.ebay.com.au/itm/100g-bag-of-Sodium-nitrite-100-Fo...
http://www.melbournefooddepot.com/buy/sodium-nitrite-powder-...


[edit] typo


[Edited on 2-10-2017 by j_sum1]

What is OTC varies a lot

Σldritch - 1-10-2017 at 22:59

I would call something OTC when i can buy without giving my credit card imformation. That includes most stores where i live.

I can buy sodium nitrates in about half of the grocery stores here yet potassium salts are hard to obtain. I think this is intentional as terrorism preventation, i doubt it is very effective though. I have not found any stores that sell sodium nitrite.

Anyway, i dont have a lot of calcium hydroxide which you would need do do this with potassium because of the similar solubility of potassium nitrite and potassium carbonate. If you do not use calcium hydroxide it is more of a tradeoff between carbonate and nitrate impurities.

Mineral salt would probably introduce more impurities than it would help reduce.

Maybe i can titrate it? I tried permangante and ammonium chloride, neither seemed to work very well.


Also i enjoy the challenge of obtaining chemicals i can not buy.

clearly_not_atara - 1-10-2017 at 23:27

You wouldn't find potassium fertilizer in a grocery store. (Neither would I) You have to go somewhere that sells fertilizer. It'd be at a hardware store for me -- the same place that sells tools and paint (Home Depot). They have giant bags marked "Potassium Sulfate" with a K2O percentage marked. You would also see it at a garden supply.

It's really strange to ban potassium because it makes a pretty crappy bomb. It doesn't react and it's heavy which reduces the temperature of mixtures containing it. And it's extremely useful as fertilizer. It's one of the most legitimate chemicals I can think of -- everyone from subsistence farmers to yuppies washing their faces with Dr. Bronner's uses potassium.

But it tastes bitter and it's not a common ingredient in food, so you won't usually see it at the store.

Broken Gears - 2-10-2017 at 05:54

Quote: Originally posted by brubei  
Sodium Nitrate is commonly sold for ceramic making.


Sodium Nitrite should be easy to find OTC, as it's used in preserving meat.
Any Home-produktion, hunters/butchers shop or DIY beef jerky shop should have it OTC.

karlos³ - 2-10-2017 at 17:56

Isn´t the stuff used for preservation of meat, "curing salt", just like 1% NaNO2 at most, the remainder being NaCl, NaNO3 for the major part?
At least in europe the curing salt contains per law less than one percent, usually.

Melgar - 2-10-2017 at 23:26

Potassium is just less common than sodium in general, and the salts are correspondingly less common. Interestingly though, plants contain almost no sodium at all, which is why we prefer the taste of sodium salt on our food, and why deer are attracted to salt licks, etc. Humans and other animals NEED sodium in our diets. Plants were able to evolve to not need sodium at all, and instead are able to extract potassium from feldspar (the most common mineral in the world), which binds potassium very tightly, but not sodium. Sodium, on the other hand, has mostly all been leached out of the ground and into the oceans eons ago.

The chemistry of feldspar is pretty neat too. Its name means "not ore" or something like that in German, and it mostly contains silica. Aluminum has a similar atomic radius as silicon, and can fit into a silica matrix, but then it has that missing electron that messes up the matrix. Unless, of course, potassium sits next to it and lends it its extra electron. Then everything works out great, since potassium fits quite well into that matrix too. And it means that terrestrial plants have access to an alkali metal as well, because this planet would be a totally different color if there wasn't one available to them.

Of course, our taste for sodium, combined with the ease of mining it from old dried up seabeds, has meant that we tend to consume it preferentially over potassium. However, it's not actually bad for us and doesn't contain any calories, so there's no reason to stop, as long as we're getting enough of all our other minerals too. Sodium nitrite though, is a whole different story. Even though it's food grade, it's a known carcinogen. However, the FDA has determined that at the very low levels that it's used in meat as a preservative, that the benefits of not getting food poisoning outweigh the (very low) risk of developing cancer from consuming it.

AJKOER - 3-10-2017 at 11:00

Here is a new approach based on my prior attempts (see http://www.sciencemadness.org/talk/viewthread.php?tid=52&... ) that I will hopefully be able to test soon.

First, as occurs in the case of the metal Aluminum (see, for example, equation (3.7) in a doctoral thesis from 2008, "Alkaline dissolution of aluminum: surface chemistry and subsurface interfacial phenomena", by Saikat Adhikari, link: https://www.google.com/url?sa=t&source=web&rct=j&...), I would argue similarly with either aluminum or zinc, the creation of the metal hydroxide directly from the metal, proceeds with the release of electrons per the reaction

Al + 3 OH- → Al(OH)3 + 3 e-
Zn + 2 OH- → Zn(OH)2 + 2 e-

Then, a possible reaction in the presence of nitrate, with either prehydrated or totally solvated electrons, being reported as readily scavenged by nitrate:

e(p)-/e(aq)- + NO3- + H2O -> NO2 + 2OH- (Source: see eq. (5) in JAERI-Conf 95-003, "5. 6 Radiolysis of Concentrated Nitric Acid Solutions R. Nagaishi" by P.Y. Jiang, et al, link: https://www.google.com/url?sa=t&source=web&rct=j&... )

Upon shaking the solution periodically, likely containing NO2 gas, in an atmosphere of pure oxygen also (see below):

2 NO2 + H2O --> HNO2 + HNO3

Implying a net reaction of in the case of Aluminum of aqueous nitrate in an alkaline solution:

2 Al + 3 NO3- + 3 H2O -- 6OH- --> 2 Al(OH)3 + 3 NO2-

Do not use more aluminum then needed. But if in excess, expect:

e(p)-/e(aq)- + NO2- + H2O -> NO + 2OH-

and shaking with O2:

2 NO + O2 --> 2 NO2

2 NO2 + H2O --> HNO2 + HNO3
-----------------------------------------------

Possible other reaction of interest would be the formation of the superoxide radical anion (from O2 + e-), which could readily react with any formed NO, creating peroxynitrate that would be converted back into nitrate.

I also suspect the reaction may proceed well with NH4NO3.

[Edited on 3-10-2017 by AJKOER]

Σldritch - 3-10-2017 at 12:41

Quote: Originally posted by AJKOER  
Here is a new approach based on my prior attempts (see http://www.sciencemadness.org/talk/viewthread.php?tid=52&... ) that I will hopefully be able to test soon.

First, as occurs in the case of the metal Aluminum (see, for example, equation (3.7) in a doctoral thesis from 2008, "Alkaline dissolution of aluminum: surface chemistry and subsurface interfacial phenomena", by Saikat Adhikari, link: https://www.google.com/url?sa=t&source=web&rct=j&...), I would argue similarly with either aluminum or zinc, the creation of the metal hydroxide directly from the metal, proceeds with the release of electrons per the reaction

Al + 3 OH- → Al(OH)3 + 3 e-
Zn + 2 OH- → Zn(OH)2 + 2 e-

Then, a possible reaction in the presence of nitrate, with either prehydrated or totally solvated electrons, being reported as readily scavenged by nitrate:

e(p)-/e(aq)- + NO3- + H2O -> NO2 + 2OH- (Source: see eq. (5) in JAERI-Conf 95-003, "5. 6 Radiolysis of Concentrated Nitric Acid Solutions R. Nagaishi" by P.Y. Jiang, et al, link: https://www.google.com/url?sa=t&source=web&rct=j&... )

Upon shaking the solution periodically, likely containing NO2 gas, in an atmosphere of pure oxygen also (see below):

2 NO2 + H2O --> HNO2 + HNO3

Implying a net reaction of in the case of Aluminum of aqueous nitrate in an alkaline solution:

2 Al + 3 NO3- + 3 H2O -- 6OH- --> 2 Al(OH)3 + 3 NO2-

Do not use more aluminum then needed. But if in excess, expect:

e(p)-/e(aq)- + NO2- + H2O -> NO + 2OH-

and shaking with O2:

2 NO + O2 --> 2 NO2

2 NO2 + H2O --> HNO2 + HNO3
-----------------------------------------------

Possible other reaction of interest would be the formation of the superoxide radical anion (from O2 + e-), which could readily react with any formed NO, creating peroxynitrate that would be converted back into nitrate.

I also suspect the reaction may proceed well with NH4NO3.

[Edited on 3-10-2017 by AJKOER]


I doubt you will get much other than ammonia similar to how you will not get alcohols from clemmensen reduction; the reactant is bonded to the metal surface until it has picked up hydrogen on all its bonds to the metal so the reaction will not stop at nitrite very often.


If you did it without a solvent it would probably just explode so that is not an option though i think chemplayer did something like this and failed in a less dramatic way:

https://www.youtube.com/watch?v=w9nhdpKhztI&t=1s

I think lead is about the strongest reducing agent that will work. Other reducing agents that work such as polysulfides and carbon are weaker reducing agents and both of them react way too fast without something to slow them down. Maybe it depends on the solubility of lead oxide in sodium nitrate/nitrite too. I doubt aluminium oxide is very soluble but zinc oxide might be.

Then there is this too, i do not know how he/she did the nitrite test but supposedly there is a lot of nitrite in it. NOt sure if if was a commercial test solution or the ferrous sulfate test though. If it was the ferrous sulfate test than there was not a lot of nitrite it seems.

https://www.youtube.com/watch?v=5Sgd1wjpywc

If you really want to use a metal then i think you should look for one less reducing like bismuth.

argyrium - 3-10-2017 at 13:18

Am I the only one who spotted this error in the original post - or am I in error?

"Then i thought i would neutralize the residual sodium hydroxide so i added sodium bicarbonate repeated the filtering again. I really should have thought about that earlier but oh well."

??

AJKOER - 3-10-2017 at 13:48

Quote: Originally posted by Σldritch  

.........
I doubt you will get much other than ammonia similar to how you will not get alcohols from clemmensen reduction; the reactant is bonded to the metal surface until it has picked up hydrogen on all its bonds to the metal so the reaction will not stop at nitrite very often.



I believe based on your comment is that it is necessary to have some excess in nitrate, at least the required amount of say NaOH, and less than indicated amount of Aluminum, plus some oxygen. The complete dissolution of the metal surface by the NaOH is likely needed. If not, likely reaction with water on the aluminum could lead to further reduction:

3 [ H2O = H+ + OH- ]

Al + 3 OH- → Al(OH)3 + 3 e-

3 e- + 3 H+ = 3 .H

.H + NO3- = OH- + .NO2

2 .NO2. + H2O = HNO2 + HNO3

.H + NO2- = OH- + .NO

.......

Physical removal of undissolved aluminum metal (actually Al foil, an alloy, with Fe,..., which could act as a galvanic couple producing reducing e- ) may be needed, else Al sitting in a slow reaction with water, or per other reactions, may be an issue. Aqueous nitrite are photoactive producing hydroxyl radicals, which could lower yield with time, so avoid prolonged light exposure.

[Edited on 4-10-2017 by AJKOER]

AJKOER - 5-10-2017 at 14:05

OK, I ran an experiment per my reaction (actually, cited in the literature under alkaline conditions, see, for example, [EDIT] "Nitrate Removal from Ground Water: A Review", by Archna, et al., E-Journal of Chemistry, 2012, 9(4), 1667-1675), link: https://www.google.com/url?sa=t&source=web&rct=j&... ):

2 Al + 3 NO3- + 3 H2O -- 6OH- --> 2 Al(OH)3 + 3 NO2-

I dissolved 7 g of KNO3 in 12O cc of distilled water. I added 1.4 g of aluminum foil (a sheet of 14 cm x 23 cm). Used an excess of NaOH (5 cc). ([EDIT] As a old prior source spoke of a pH around 9.0, a better procedure may be to add the stoichiometric dose of NaOH over divided doses, with stirring and stopping before if all the aluminum is dissolved). All the aluminum dissolved. Left a fine black suspension of which I was able to filter most out of the very alkaline solution. See picture of pre-filtered solution below:

Added the hopefully now nitrite rich mix with added sea salt to 97% ethanol ([EDIT] Ever Clear, not Evergreen), and currently awaiting sunlight to breakdown the alcohol (smell change) via photolysis of aqueous nitrite/sea salt (reference: please see http://onlinelibrary.wiley.com/doi/10.1029/JC086iC04p03173/a...). Not a classic test for nitrite, but it is one of my intended uses (for photolysis).

20171005_100153-640x480.jpg - 31kB

[Edited on 5-10-2017 by AJKOER]

[Edited on 6-10-2017 by AJKOER]

AJKOER - 5-10-2017 at 14:11

Photolysis run:


20171005_141947-480x640.jpg - 46kB

Melgar - 5-10-2017 at 19:14

Quote: Originally posted by AJKOER  
OK, I ran an experiment per my reaction (actually, cited in the literature under alkaline conditions, see, for example, file:///home/chronos/u-6092dab7e8781d5c630e3fdaff87bc2dff6db2e0/Downloads/154616.pdf ):

2 Al + 3 NO3- + 3 H2O -- 6OH- --> 2 Al(OH)3 + 3 NO2-

Have you realized that that's a link to a file on your local computer? It seems to indicate that you have a Unix-like filesystem, and that your username (or the system name) is "chronos". Possibly on a public computer, since your files are in a folder with what appears to be an MD5 hash in the username, and may be a way of allowing guest users to save files locally.

Quote: Originally posted by AJKOER  
I dissolved 7 g of KNO3 in 12O cc of distilled water. I added 1.4 g of aluminum foil (a sheet of 14 cm x 23 cm). Used an excess of NaOH (5 cc). All the aluminum dissolved. Left a fine black suspension of which I was able to filter most out of the very alkaline solution. See picture of pre-filtered solution below:

Added the hopefully now nitrite rich mix with added sea salt to 97% ethanol (Evergreen), and currently awaiting sunlight to breakdown the alcohol (smell change) via photolysis of aqueous nitrite/sea salt (reference: please see http://onlinelibrary.wiley.com/doi/10.1029/JC086iC04p03173/a...).]

Evergreen? You sure you don't mean "Everclear"? Maybe it's time to get some sleep now, eh?

XeonTheMGPony - 6-10-2017 at 03:35

or stop taste testing the ever clear for potency!

AJKOER - 6-10-2017 at 04:40

Melgar:

Fixed the link on another machine and inserted article title.

Thanks. I was using an alternate computer (Acer Chrome book). Apparently, just copying the url as displayed on that machine for certain links (like to locally stored downloaded files) is problematic for the other computers. Lots of new things with the Chrome book computer got to get acquainted with, but it does have a low price, large screen and even HMDI ports to play online movies onto big screen TVs,...... Recommend it for word processing (talk and it enters your text fairly accurately based on context), research,..., but not for anything like online games and such.

I don't drink the alcohol, else I would at least known what to call it if I have to buy more!

Cheers!

[Edited on 6-10-2017 by AJKOER]

Magpie - 6-10-2017 at 07:39

NaNO2 can be bought very cheaply at Ace hardware where salts are used to brine salmon eggs.

[Edited on 7-10-2017 by Magpie]

AJKOER - 6-10-2017 at 08:02

Pure KNO3 is sold as stump remover aid in stores with home garden sections (Home Depot,...).

AJKOER - 6-10-2017 at 13:47

Updated picture following photolysis in sunlight for 6 hours:

The reaction mix is now more intensely colored (resembling olive oil) together with a diminished smell from the former strong scent of the EverClear.

Some photochemical reaction, in alkaline conditions, has apparently occurred, which may be supportive of the claim of the initial nitrite presence given the short time frame of treament. The latter with sea salt, alcohol and distilled water in the presence of strong sunlight, may have produced hydroxyl radicals, as would be expected per my prior cited source, thereby further producing new products. Definitely, no smell of NH3.

20171006_174244.jpg - 344kB

[Edited on 7-10-2017 by AJKOER]

Σldritch - 9-10-2017 at 09:17

Tried preparing Isopropyl nitrite from the nitrite i made. Yield was 30%.

Melgar - 9-10-2017 at 13:44

Quote: Originally posted by AJKOER  
Photolysis run:



Nitrates and nitro groups are rarely very active at all toward reduction in strongly alkaline solutions. I'm pretty sure that a H+ ion would be necessary for reducing NO3-, and those are hard to come by in a solution that alkaline. Not to mention, aluminum would be acting as an acid, and forming aluminate salts with your alkalis. I think that in the reaction you cited, what must be happening is that eventually aluminum neutralizes the pH, at which point it may be possible for it to reduce nitrates selectively, since the aluminum/aluminate would be able to buffer the pH. But since your solution was strongly alkaline, I'd expect that you still have nitrates, rather than nitrites. You can always test by adding a strong acid and checking for brown fumes, which would mean nitrite. I suspect you don't actually have any though.

Rhodanide - 10-10-2017 at 05:53

Quote: Originally posted by j_sum1  
I'm with Sulaiman on this one. In my world, sodium nitrite is OTC. And easier to obtain than nitrates.

http://www.ebay.com.au/itm/100g-bag-of-Sodium-nitrite-100-Fo...
http://www.melbournefooddepot.com/buy/sodium-nitrite-powder-...


[edit] typo


[Edited on 2-10-2017 by j_sum1]



RIGHT?!
I can buy NaNO2 by the POUND, but NaNO3 is IMPOSSIBLE to find!!! Or any Nitrate for that matter, besides NH4NO3 from instant cold packs.

AJKOER - 12-10-2017 at 03:24

My claimed alteration of NH3 generation is cited as likely correct (see reaction 1.6 below). Here is an extract from a source, page 1.12, "Mitigation of Hydrogen Gas Generation from the Reaction of Water with Uranium Metal in K Basin Sludge", by SI Sinkov, et al, January 2010, to quote:

"2 Al + 2 NaOH + 6 H2O → 2 NaAl(OH)4 + 3 H2

The evolution of H2 was moderated by the addition of NaNO3 to the cladding removal solution to form ammonia. The chemical reduction of the nitrate to ammonia occurs by the following stoichiometry:

8 Al + 5 NaOH + 3 NaNO3 + 18 H2O → 8 NaAl(OH)4 + 3 NH3 Reaction 1.5

With higher sodium nitrate concentrations, ammonia decreases and NaNO2 is favored:

2 Al + 2 NaOH + 3 NaNO3 + 3 H2O → 2 NaAl(OH)4 + 3 NaNO2 Reaction 1.6

Systematic study of the effects of NaOH concentration and the NaNO3:Al ratio were undertaken to optimize the cladding removal process to minimize H2 release and decrease the unwanted production of NH3 (Gresky 1952). The reactions showed reasonable adherence to stoichiometry, as the NaNO3:Al ratio was varied, particularly at lower ratios. However, as shown in Figure 1.4, the release of NH3 could not be completely supplanted by NaNO2, even at high NaNO3:Al mole ratios.

Testing also showed that NaNO3 concentrations above ~1 M (85 g NaNO3/liter) had little further effect in decreasing the H2 yield (Figure 1.5). At high NaNO3 concentrations, the H2 yield was ~2 mL of gas (~8.3×10-5 moles) per gram (3.7×10-2 moles) of aluminum or 2.2×10-3 moles of H2 per mole of Al. This is about 0.15% of the 1.5 moles H2 per mole of Al yield that would have occurred in nitrate-free alkaline solution or an attenuation factor of 1/0.0015 (~670).
.......
The joint evolutions of H2 and NH3 were found to be at a practical minimum under plant conditions when the nitrate and aluminum mole quantities were nearly equal (Gresky 1952):

20 Al + 17 NaOH + 21 NaNO3 + 36 H2O → 20 NaAl(OH)4 + 18 NaNO2 + 3 NH3 Reaction 1.7"

Source link: http://r.search.yahoo.com/_ylt=A0LEV1L8gNxZJTAA.mnBGOd_;_ylu...

Note, my prior work above suggested a reaction of:
2 Al + 3 NO3- + 3 H2O -- 6OH- --> 2 Al(OH)3 + 3 NO2-

As compared to:
2 Al + 2 NaOH + 3 NaNO3 + 3 H2O → 2 NaAl(OH)4 + 3 NaNO2 Reaction 1.6"

[Edit] I have happily surprised that my reaction mechanics, attributed to likes of Mg, Al and Zn, apparently apply also to uranium, to quote from the same source, page 1.2:

"Uranium metal is highly electropositive, reacting with water to produce hydrogen radicals (H·) and UO2. The reactive hydrogen radicals can combine to form H2:

U + 2 H2O → UO2 + 4 H· → UO2 + 2 H2 Reaction 1.1

The H2 dissolves in water and, upon water saturation, forms bubbles that are released into the gas phase.

The hydrogen radicals or H2 also can react with uranium metal to form UH3:

U + 3H· (or 1.5 H2) → UH3 Reaction 1.2

The UH3 then can react with water to liberate hydrogen radicals or H2:

UH3 + 2 H2O → UO2 + 7 H· (or 3.5 H2) Reaction 1.3 "

[Edited on 12-10-2017 by AJKOER]

AJKOER - 14-10-2017 at 06:38

Quote: Originally posted by Melgar  

......
Nitrates and nitro groups are rarely very active at all toward reduction in strongly alkaline solutions. I'm pretty sure that a H+ ion would be necessary for reducing NO3-, and those are hard to come by in a solution that alkaline. Not to mention, aluminum would be acting as an acid, and forming aluminate salts with your alkalis. I think that in the reaction you cited, what must be happening is that eventually aluminum neutralizes the pH, at which point it may be possible for it to reduce nitrates selectively, since the aluminum/aluminate would be able to buffer the pH. But since your solution was strongly alkaline, I'd expect that you still have nitrates, rather than nitrites. You can always test by adding a strong acid and checking for brown fumes, which would mean nitrite. I suspect you don't actually have any though.


As I noted previously on page 1 of this thread, "possible reaction in the presence of nitrate, with either prehydrated or totally solvated electrons, being reported as readily scavenged by nitrate:

e(p)-/e(aq)- + NO3- + H2O -> NO2 + 2 OH- (Source: see eq. (5) in JAERI-Conf 95-003, "5. 6 Radiolysis of Concentrated Nitric Acid Solutions R. Nagaishi" by P.Y. Jiang, et al, link: https://www.google.com/url?sa=t&source=web&rct=j&... )"

which would seem to suggest a possible shift to say partially solvated electrons in place of .H in less acidic conditions as a path to aqueous NO2 (and some NO2- + NO3- therefrom).

My rough recollection of the literature was that the vigor of Al/NaOH reaction was possibly a factor in the effectiveness of any reductive process. This could imply that oxygen from air or dissolved in solution could be entering the reaction and producing the superoxide radical anion (or just referred to as superoxide), via:

e(p) + O2 = .O2-

Given the apparent affinity of superoxide with nitric oxide to form peroxonitrite in alkaline aqueous solution (see, for example, "Reaction of superoxide with nitric oxide to form peroxonitrite in alkaline aqueous solution", Inorganic Chemistry (ACS Publications), pubs.acs.org/doi/abs/10.1021/ic00216a003, by NV Blough (1985), http://pubs.acs.org/doi/abs/10.1021/ic00216a003 ), a further reaction may be occurring with the stable NO2 radical also, which I would state as:

.O2- + .NO2 = O2 + NO2- (Source: "Table 1: Initial Concentrations for three scenarios under polluted continental (urban), unpolluted continental remote)", R48 at http://www.google.com/url?sa=t&rct=j&q=e(p)-%2B%20NO3-%20%2B%20H2O%20%3D%20NO2-%20%2B%202%20OH-&source=web&cd=17&ved=0ahUKEwj_js-k nvHWAhUC4SYKHaP0Cr04ChAWCC0wBg&url=http%3A%2F%2Fprojects.tropos.de%2Fcapram%2Fcapram23.pdf&usg=AOvVaw1t2DhghrOHnYru_phBHNYk )

The net of the last three reactions could then be:

2 e(p)- + NO3- + H2O -- O2 -> NO2- + 2 OH-

which I have also seen reported in the literature (it is also a cited half cell reaction, see, for example, http://www.google.com/url?sa=t&rct=j&q=e-%20%2B%20NO... ).

In any event, the formation of a reductive species (.H or e-(p) ) appears to occur at both low and high pH.

[Edited on 15-10-2017 by AJKOER]

AJKOER - 14-10-2017 at 08:24

Some interesting observations from this 1921 paper (please ignore the theory), "THE MECHANISM OF REDUCTION OF NITRATES AND NITRITES IN PROCESSES OF ASSIMILATION.", by OSKAR BAUDISCH, 1921, link: http://www.google.com/url?sa=t&rct=j&q=THE%20MECHANI... . Some interesting comments to quote:

"This dissociation of nitrate into oxygen and nitrite can also be brought about by means of metallic iron as well as under the influence of the energy of light. If a neutral oxygen-free solution of potassium nitrate be shaken in a vacuum with active iron prepared by reduction with hydrogen, the supernatant liquor obtained after the iron powder has been allowed to settle will give every reaction applicable for the detection of nitrous acid. In other words, metallic iron will easily reduce potassium nitrate to potassium nitrite in the cold in the absence of every trace of oxvgen, .."

My take on using a boiled aqueous nitrate solution (removing oxygen) to which is added fresh iron filings in an a sealed O2 free vessel, as a possible path to nitrite:

2 [ H2O = H+ + OH- ]
Fe + 2 OH- → Fe(OH)2 + 2 e-
2 [ e- + H+ = .H ]
2 [ .H + NO3- = OH- + .NO2 ]
2 NO2. + H2O = 2 H+ + NO2- + NO3-

Adding reactions:
Fe + 3 H2O + 2 NO3- → Fe(OH)2 + NO2- + NO3- + 2 H2O

Upon cancelling, my estimate of the overall slow net reaction is (which implies equal moles of iron metal powder and an available nitrate):

Fe + H2O + NO3- → Fe(OH)2 + NO2-

Note, avoid an excess of iron metal and water as:
.H + NO2- = OH- + .NO
......

[Edit] In fact, a source notes the following:

"Nitrate reduction can be induced under basic pH according to the following reaction10:

3NO3- + 8Fe (OH)2 + 6H2O → NH3 + 8Fe(OH)3 + OH-

Experimental results showed that a Fe: NO3- ratio of about 15: 1 was required in the presence of copper catalyst for the reaction to proceed"

Source: "Nitrate Removal from Ground Water: A Review", by Archna, et al., E-Journal of Chemistry, 2012, 9(4), 1667-1675), link: https://www.google.com/url?sa=t&source=web&rct=j&...

A problematic side reaction is possibly the formation of hydrogen gas (which also suggests employing an expandable vessel to avoid spillage):

.H + .H = H2 (g)

Use of a Magnetizer may likely accelerate the reaction also (see https://www.sciencemadness.org/whisper/viewthread.php?tid=77...).

[Edited on 14-10-2017 by AJKOER]

Σldritch - 14-10-2017 at 09:10

1. Metal powders are hard to make and/or expensive.

2. The reactions requires an excess of nitrate which...

3. is hard to separate and...

4. produces lots of byproducts such as...

4. nitric oxide produced by the reaction of ferrous with nitrates/nitrites and ammonia. (http://pubs.acs.org/doi/abs/10.1021/ja01331a020?journalCode=...)

The carbohydrate-nitrate-base route was at least well established and used industrially for quantitative nitrite production. It really seems like the best route to make nitrite to me unless you just want to have fun with the metal reduction.

If you want really pure nitrite i think converting it to an alkyl nitrite and hydrolysing it will do it.

AJKOER - 14-10-2017 at 10:50

Quote: Originally posted by Σldritch  
1. Metal powders are hard to make and/or expensive.

2. The reactions requires an excess of nitrate which...

3. is hard to separate and...

4. produces lots of byproducts such as...

4. nitric oxide produced by the reaction of ferrous with nitrates/nitrites and ammonia. (http://pubs.acs.org/doi/abs/10.1021/ja01331a020?journalCode=...)

The carbohydrate-nitrate-base route was at least well established and used industrially for quantitative nitrite production. It really seems like the best route to make nitrite to me unless you just want to have fun with the metal reduction.

If you want really pure nitrite i think converting it to an alkyl nitrite and hydrolysing it will do it.


In my opinion, none of your cited points are valid for the claimed oxygen-free iron metal approach (granted, to be verified and not likely very large scale as keeping oxygen free is likely increasingly difficult upon scaling up).

1. I have prepared iron filings for experiments in under 5 minutes with a medium sized file acting on a cast iron rode.

2. My indicated net reaction indicates just 1 mole of nitrate to produce 1 mole of nitrite.

3. Fe(OH)2 is not soluble in near neutral conditions, so no separation issue.

4. No byproducts except perhaps a very small amount of H2 or NH3.

5. Normally, per your link, the reaction of a ferrous salt in highly acidic (not neutral conditions) acting on nitrate can lead to NO. Also, using a very high (15:1) Fe to NO3- ratio along with a copper catalyst (as I noted above in a reference) may enable reduction to ammonia. The latter reference applies to commercial nitrate removal from ground water.
---------------------------------------------------------

[Edit] My speculation as to why air/oxygen is such a problem, even in trace amounts, with respect to the Iron metal/H2O/Nitrate process:

First, we want H+ + e- = .H to take place.

But, O2 can steal the the e- forming the superoxide radical anion, .O2-, resulting in the loss of one potential .H

Also, the .O2- + .H = HO2- , resulting in the loss of an existing .H (source: see
https://images.search.yahoo.com/search/images;_ylt=AwrBT89Dq... ).

Also, in the presence of CO2 in air, creating soluble ferrous bicarbonate, we could have Fe(ll) + O2 = Fe(lll) + .O2- , which is the so called metal auto oxidation reaction, regenerating the superoxide to remove another .H

And finally, Fe(lll) + HO2- = Fe(ll) + H+ + .O2- (pH >4.8), which recycles any soluble ferric to ferrous (and also creates another superoxide), thereby resulting in a cyclic chain reaction consuming any created .H (or e-) reducing radicals.

[Edited on 15-10-2017 by AJKOER]

unionised - 14-10-2017 at 10:56

Quote: Originally posted by Melgar  
Quote: Originally posted by AJKOER  
Photolysis run:



Nitrates and nitro groups are rarely very active at all toward reduction in strongly alkaline solutions. I'm pretty sure that a H+ ion would be necessary for reducing NO3-, and those are hard to come by in a solution that alkaline. Not to mention, aluminum would be acting as an acid, and forming aluminate salts with your alkalis.

Nonsense.
https://en.wikipedia.org/wiki/Devarda%27s_alloy

Antiswat - 10-1-2018 at 04:18

iron oxalate, very low solubility in water, as oxalic acid is a strong acid pretty much any soluble iron salt you can find of iron will precipitate iron oxalate when mixing up solutions of oxalic acid and iron salt, its easily purified by decantation, barely takes minutes to settle, its decomposition point is around the temperature of melting point for sodium nitrate
mix the two and heat up, nanoiron will form, along with a bit of carbonate and hydroxide of iron oxides, the nanoiron should react so vigously that you wouldnt need to heat the mixture very forcibly, likely it would act slightly pyrotechnic even supplying itself with energy, in theory this works, in practice i have zero experience yet
starting materials are relatively easy to get around, procedure should be simple and isolation of reaction products should also be quite doable

sodium nitrite is sparingly soluble in ethanol, cant find much about sodium nitrates solubility in ethanol however, this may be exploitable.
sodium nitrate is insoluble in acetone, nitrite in acetone? i'd test solubilities out in common solvents but some very organized thieves got this idea that sodium nitrite is super valuable in producing explosives so they had to my neat little bottle of well labeled sodium nitrite.

barium nitrate-nitrite could be a plausible way to get rid of nitrate, although it comes close to solubility different of sodium nitrate-nitrite, ~5g Ba(NO3)2 vs 50g Ba(NO2)2 100mL @0*C

what using FeOx could offer would be a quite clean process without hazardous lead fumes, possibly some carbon monoxide and some iron, but likely dodging the brutal mess of charcoal, for dealing with alkali carbonate i'd suggest reacting the mess with HCl as NaCl has only a few grammes of solubility difference from 0-100*C making it ideal for fractional crystallization
on a sidenote IPN is quite worthy for FAE

GrayGhost- - 10-1-2018 at 09:04

I used Al flakes ( used to applied to face ) + potassium nitrate both dry, and firing. I was obtain many smoke and scrap.:(

Master of the Elements - 10-1-2018 at 18:07

It is sold at some sporting goods shops as a bait preservative, usually right next to the borax and sodium sulfite.

Antiswat - 4-1-2019 at 13:01

alright. i think ive just came across the ultimate preparation of nitrite, works with sodium nitrite. ive just tested it using dilute H2SO4 and a bit of the suspected material with a bit of iron sulfate solution. immediate black precipitate

its ridiculously simple: NaNO3 decomposes into mainly NaNO2 at temp of 300-500*C. a hotplate can easily reach glowing red hot temperatures, my infrared thermometer said hotplate was 450*C hot but having worked with steel i'd say easily 600*C
shortly after it melted an aroma of pyrotechnics was present, possibly from some NOx, i wouldnt call it NO2 however, Na2O was mentioned as one minor decomposition product in a study about it.
i reacted the finished product with H2SO4 and a bit of copper, around 10mL H2O to 1mL conc H2SO4, faint bubbling, no visible NO2
it was heated for something past 30 minutes in a small stainless steel tray, covered with aluminium foil
the aluminium foil was attacked by something, seemed very fragile but mostly intact, once it has been heated for a while the pyrotechnic smell is no longer present and this could be a hint at the reaction being mostly over
once the thing finished crackling a solid mass could be broken apart and chipped out

essentially you may be left with a bit of impurities, mostly NaNO3, but relatively pure from something as lazy as throwing it on a hotplate
i didnt measure out exact amounts before and after, finished product was quite dense, but the stainless steel appears to have been darkened greatly, most likely from either NOx or oxygen along with intense heating, nice matt dark.
low viscocity NaNO2 + ~800*C hotplate https://i.gyazo.com/bf38097573567358c1d5f48eb38b0934.jpg

i can see NaNO3 is insoluble in acetone, but i cant find anything on NaNO2, so for purification it may be easier to do it with potassium nitrate, as theres major difference in potassium nitrite/nitrate solubility in water, KNO2 is also soluble in ethanol where KNO3 isnt really.

tl;dr thermal decomposition of NO3 = NO2, KNO2 can be extracted with ethanol, KNO3 cant