Sciencemadness Discussion Board - Preparation of ionic nitrites

Preparation of ionic nitrites

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woelen - 16-4-2012 at 22:19

Trying to use iron in aqueous solution will not produce any nitrite. Nitrate ion is remarkably inert in aqueous solution, except in strongly concentrated acidic environments or at very high pH in combination with magnesium and aluminium. In the acidic environments it forms NOx with suitable reductors, in the strongly alkaline environment it forms ammonia.

I have no experience with mixing iron and KNO3 and trying to react that. Just try it on a small scale and report back on your results.

dann2 - 17-4-2012 at 13:28

From another thread.

Making Nitrite from Nitrate + Lead

Attachment: nitrite.zip (213kB)
This file has been downloaded 846 times

chochu3 - 17-4-2012 at 14:35

If ionic solutions are preferred some older text which I do not remember have stated it can be done with hydriodic acid and some other substrates, if I can find my journal with references I have other written formulas which I will post later.

AJKOER - 18-4-2012 at 07:59

Expanding on a previous reference to a 2009 Sciencemadness thread containing a reply by Madcedar entitled "NaNO2 from Ethanol, HNO3 and NaOH"
(https://www.sciencemadness.org/whisper/viewthread.php?tid=52... ), here is an old work dedicated to Ethyl nitrite. Reacting with KOH in alcohol, for example, yields KNO2.

Source: "Spirit of nitrous ether and ethyl nitrite: a bibliographical index of ...", by W. O. Richtmann, J. A. Anderson. Some extracts:

Link:
http://books.google.com.bz/books?pg=PA34&id=7yfrAAAAMAAJ...

[SEE PAGE 36]
"Persoz prepares ethyl nitrite by the action of pure fuming nitric acid upon absolute alcohol, well cooled, to prevent explosion. The ethyl nitrite is removed, and purified in the usual manner. 1862."

"The ethyl nitrite is decomposed by caustic potash (especially in alcoholic solution) and yields potassium nitrite and alcohol, proving that ethyl nitrite is a true "ether." Kopp, E. 1851."

Revue scientif., 27, p. 273. [Jour. de Pharm., 11, p. 320; Ann. d. Chem., 77, p. 332; Neues Jour. de Pharm, Jl, p. 320; Gmelin, vol. 8, p. 470; Jahresber. u. d. Ftschr. d. Pharmacog.,

2, p. 325; Am. Jour. Pharm., 59, p. 484.] Salpetrigsaures Aethyloxyd.

"Kopp prepares ethyl nitrite by using a mixture of alcohol, nitric acid, copper turnings, and passing the gas through a solution of ferrous sulphate to remove nitrous acid and other oxides of nitrogen. Gerhardt. 1851."

[SEE PAGE 35]
"Gmelin, vol. 8, p. 468. Nitrite of Ethyl.

History, synonomy, methods of formation, preparation—

1. Fuming nitric acid and alcohol, cold.

2. Fuming nitric acid, alcohol and water.

3. Nitric acid and alcohol, distilling.

4. Sulphuric acid, alcohol and nitrate.

5. Nitric acid, alcohol and reducing substances.

6. Nitrous acid on alcohol direct.

7. Sulphuric acid, alcohol and saltpeter (ord. temp.), purification, properties, decomposition and combinations.

Lea, M. C. 1861.

Sillim. Am. Journ. Sc, 32, p. 177. [Chem. Centrbl., 33, p. 688; Am. Journ. Ph., 34, p. 69; Proc A. Ph. A., 10, p. 160; Ann. der Pharm., 165, p. 58; Drug. Circ, 5, p. 222; Cans. Jahresb. d. Pharm., 1862, p. 182.] On the preparation of nitrate and nitrite of ethyl.

The reaction between nitric acid and alcohol in preparation of ethyl nitrite is moderated by ferrous sulphate. "

Those wishing to explore this route should be well acquainted with Ethyl nitrite's MSDS (a link provided) given its poisonous and explosive properties, and handle only in alcoholic solutions.
http://www.msdshazcom.com/REFERENCE/HAZARDS/ETHYL%20NITRITE....

[Edited on 18-4-2012 by AJKOER]

franklyn - 14-10-2012 at 23:55

A lttle paper from 1905 on inorganic nitrites and their decomposition by heat.
How to derive by ion exchange in solution with silver nitrite.

Attachment: Nitrites & decomposition by Heat.pdf (738kB)
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Zinc Nitrite , Zn(NO2)2
CAS 10102-02-0
Literature References: Prepared by treating Sodium Nitrite with Zinc Sulfate in
alcohol: Ephraim, Bolle, Ber. 48, 643 (1915).
Properties: Hydrolyzes so quickly that it cannnot be prepd from water, from
which only basic salts will separate.

Magnesium Nitrite , Mg(NO2)2
is prepared by a solution of Barium Nitrite and Magnesium Sulphate, and has also
been obtained by treating Silver Nitrite with Magnesia or with Magnesium Chloride.
The trihydrate, Mg(NO2)2.3H2O, crystallises in deliquescent leaflets which are
yellowish or snow white. It slowly decomposes in a stoppered bottle, and its
solution decomposes, evolving nitric oxide when evaporated on the water bath.
Careful concentration under diminished pressure , or under diminished pressure
over sulphuric acid , results finally in the dihydrate. This has also been obtained
by digesting magnesium sulphate and sodium nitrite with 94 per cent , alcohol
and evaporating under reduced pressure. It occurs as a hard , efflorescent ,
white mass or as clear crystals , and is much more stable than the trihydrate.
It does not usually dissolve to a clear solution , and may partially decompose
on continued dehydration.

.

[Edited on 15-10-2012 by franklyn]

Natures Natrium - 18-2-2013 at 20:40

Just wanted to jot down a few quick observations I have made.

10g of KNO3 (directly from stump remover bottle) was held over a propane fueled bunsen burner flame in a 50mL procelain crucible for about 2 hours, during which time the steel tongs which held the crucible were glowing bright red but the crucible did not appear to be. A test with sulfuric acid did not indicate the presence of any substantial quantity of nitrite. A comparison with a pre-burn weight showed no detectible (> .1g) weight loss.

A control test tube with KNO3 had exactly the same non-reaction to sulfuric acid.

Next, 1g of KNO3 in a test tube was slowly warmed at first, and then strongly heated in the flame of a propane blow torch. The salt foamed a bit on melting, then became water clear. On further heating, when the glass and molted salt were both cherry red, the salt began to evolve gas. This was not timed, but continued for roughly about 5 to 10 minutes. At this point the glass had sagged slightly due to gravity, and was allowed to cool slowly by "flame polishing" until all visible glow had ceased. It was then allowed to cool, and after the salt solidified it held the glass together as the glass in the heating zone cracked into dozens of tiny pieces. A sulfuric acid test was largely positive for nitrites, emitting a good deal of brown oxides of nitrogen, and most of the salt dissolving.

Interestingly enough, the test tube also tested negative for weight loss, but this may be due to the inaccuracy of the old triple beam to measure weight differences in the 0.01g to 0.1g range.

Just thought I would share that the thermal decomposition works at least qualitatively, but takes very high temperatures; thus standard borosilicate is unsuitable for the task.

Edit:Second Attempt:

The same 10g of KNO3 in the same crucible was subjected to a propane bunsen burner flame. This time, the tongs held onto the crucible much further up the body of the curcible, and the crucible was better situated in the hottest part of the flame (took a bit to relearn my bunsen burner skills). Also a sheet of aluminum foil in a cone shape was used to help reflect heat inward to the crucible.

After a bit the crucible was glowing orange-red in the dark, and small bubbles could be seen making their way to the surface. It was left in this state for roughly half an hour, at which point the supposed oxygen evolution had slowed considerably.

At this point the molten salt was poured directly into a clean but old and cracked beaker. After cooling some salt was scrapped up and tested for nitrite with sulfuric, and the salt tested negative. The beaker (which cracked plenty more) was broken and the salt harvested, several pieces of which tested negative for nitrite.

Well, I am at a loss. Thermal decomposition of the potassium salt of nitrate clearly works, but also doesn't seem to be very reliable. A reducing agent seems like it would make things much more dependable. I have some lead shot around here that I may give a try with at some point.

[Edited on 20-2-2013 by Natures Natrium]

AJKOER - 19-2-2013 at 19:42

Quote: Originally posted by JohnWW

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.
.......
John W.

Yes, there is some truth in this statement. Per recent research, apparently in an aqueous environment, the products are either nitrite and nitrate ions or just nitrite ions depending on pH at room temperature. To quote (source: "Mechanism of the NO2 conversion to NO2- in an alkaline solution" by Chen X, Okitsu K, Takenaka N, Bandow H. at the Department of Applied Materials Science, Graduate School of Engineering, Osaka Prefecture University):

"The reaction of NO2 and NaOH aqueous solution at room temperature was studied for elucidating the behavior of gaseous NO2 in an alkaline solution. Experimental runs related to NO2 absorption have been carried out in various pH solutions. The nitrite and nitrate ions formed in these absorption solutions were quantitatively analyzed. In the case of pH 5-12, both of the nitrite and nitrate ions were formed simultaneously. On the other hand, only the nitrite ion was formed when the pH of the absorption solution was higher than 13. In this paper, a new reaction mechanism was proposed to explain the selective formation of nitrite ion in the 10 M alkaline solution. In order to confirm the new reaction mechanism, H2(18)O was used as part of the absorption solution for detecting oxygen gas production. The amounts of reaction products: (18)O(18)O, (18)O(16)O and (16)O(16)O, were quantitatively determined. It was confirmed that the new reaction proceeds mainly in the 10 M alkaline solution."

Link: http://www.ncbi.nlm.nih.gov/pubmed/15636532

Now, if solid moist NaOH is treated with NO2, in this high pH condition, I suspect that primarily NaNO2 may form. However, this study using Soda Lime ("Elimination of nitrogen dioxide and nitric oxide by soda lime and its components", original in Chinese by Zhang D, Hu X, Liu J at the Department of Anesthesiology, Fu Wai Hospital, Chinese Academy of Medical Sciences, Peking Union Medical College) for scrubbing gases notes that NaOH or KOH by themselves are not effective in removing NO and NO2 unless Ca(OH)2 is present. Link: http://www.ncbi.nlm.nih.gov/pubmed/9772493

As such, I would also consider adding Ca(OH)2 to increase reactivity between the moist NaOH and gaseous NO2 at a high pH. I would expect the primary product to be NaNO2 (and some Ca(NO2)2). My take on the reaction:

4 NaOH + 4 NO2 --High pH--> 4 NaNO2 + 2 H2O + O2

jock88 - 25-2-2013 at 16:30

Some good info. in this thread about using Tin instead of lead
for making nitrites.

http://www.sciencemadness.org/talk/viewthread.php?tid=23485

I guess a mod should merge threads??

AJKOER - 6-4-2013 at 06:13

Here is a commercial based method, using Cadmium (and Cadmium treated with Copper) that is complexed either by NH4Cl or the Sodium salt of EDTA. In a neutral to basic solution the reaction is given by:

[NO3]- + H2O + 2 e- ---> [NO2]- + 2 [OH]-

and the oxidation of Cadmium by:

Cd + 1/2 O2 + H2O --> Cd(OH)2

or, if using EDTA:

[NO3]- + Cd + [EDTA]4- + H2O ---> [NO2]- + Cd[EDTA]2- + 2 [OH]-

The method is claimed to have near quantitative reduction of nitrate to nitrite.

Source: Journal of the Marine Biological Association of the United Kingdom / Volume 47 / Issue 01 / February 1967, pp 23-31. Link to full text: http://www.google.com/url?sa=t&rct=j&q=determination...

For the current application of processing large amounts of NaNO3, I would start by dissolving NaNO3 in hot water and add powdered Cd, and further heat. One could improve the reactivity of the Cadmium by employing the thermal decomposition product of Cd oxalate in nitrogen. To quote: "It is suggested that the decomposition of cadmium oxalate to cadmium and cadmium oxide is ... and stated that the formation of metal is the result of reduction of oxide". A reference link to the entire article (for a fee) can be found at http://www.sciencedirect.com/science/article/pii/00406031828... . I would prefer this route over complexing with NH4Cl as any excess Ammonium chloride could act on the newly created NaNO2 as follows reducing yield:

NH4Cl + NaNO2 --> NaCl + NH4NO2

as the Ammonium nitrite formed is unstable decomposing into water and N2 (and in concentrated/acidic solutions in an energetic manner).

[Edited on 6-4-2013 by AJKOER]

AJKOER - 8-4-2013 at 06:58

For those wishing to work with more friendly metals and readily available to turn nitrates into nitrites, consider Aluminum (Al foil no less). Per one source (see equation [1] under Section 6.4, link:
https://docs.google.com/viewer?a=v&q=cache:Jz8OCxPNXSoJ:...), to quote:

"1. 3NO3- + 2Al + 3H2O → 3NO-2 + 2Al(OH)3

2. NO-2 + 2Al + 3H2O → NH3 + 2Al(OH)3 + OH-

3. 2NO-2 + 2Al + 4H2O → N2 + 2Al(OH)3 +2OH-

Nitrate reduction was found to be pH dependant. At pH values less than eight
no nitrate reduction took place. Above pH 10.5 nitrate was reduced upon addition of
the aluminium powder. Aluminium powder has been suggested for the denitrification
of sodium-based nuclear wastes, employing the nitrate to ammonia and ceramic
(NAC) process (Mattus et al., 1993, 1994)."

Here are encouraging comments from another source (see http://en.wikibooks.org/wiki/Inorganic_Chemistry/Qualitative... ):

"The Nitrate ion can easily be reduced to ammonia with either Devardas Alloy or Aluminium Foil. The aluminium is a very powerful reducing agent, and this combined with heating causes the nitrate ions to form ammonia gas. This can be tested for by holding a piece of damp red litmus paper over the end of the test tube. The ammonia will form alkaline ammonium ions in the water and turn the paper blue.

4NO3-(aq) + 6H20(l) -> 4NH3(g) + 9O2(g)

Aluminium powder is not shown as it merely catalyses the reaction."

As a third, and some would claim as less credible authority, I present my analysis on the mechanism on exactly how Aluminum is involved in reaction [1]. First, as the Al only initiates the reaction in only more alkali conditions, I suspect the following occurring:

2 Al + 6 H2O --High pH Weakening Al2O3 Coating--> 2 Al(OH)3 + 3 H2 (g)

Or, more directly per the action of the strong alkali, for example;

2Al(s) + 2NaOH + 6H2O → 2Na[Al(OH)4] + 3H2(g)

That is, I would claim, it is nascent Hydrogen (label it H*) responsible and as some support, I do recall reading in an old text that freshly prepared Hydrogen can reduce nitrates (I look for the reference). So, a possible reaction, for example:

3 NaNO3 + 6 H* --> 3 NaNO2 + 3 H2O

So, in my opinion, it is the nascent H2, and not directly the Aluminum, that is the so called powerful reducing agent here. Following this point, nascent Hydrogen generation say from an external metal/acid reaction (like Zn/HCl) may also prove to be successful. Here is a source on nascent H2 (see http://en.wikipedia.org/wiki/Nascent_hydrogen ), to quote:

"According to one claim, nascent hydrogen is generated in situ usually by the reaction of zinc with an acid, aluminium (Devarda's alloy) with sodium hydroxide, or by electrolysis at the cathode.[citation needed] Being monoatomic, H atoms are much more reactive and thus a much more effective reducing agent than ordinary diatomic H2, but again the key question is whether H atoms exist in any chemically meaningful way under the conditions claimed. The concept is more popular in engineering and in older literature on catalysis."

A second, and very important point, is that per the three equations given above, and the evidence of ammonia formation per the 2nd author, one cannot employ an excess of Aluminium (else further reduction to NH3 and N2 could occur), and also, I would recommend stirring. The good news is that the reaction proceeds on gentle heating.

[EDIT] I found an old reference, "A manual of chemical analysis, qualitative and quantitative", by George Samuel Newth, page 146, link: http://books.google.com/books?pg=PA146&lpg=PA146&dq=... ) to quote:

"Reduction by Nascent Hydrogen. (1) With Formation of Nitrite.—When a nitrate in solution is exposed to the gentle action of nascent hydrogen—derived by the action of sodium amalgam, zinc amalgam, or copper-zinc couple—the nitrate is reduced to nitrite—

KN03 + H2 = H20 + KN02

This test is very delicate, and may be carried out as follows: A small piece of zinc foil (or granulated zinc) is placed in the solution of the nitrate in a test-tube, and one drop of copper sulphate added (this causes the deposition of a minute quantity of copper upon the zinc, thus creating the "copper-zinc couple"). The mixture is gently boiled for a minute or two. 0ne drop of the liquid (after cooling) is placed upon a piece of potassium-iodide-and-starch paper, and then touched with a glass rod moistened with dilute sulphuric acid. The paper will be instantly stained blue by the liberation of iodine and formation of iodide of starch (test for a nitrite).

[foot note]
* The mechanism of this reaction is sometimes explained by supposing that the nitric acid, set free from the nitrate, acts upon the copper according to the familiar equation 3Cu + 8HN03 = 3Cu(N03)2 + 4H20 + 2N0. But, as a matter of fact, copper sulphate, not nitrate, is found in solution, the whole of the nitrogen being converted into nitric oxide. Moreover, charcoal may be substituted for the copper. If dilute nitric acid be boiled with charcoal no brown fumes are formed, but on the addition of a little sulphuric acid they at once appear, owing to the action of the sulphur dioxide which is evolved from carbon and sulphuric acid."

My new take on the reputed creation of nascent H2 by the several mentioned paths is that it appears to me that there may be, in fact, an electrochemical connection (note, my quote source above on nascent Hydrogen citing one possible formation also "by electrolysis at the cathode"). Also, chemically pure Aluminum (or Zinc) may be inferior to say Al foil, an amalgam.

A note on Devarda's alloy (aluminium (44% – 46%), copper (49% – 51%) and zinc (4% – 6%)), to quote (see http://en.wikipedia.org/wiki/Devarda's_alloy ):

"Devarda's alloy is used as reducing agent in analytical chemistry for the determination of nitrates after their reduction to ammonia under alkaline conditions. It owes its name to the Italian chemist Arturo Devarda (1859–1944), who synthezised it at the end of the 19th century to develop a new method to analyze nitrate in Chile saltpeter.[2][3][4]"

Also, "The reduction of nitrate by the Devarda's alloy is given by the following equation:

3 NO3− + 8 Al + 5 OH− + 18 H2O → 3 NH3 + 8 [Al(OH)4]− "

This reaction with excess Aluminum is in complete agreement with a prior comment by Woelen with a qualification that I am suggesting by limiting the Al, per equation [1] above, and per an old qualitative test for nitrite, one may be able to successfully reduce nitrates to nitrite. To quote:

Quote: Originally posted by woelen

Trying to use iron in aqueous solution will not produce any nitrite. Nitrate ion is remarkably inert in aqueous solution, except in strongly concentrated acidic environments or at very high pH in combination with magnesium and aluminium. In the acidic environments it forms NOx with suitable reductors, in the strongly alkaline environment it forms ammonia.
.....

[Edited on 8-4-2013 by AJKOER]

Oscilllator - 4-8-2013 at 01:54

A couple of my own observations regarding potassium nitrate -> nitrite through thermal decomposition:
Over several batches, approx 250g KNO3 was boiled in a stainless steel saucepan for about 15 minutes. The source of the heat was a coke-fired forge that had previously proved itself capable of melting steel During the boiling, copious quantities of smoke was evolved.

When the molten salt was cooled down, it formed a distinctly green solid (greener than it looks in the photo). perhaps this is some kind of iron compound from the stainless steel?

This was then dissolved in approx 500ml of water, heated by placing it on the forge. This brought the 500ml of water to the boil in about 10-15 seconds

. The resultant brown muddy liquid was filtered to obtain a clear yellow solution that formed brown fumes upon addition of HCl. When the liquid was cooled down, large quantities of needle-shaped crystals formed that proved to be potassium nitrate :mad:

This solution was then boiled down and a second crop of crystals were formed. They had a slightly different shape, and I am not certain that they are KNO3 or KNO2, but I'm afraid the odds favour the KNO3. Irrespective of that, large quantities of KNO2 were formed in this process, as least enough that the evolution of NO2 was so great that the test tube almost bubbled over.
It is interesting to note that wikipedia and a number of other sources say that KNO2 explodes when heated above around 700 degrees. I am 90% sure this is not the case based on the fact that a rounded tablespoon of KNO3 was fully decomposed all the way to some black gunk that bubbled when water was dropped on it. Presumably this black substance is K2O. At any rate, it was sitting on top of a red-hot piece of steel and didn't do anything that could be described as an explosion

.

I do have a hypothesis regarding the entire method of producing nitrites by thermal decomposition: Is it possible that as the nitrate decomposes the formed nitrite also decomposes, so at no point in time is there a solution (if thats the right word) of 100% nitrate?

plante1999 - 4-8-2013 at 02:07

Nice try, but it was known thermal process works so so and need multiple fractional crystallization. Try to use lead or copper powder next time.

S.C. Wack - 4-8-2013 at 05:29

Quote: Originally posted by Oscilllator

It is interesting to note that wikipedia and a number of other sources say that KNO2 explodes when heated above around 700 degrees.

Boiling point 537 °C (explodes) may be total bullshit. Perhaps 700C is what it takes for this to "work" (15 minutes is probably still not long enough), as has been noted in this thread.
http://www.sciencemadness.org/talk/viewthread.php?tid=52&...

PS My summary of all this is: probably best done with some form of protection from O at a temperature not much below where N in some form is lost. No explosions IME.

[Edited on 4-8-2013 by S.C. Wack]

watson.fawkes - 4-8-2013 at 06:24

Quote: Originally posted by Oscilllator

When the molten salt was cooled down, it formed a distinctly green solid (greener than it looks in the photo). perhaps this is some kind of iron compound from the stainless steel?

Chromium oxide, I'd guess. Stainless steel is stainless because it forms passivating surface coatings of chromium oxides. I'd guess you've liberated some of the Cr from the outer layer of the pot.

blogfast25 - 4-8-2013 at 07:20

Quote: Originally posted by watson.fawkes

Chromium oxide, I'd guess. Stainless steel is stainless because it forms passivating surface coatings of chromium oxides. I'd guess you've liberated some of the Cr from the outer layer of the pot.

Yes, that happens each time I use a SS 'crucible' with some alkaline fusion. The iron itself doesn't seem to be affected much but Cr is amphoteric.

AJKOER - 25-12-2013 at 13:59

Quote: Originally posted by AJKOER

For those wishing to work with more friendly metals and readily available to turn nitrates into nitrites, consider Aluminum (Al foil no less). Per one source (see equation [1] under Section 6.4, link:
https://docs.google.com/viewer?a=v&q=cache:Jz8OCxPNXSoJ:...), to quote:

"1. 3NO3- + 2Al + 3H2O → 3NO-2 + 2Al(OH)3

2. NO-2 + 2Al + 3H2O → NH3 + 2Al(OH)3 + OH-

3. 2NO-2 + 2Al + 4H2O → N2 + 2Al(OH)3 +2OH-

Nitrate reduction was found to be pH dependant. At pH values less than eight
no nitrate reduction took place. Above pH 10.5 nitrate was reduced upon addition of
the aluminium powder. Aluminium powder has been suggested for the denitrification
of sodium-based nuclear wastes, employing the nitrate to ammonia and ceramic
(NAC) process (Mattus et al., 1993, 1994)."
.........

A week ago, I tried to convert some of my KNO3 to KNO2 using Al foil per Equation [1]. I am now reporting some of the results, for those interested, to the best of my recollection.

I restricted the aluminum to the amount required for just Equation [1] per above:

3NO3- + 2Al + 3H2O → 3NO-2 + 2Al(OH)3

and also added Na2CO3 to make the solution alkaline as was required per the author's discussion. After the 1st day, visible signs of Al(OH)3. I extract some of the solution to experiment further with the nitrite (another story).

What is interesting is what happened in the remaining solution in the following day. A significant gas formation became evident. When I opened the vessel, I was greeted by a rush of ammonia fumes. This was in accord with Equation [2] above, namely:

NO-2 + 2Al + 3H2O → NH3 + 2Al(OH)3 + OH-

My conclusion, was that removing of the some of aqueous KNO3/KNO2 may have created an excess of Aluminum which relatively rapidly (as compared to Eq [1]) formed the NH3 per above.

If I repeat the experiment again, I will certainly use an excess of KNO3. Clearly, the significant formation of ammonia implies that the nitrate can be converted to nitrite and then ammonia by this method. I am not certain, however, as to precise yield measures on the KNO2.

The purpose of my experiment was to test the possibility of nitrite only formation using Aluminum foil. I am more confident now that it may be possible.

[Edited on 25-12-2013 by AJKOER]

thebean - 4-1-2014 at 13:08

I had a semi successful non stoichiometric micro scale reaction. I put some potassium nitrate in a test tube with iron wool and heated until bubbles began forming. Small isolated incidents of ignition occurred in the iron but I would remove heating to stop the combustion. After the reaction stopped I was left with yellowish crystals which I suspect were nitrite. To test for presence of nitrite, I added some hydrochloric acid which produced brown fumes of what can only be NOx. I wafted a little and it had the classic NOx scent with a hint of chlorine. If one had proper amounts of nitrate and iron (preferably powder) and placed it in a small FBF with gentle heating up to the melting point of the nitrate and maintain that temperature until all of the iron appeared to have reacted you could produce decent amounts of the corresponding alkali nitrite, of course I would guess that an order of nitrite would be cheaper. It appears that you can replace the lead in the typical reaction with a variety of metals, the easiest to acquire being iron but lead having the best results.

AJKOER - 4-1-2014 at 17:16

Here is a source confirming the use of a variety of metals. To quote (see https://www.google.com/url?sa=t&rct=j&q=&esrc=s&... ):

"Potassium nitrate decomposes to the nitrite when heated to its melting point. Reducing
agents can also be added to facilitate the reaction to completion these are Pb, Fe2S, Fe2O, Cu, MnO2,
Fe, chromic, and sulfur. Best results were from MnO2 at 650 ° C.

The nitrite can also be formed by reduction of a nitrate with a metal in their aqueous
solutions. Examples of the reducing agents are Zn, K, Na, Pb, or Zn-Amalgam. The problem with
these reactions is the nitrite when heated in the aqueous solution will absorb oxygen from the air
and water, which reforms the nitrate. The electrolysis of sodium nitrate with amalgated copper
cathode. Silver nitrate solution reacted with silver and heated yields silver nitrite.

Ammonia can be oxidized to ammonium nitrite by the addition of hydrogen peroxide or
potassium permanganate.
In the reaction of a nitrate with granulated lead, the reducer should be in 15% excess. The
reaction commences at the boiling point and after 2 hours at 420 °C, complete reduction is obtained
(example with sodium nitrate).

KNO3 + H3PO3 → KNO2 + H3PO4
NaNO3 + CaO + SO2 → CaSO4 + NaNO3
2KNO3 → 2KNO2 + O2 with heat
NaNO3 + R → NaNO2 + RO where R is the metal reducing agent mentioned above

Pb + NaNO3 → PbO + NaNO2 "
--------------------------------------

Here is another old source, "Journal of the Society of Chemical Industry", Volume 27, based on commercial processes with some very insightful commentary on processes (see pages 484 to 485, link: http://books.google.com/books?pg=PA484&lpg=PA484&dq=... ). For example:

"Thermal decomposition.—The alkali nitrates when heated above their fusion point evolve oxygen and furnish nitrites, but this reaction is, of no practical importance owing to the simultaneous occurrence of a further decomposition to oxide. "

There is also an interesting aqueous method mentioned. To quote:

"The heating of a concentrated solution of lead nitrate with finely divided lead leads to the formation of insoluble basic lead nitrite, which furnishes sodium nitrite by double decomposition with sodium carbonate. This reaction is, however, only of theoretical interest. "

Per the prior source, one can employ Zinc in place of Pb in this aqueous reaction. Apparently, Zn dust is effective in reducing nitrate to nitrite (see, for example, http://www.microbelibrary.org/library/laboratory-test/3660-n... ). To quote:

"FIG. 5. Zinc dust will reduce nitrate to nitrite, but will not further reduce nitrite to nitrogen gas or other nitrogenous by-products when used sparingly."

The author also notes, to quote:

"3. Some authors recommend adding zinc to colorless NO2- reactions that do not contain gas to make sure that the NO2 has not been oxidized to NO3 rather than having been reduced to a nitrogen product other than N2 gas (21), but that reaction is rare."

which suggests to me the need to remove O2 from the concentrated aqueous nitrate solution by pre-boiling and avoid further exposure to air (perhaps by adding CO2,..). In place of Zn dust, I might use (when I perform this reaction) a Zn colloidal suspension in an O2 free solution (say from the action of Aluminum foil on an aqueous Zinc ammonium salt).

[Edited on 5-1-2014 by AJKOER]

Alyosha Karamazov - 13-1-2014 at 15:56

The J. Soc. Chem. Ind. (1908) review on nitrite production mentioned by AJKOER is a very interesting read. Unfortunately I can't view it using google books in my country, but archive.org is more lenient when it comes to copyright restrictions, I downloaded the journal issue from there and snipped the article so I could provide it here.

Also included is the Compt. Rend. article from 1889 (and not 1900 stated in the review) describing briefly the reduction of sodium nitrate by heating with barium sulfide, the mixture diluted with barium sulfate to make it less violent and more controllable. Separation of the nitrite is easy and the formed BaSO₄ could be used to produce the sulfide again, as suggested by the author.

Also provided here is a BASF patent (FR363643) describing the oxidation of nitrogen in air by electrical discharge. Interesting is that the author notes that while the nitric oxide produced readily oxidises in air to nitrous anhydride (N₂O₃), the latter oxidises much less readily to nitrogen dioxide at elevated temperatures. The rate of oxidation of N₂O₃ to NO₂ is 20x higher at 0°C compared to 100°C, and at 300°C it is virtually non-existant. Hence the author suggests working at that temperature to transfer and absorb the nitrous anhydride in basic solutions.

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S.C. Wack - 4-7-2015 at 07:46

Quote: Originally posted by komodo13

J. Soc. Chem Ind., 27, 483-5 (may30)-The author reviews various methods for the production of nitrite from nitrate by reduction with metals

Yesterday I also happened to independently find this article while searching for something totally unrelated. Searching the author's name just now shows that just a couple days before that, a different article by the same author was requested in the references thread. And, this person here who signed up a month after me had come by to drop the name, made one more post, then disappeared. I'm not sure if posting the full text, and the author's name for the third time onsite, is the last thing necessary to start the Apocalypse or what, but now I'm afraid not to. I better include the abstract of his obit in a 1940 issue of Nature: SIR GILBERT MORGAN died after a very short illness on February 2 in his seventieth year. For fifty years he had been engaged in chemical research and probably no other chemist had such wide knowledge of scientific and applied chemistry. A fitting mark of the Jubilee was the presentation to him in July last of the medal of the Society of Chemical Industry, its highest honour; on this occasion he gave an account of his career and an outline of his researches.

THE MANUFACTURE OF SODIUM NITRITE.
GILBERT T. MORGAN
J. Soc. Chem. Ind. 27, 483 (1908)

Sodium nitrite is practically the only salt of nitrous acid which is prepared on a manufacturing scale and it finds extensive use in the production of several classes of artificial colouring matters and also in the preparation of various pharmaceutical products and other fine chemicals.

I. Production from sodium nitrate.

Hitherto nitrite has generally been prepared from sodium nitrate, and the following are some of the principal methods by which this chemical change can be effected.

1. Thermal decomposition.—The alkali nitrates when heated above their fusion point evolve oxygen and furnish nitrites, but this reaction is of no practical importance owing to the simultaneous occurrence of a further decomposition to oxide. The difference in the behaviour of the two alkali nitrates is of some theoretical interest, the potassium salt giving a final residue of the dioxide, whilst the sodium salt yields the monoxide.

2. Reduction by metals.—Finely divided copper has been suggested by Persoz and by Muller and Pauly, but owing to the infusibility of this metal at the temperature of molten nitre it is difficult to ensure a uniform reduction throughout the heated materials, and, moreover, this process involves the regeneration of the comparatively expensive copper from the resulting copric oxide. Zinc dust has been tried but I am not aware that a successful method has been based on the use of this metal. Sturm (Fr. Pat. 321,498 of 1902; this J., 1903, 212) obtains nitrite by heating sodium nitrate with finely divided metals in a muffle furnace. More recently iron has been recommended as a reducing agent for sodium nitrate.

The heating of a concentrated solution of lead nitrate with finely divided lead leads to the formation of insoluble basic lead nitrite, which furnishes sodium nitrite by double decomposition with sodium carbonate. This reaction is, however, only of theoretical interest.

Lead is undoubtedly the most convenient metal for the reduction of sodium nitrate; it is comparatively cheap, and its melting point (330—335° C.) lies so close to that of the nitrate (314° C.), that the two reagents can be brought together in the liquid condition at a temperature considerably lower than that at which the thermal decomposition of nitrite occurs. (Hampe, Annalen, 1863, 125, 336.) The reduction is effected on a manufacturing scale in shallow cast-iron pans, 4 ft. in diameter and about 18 inches deep, fitted with a stirring gear which agitates thoroughly the whole mass of molten material. The pans are supported on perforated firebrick arches, arranged so that the products of combustion of the coal fires circulate uniformly round the pan before passing to the fines. Each pan is charged with 200 lbs. of sodium nitrate and 3 cwts. of lead, which are heated and stirred until all the metal is oxidised. At this stage 50 lbs. of nitre are added and thoroughly stirred in until the mixture is of uniform consistence when 3 1/2 cwts. of lead are gradually added, with constant agitation, the mechanical stirring being supplemented by the use of long-handled rakes employed to “pull out" the mixture of metal and salt from the central and hottest part of the pan. The success of the reduction depends very largely on the skill and experience of the workman, who generally controls two pans, and whose duties are to regulate the firing of the pan and the rate of addition of the metal.

The reaction occurring in the nitrite pan is not so simple as that represented bv the equation, Pb + NaNO3 = PbO + NaNO2, for the higher oxides of lead are also produced both by aerial oxidation and by further reaction with the nitrate. In this connection it should be noted that litharge itself has been suggested as a reducing agent for nitrate, 3PbO + NaNO3 = NaNO2 + Pb3O4, although the large proportion of this oxide required would militate against its adoption.

During the reaction a portion of the lead becomes converted into a singularly inert substance of high specific gravity; this product is regarded as a sub-oxide, but may be a passive form of the metal contaminated with higher oxides and sodium plumbite.

The fusions are allowed to run for one-half to three quarters of an hour after all the lead has been added, when the product, which now has a yellowish brown tint, is tested for nitrite. With careful working the soluble constituent of the melt should contain 90 per cent. or even more of sodium nitrite. When rich in nitrite the cooled melt has on its surface a characteristic crystalline incrustation which is never noticed on specimens containing a relatively small percentage of the required salt.

Satisfactory results are, however, only obtained when both the lead and nitrate are of good quality. The former should be good commercial lead which has been remelted, skimmed and cast into small bars. The latter should be crystallised Chili saltpetre of the best quality. The presence of sodium iodate is especially harmful, as this salt appears to act catalytically in promoting the destruction of the molten nitrate and nitrite. The appearance of the characteristic violet vapour of iodine arising from the melt shows that inferior nitrate is being employed and in these circumstances the salt should be recrystallised, when the harmful impurities are eliminated.

The reduction being complete, the molten contents of two pans are ladled into 120 gallons of warm water contained in a covered cylindrical washing box fitted with powerful stirrers rotating on a horizontal axis. The curved lid of this box contains two small circular apertures for the introduction of the melt, and within, the soluble nitrite is separated from the litharge by agitating the mixture for one hour and then allowing the precipitate to subside. The clear liquor is then run off into a neutralising tank and here the solution, which is distinctly alkaline, is neutralised with dilute sulphuric acid or, if possible, with the solution obtained by absorbing nitrous fumes in water. These fumes are obtained in such nitric acid oxidation processes as the manufacture of arsenic acid.

This neutralisation causes the decomposition of sodium plumbite and the precipitation of a "small amount" of lead hydroxide. The neutralised liquors are now evaporated in wrought iron pans heated either directly or with internal steam pipes. At 45° Bé. the concentrated solution shows a thin film of nitrite on its surface and is then run off into rectangular cast iron crystallisers and left for at least 12 hours, when the first crop of sodium nitrite crystals is collected, the mother liquors being again concentrated and allowed to crystallise. These crystals, when dried in a centrifugal hydro-extractor and then in an air-oven, should contain 96 per cent. of NaNO2.

The washing of the oxides of lead is repeated, the more dilute washing liquors being flushed off together with the litharge into settling tanks. These liquors are used to lixiviate subsequent melts, whilst the litharge is either dried for sale or mixed with the “sub-oxide” and smelted to lead in a small blast furnace.

As received from the settling tanks the pasty litharge contains varying proportions of the higher oxides of lead. It may, however, be rendered more uniform by conversion into flake litharge or red lead in suitable reverberatory furnaces.

The working up of these large quantities of lead compounds is one of the chief disadvantages of this process, another is the baleful effect of the lead on the workmen, a certain incidence of plumbism being almost unavoidable.

3. Reduction by non-metals.—The well-known detonation of charcoal and nitre leads to the formation of a carbonate and only small quantities of nitrite. When brought under control by the addition of caustic soda and lime the reaction between sodium nitrate and graphite has been patented as a process for preparing sodium nitrite (Grossmann, Eng. Pat. 1452 of 1904; and also Knop, Eng. Pat. 4747 of 1897).

A similar process involving the use of sulphur, which has been successfully worked out on a manufacturing scale by Messrs. Read Holliday and Sons, of Huddersfield, is based on the following reaction:—
3NaNO3 + S + 2NaOH = Na2SO4 + 3NaNO2 + H2O.
The fusion is carried out in open pans fitted with stirring gear, but of larger capacity than those employed in the lead process. The nitrate containing a portion of the caustic soda (about 1/10th) is melted and treated alternately with sulphur and more molten caustic soda until the nitrate is practically all reduced. The fused product is added while still hot to sufficient warm water to dissolve the whole of the nitrite and only a portion of the sodium sulphate, a large proportion of which is left behind in a granular condition. The liquor is drained through a vacuum filter and evaporated to a smaller bulk when a further portion of sulphate separates and the solution on cooling deposits sodium nitrite while the final mother liquors furnish more Glauber’s salt.

This process yields without troublesome by-products a nitrite of good quality, which is obviously quite free from lead, and apart from the separation by fractional crystallisation of the nitrite from the dissolved sulphate, the operations involved present no serious practical difficulties.

Carbon monoxide either pure or in the form of producer gas has no action on molten sodium nitrate but in the presence of fused caustic soda there in an intermediate formation of formate which then reduces the nitrate in the following manner:—
NaNO3 +HCO2Na + NaOH = NaNO2 + Na2CO3 + H2O.
(Goldschmidt, Eng. Pat. 17066 of 1895.)

4. Reduction by metallic sulphides and sulphites.—Etard formerly recommended sodium sulphite as a reducing agent for sodium nitrate (Bull. Soc. Chim., 1877, 27, 434), and a modification of his process has recently been patented (D.R.P. 138,029). The percentage yield of nitrite is excellent but the high proportion of sodium sulphate in the melt - about two-thirds of the total — renders the separation of sodium nitrite somewhat troublesome. Closely allied to this method is the process devised by the firm of Gebruder Flick, which consists in passing sulphur dioxide over sodium nitrate and calcium hydroxide heated in retorts. The nitrite is then readily separated from the sparingly soluble calcium sulphate.

The commonly occurring sulphides react with fused sodium nitrate, furnishing nitrite. Messrs. McGougan have patented the use of galena which gives a melt containing litharge, sodium nitrite and sulphate with a small amount of sodium plumbite (Eng. Pat. 7715 of 1897). I have noticed that stibnite, the fusible sulphide of antimony, very readily reduces the nitrate giving a high percentage of nitrite but the cost of this reducing reagent is prohibitive.

Sodium sulphide and nitrate interact energetically forming a nitrite melt which contains only a relatively small proportion of sodium sulphate. Le Roy has advocated the use of barium sulphide, a mixture of this substance and sodium nitrate being heated in an iron dish when a vigorous reaction sets in and sodium nitrite and barium sulphate result. The intensity of this reduction is moderated by the admixture of barium sulphate (Compt. rend., 1900, 108, 1251).

II. Production of nitrite from nitrous fumes.

It was shown conclusively by Divers (Trans., 1899, 75, 85) that pure sodium nitrite could be readily prepared by absorbing nitrous fumes in aqueous sodium carbonate or hydroxide, provided that these gases contain a slight excess of nitric oxide. Excess of nitrogen peroxide would result in the formation of nitrate. Nitric oxide itself was shown by Debray to unite with barium peroxide forming barium nitrite, a similar reaction with the alkali peroxides would lead to sodium and potassium nitrites.

Raschig‘s observation that nitric oxide combines very rapidly with oxygen to form nitrous anhydride whilst the further change of the latter oxide to nitrogen peroxide occurs comparatively slowly suggests a method of utilising atmospheric nitrogen in the production of nitrite. The absorption of the nitrous fumes within a few seconds of their formation in the electric arc is an operation involving considerable practical difficulty, which however has to some extent been overcome by the method recently patented by the Badische Anilin- und Soda-fabrik (Fr. Pat. 363,643 Of 1906). According to this patent the nitrous fumes are maintained at a temperature of 300° C. until they are absorbed by an alkaline solution of sodium nitrite from a former operation. A strong solution of this salt is employed in order to reduce as far as possible the vapour pressure of the liquid, and thus minimise the dilution of the hot reacting gases with steam.

According to Eyde (Eng. Pat. 28,613 of 1904) the gases of the electric furnace containing much air when quickly brought into contact with the hydroxides of the alkalis or alkaline earths, yield nitrites
2NO + NaOH + O = 2NaNO2 + H2O.

Electrolytic reduction of nitrates.-Various attempts have been made to utilise the electric current in the reduction of nitrates to nitrites. Among the most recent are the experiments made by E. Muller and F. Spitzer (Zeit. Elektrochem., 1905, 11, 509) with cathodes of different metals, the most favourable results being obtained with spongy silver.

Miscellaneous agents.-The interaction between barium hydroxide, manganese dioxide and sodium nitrate has been patented by Huggenberg. Zinc and ammonia have been employed by Stahlschmidt. The oxidation of ammonia in the presence of metallic oxides at 650-750° C. leads to the production of “nitrous anhydride” which is absorbed by alkalis (U.S. Pat. 763,491 of 1904). This oxidation of ammonia to nitrite has also been effected electrolytically in aqueous solution in the presence of sodium and cupric hydroxides (Traube and Biltz, Ber., 1904, 37, 3120).

III. Production of nitrite from calcium nitrate.

As it has been predicted that calcium nitrate will gradually displace the sodium salt as the commercial source of nitre, I have made some experiments on the production of sodium nitrite from calcium nitrate, or from mixtures of this salt with calcium nitrite.

Calcium nitrate melts in its water of crystallisation becomes solid again at higher temperatures and finally fuses. When maintained in a pasty state for some time the anhydrous salt loses oxygen and oxides of nitrogen; some nitrite is produced but the yield is very small. Reducing agents increase the production of nitrite very considerably and when mixed with sodium sulphite and sulphide the calcium nitrate on heating furnishes a yield of more than 60 per cent. of the calculated amount of sodium nitrite. The object of taking the two reducing agents in these proportions is to ensure the conversion of both sulphur compounds into sparingly soluble calcium sulphate.
i. 2Ca(NO3)2 + Na2S = CaSO4 + Ca(NO2)2 + 2NaNO2
ii. Ca(NO3)2 + 2Na2SO3 = CaSO4 + Na2SO4 + 2NaNO2
The combined changes may be represented as follows:-
3Ca(NO3)2 + Na2S + 2Na2SO3 = 6NaNO2 = 3CaSO4

By taking the sulphide and sulphite in these proportions the product after lixiviation consists chiefly of very soluble nitrite and sparingly soluble gypsum which are readily separated. The sulphide and sulphite are melted together until their water of crystallisation is driven off and the residue intimately mixed with the calcium nitrate. This mixture is heated until the water contained in the last salt is eliminated. A portion of the dried mixture is then heated strongly until a reaction sets in accompanied by incandescence and the remainder is added sufficiently rapidly to ensure the continuance of this interaction. The greyish white product is lixiviated with warm water, the sulphate removed and the nitrite obtained from the solution.

With a mixture of nitrate and nitrite the sulphide may be omitted as in this case sulphite alone suffices to convert all the calcium into sulphate.
Ca(NO2)2 + Ca(NO3)2 + 2Na2SO3 = 4NaNO2 + 2CaSO4.

Instead of the sulphide and sulphite, a mixture of sulphur, caustic soda and calcium nitrate may be employed.
3Ca(NO3)2 + 2S + 6NaOH = 2CaSO4 + Ca(OH)2 + 6NaNO2.

The introduction of carbon dioxide or dilute sulphuric acid into the aqueous solution of the melt ensures the precipitation of the calcium hydroxide in the form of calcium carbonate or sulphate.

[Edited on 4-7-2015 by S.C. Wack]

glassplass - 28-4-2016 at 05:51

Quote: Originally posted by madscientist

.....The first method was heating calcium sulfite and sodium nitrate together. This had seemingly good yields of sodium nitrite..... I prepared the CaSO3 from NaHSO3 and CaCl2.....

Can you use potassium metabisulfite or potassium sulfite instead of calcium sulfite?Probably you can but both nitrite and sulfite ar soluble in water,and its dificult to separate them...
I dont have any sodium sulfites so i look for an alternative.
Can you directly buble SO2 gas into the solution of Ca(OH)2 or CaCO3,but aggain solubility in water is a problem so it will be ineffective,or?
If you cant do anything,than SO2 needs to be bubled into solution of NaOH to get mixtures of sodium hydrogen sulfite and sodium sulfite,but i dont know when is enough,it will be contaminated with hydroxide.
Sorry for confusion,help!!

AJKOER - 4-6-2016 at 15:02

Quote: Originally posted by AJKOER

.........

Here is another old source, "Journal of the Society of Chemical Industry", Volume 27, based on commercial processes with some very insightful commentary on processes (see pages 484 to 485, link: http://books.google.com/books?pg=PA484&lpg=PA484&dq=... ). For example:

"Thermal decomposition.—The alkali nitrates when heated above their fusion point evolve oxygen and furnish nitrites, but this reaction is, of no practical importance owing to the simultaneous occurrence of a further decomposition to oxide. "

There is also an interesting aqueous method mentioned. To quote:

"The heating of a concentrated solution of lead nitrate with finely divided lead leads to the formation of insoluble basic lead nitrite, which furnishes sodium nitrite by double decomposition with sodium carbonate. This reaction is, however, only of theoretical interest. "
......

Apparently, aqueous nitrate can be reduced by Pb to nitrite:

Pb + NO3- → PbO + NO2- (see reference [22] below)

Also, nitrite then reduced further by Pb to N2 or NH3 (depending on pH).

2 NO2- + 3 Pb + H2O → N2 + 3 PbO + 2 OH-

NO2- + 3 Pb + 2 H2O → NH3 + 3 PbO + OH-

Reactions reference:: https://www.google.com/url?sa=t&source=web&rct=j&...

Cited reference[ 22] is Uchida, Miho; Okuwaki, Akitsugu. (1998), "Decomposition of nitrate by in situ buff abrasion of lead plate", published in Hydrometallurgy, 49, 297-308. To quote the abstract:

"A new approach to the decomposition of nitrate ion using Pb metal has been developed. Removal of the oxide layer formed on the surface of Pb plate and the production of Pb powder have been achieved by in situ abrasion of Pb plate in NH4NO3 solution with an abrasive buff. NO3− was reduced to NO2− at initial concentrations of NH4NO3 in the range 0.01–0.2 M, rotational speed of buff, 100–800 rpm and temperature, 25–80°C. Complete reduction of NO3− in 0.05–0.2 M NH4NO3 solution was achieved at 80°C within 4 h. As the temperature increases, the reduction rate of NO3− to NO2− increases abruptly. The reduction rate increases gradually with rotational speed. The formation of NO2− is almost independent of the initial NO3− concentration. Reduction of NO3− to NO2− is related to corrosion of Pb in NH4NO3 solution."

Link: http://www.sciencedirect.com/science/article/pii/S0304386X98...

So apparently buffing the Pb electrode to remove PbO avoids the formation of the basic lead nitrite and leaves aqueous nitrite.

So no extreme temperatures or lead vapors are required, just combine a buffing hand tool with a galvanic corrosion reaction (which, I suspect would be accelerated by the addition of NaNO3 to the aqueous NH4NO3).

Preparation of sodium nitrite

aga - 28-8-2016 at 07:43

To attempt to follow the Congo Red sysnthesis, 0.5g of sodium nitrite is required, hence this preparation.

According to this procedure :
http://www.prepchem.com/synthesis-of-sodium-nitrite/

71g of sodium nitrate and 172g of elemental lead was added to an iron vessel, then heated strongly with a blowtorch.
(this represents an excess of lead).

The mixture was stirred with an iron bar throughout the process.

After 20 minutes of heating & stirring, a pink scum was evident on the surface of the liquid, presumably PbO.

After 30 minutes the mixture began to take on the appearance and texture of pink blancmange, with many bubbles in the structure: a foam.

At 40 minutes the centre of the mass began to glow red, so heating was discontinued, although manual stirring was maintained.

After 1 minute the mass began to solidify, so stirring was continued to try to achieve a small fragment size as per the reference.

The cooled mass appeared as a yellow/ochre 'rocks' with some pink showing, mostly in areas furthest away from the flame.

The product was extracted with a total of 500ml of hot water, which had to boiled down to about 50ml before it would crystallise out.

Commercial test strips were used to test for the presence of Nitrites

The plastic cup used for weighing the NaNO3 had 100ml of DW added, and showed positive for Nitrates, negative for Nitrites.

1.5ml of the supernantant liquid from the interim product was added to the same plastic cup, then the test was repeated with this mixture and a very strong Nitrite result was seen.

(The nitrite block is second from the left.)

The crude product was recrystallised and dried, giving a yield of 26.32g (44%)

Σldritch - 16-9-2016 at 06:02

Im trying too make nitrite by reduction with sulfur (2 NaNO3 + S + NaX --> 2 NaNO2 + NaSO4 + X2) and i have had some minor success with igniting what is essentially yellow powder with extra nitrate but i am trying to do this in solution now because of the explosion hazard.

When i tried melting sodium hydroxide and sulfur together i got red to orange goop and i poured it out on a rusty iron plate to cool to later add it to a solution of potassium nitrate. When i did i obtained a very dark green solution which mostly dispeared on boiling down the solution.

So what is the green compound? Iron complex with polysulfide or nitrite? Iron nitrosyl is brown and the ph is really high so i doubt that it would exist in the solution. I will post progress if i make any.

Sorry if this is supposed too go in beginnings not sure tough.

Melgar - 16-9-2016 at 21:43

Quote: Originally posted by Σldritch

Im trying too make nitrite by reduction with sulfur (2 NaNO3 + S + NaX --> 2 NaNO2 + NaSO4 + X2) and i have had some minor success with igniting what is essentially yellow powder with extra nitrate but i am trying to do this in solution now because of the explosion hazard.

When i tried melting sodium hydroxide and sulfur together i got red to orange goop and i poured it out on a rusty iron plate to cool to later add it to a solution of potassium nitrate. When i did i obtained a very dark green solution which mostly dispeared on boiling down the solution.

So what is the green compound? Iron complex with polysulfide or nitrite? Iron nitrosyl is brown and the ph is really high so i doubt that it would exist in the solution. I will post progress if i make any.

Sorry if this is supposed too go in beginnings not sure tough.

Blue Fe(II)SO4 mixed with yellow sulfur would give a green color. Eventually, the iron II sulfate would oxidize to iron III, which is more of a yellow/orange/brown color depending on concentration.

Σldritch - 17-9-2016 at 13:13

Quote: Originally posted by Melgar

Blue Fe(II)SO4 mixed with yellow sulfur would give a green color. Eventually, the iron II sulfate would oxidize to iron III, which is more of a yellow/orange/brown color depending on concentration.

I dubt that is what caused the color because it was very dark green to the point that after i washed out the beaker and neutralized some sodium bisulfate in it the color returned albeit not nearly as dark and it seemed to move down to the bottom of the solution.

Anyway the aqeous reduction with liver of sulfur was a failure as expected, worth a try atleast.

theAngryLittleBunny - 22-7-2017 at 14:49

I tried reducing KNO3 with tin today, and I think I should tell you what happend and why it isn't a good idea.

So I melted 62g of KNO3 in a metal can on a gas burner and put a 35g piece of tin into it. It obviously quickly melted and on stirring, within a few minutes a lot of gray SnO appeared and it seemed to be going well, the mixture became quite thick from the SnO. But after a substantial amount of SnO built up, the KNO3/KNO2 started reacting with it in a thermite like way, a ton of SnO2 smoke was released from the steel can, and as soon as I saw that the can started to melt, I knocked it from the gas burner and it cuntinued reacting violently for a like 30 seconds and melted completely through the container.

So just don't use tin for this, it's not a good idea.

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 ?

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).

[Edited on 5-10-2017 by AJKOER]

[Edited on 6-10-2017 by AJKOER]

AJKOER - 5-10-2017 at 14:11

Photolysis run:

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.

[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:

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

Fantasma4500 - 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.

Fantasma4500 - 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

Fyndium - 4-11-2020 at 13:04

When using KNO3 in place of NaNO3 in the OP reaction consisting of a nitrate, a base, and starch(in this case, potato or corn starch powder), is there a risk of forming cyanides?

I ask this because of this article in SM wiki:

http://www.sciencemadness.org/smwiki/index.php/Potassium_nit...

Quote:

A less known reaction is the synthesis of potassium cyanide, by reacting a mixture of potassium nitrate and charcoal in a cast iron bowl, in an inert atmosphere to prevent combustion or oxidation to potassium cyanate:

2 KNO3 + 7 C → KCN + KCNO + 5 CO[2]
If you attempt to try this reaction, AVOID ADDING ACID TO THE RESULTED SLAG AS IT WILL GIVE OFF HYDROGEN CYANIDE GAS WHICH CAN BE DEADLY (see the Sciencemadness thread below).

This is, because I don't have Na but K nitrate, and in theory, upon forming, potassium and sodium ions should precipitate sodium carbonate and potassium nitrate out easily when recrystallizing, and leave potassium nitrite in solution, because it has by far the highest solubility, hence the purification should be more convenient than with Na.

Fyndium - 5-11-2020 at 05:08

I performed a test with this. I ground up 100g of KNO3, 48g of NaOH and 20g of potato starch and placed the mix in steel pot and placed a steel plate as a lid. I heated it slowly with gas burner, and it started to buff up, smoke, and then a controlled reaction kicked in and it appeared to be burning for several seconds, generating white smoke, and then it stopped and ceased totally. I let it cool, and yellow rock hard mass was formed with slight greenish tint, not sure if it was from the red impurity of KNO3.

I dissolved it with 100mL of boiling water and added a little, but part was left undissolved. I filtered it clear, orange liquid, and cooled it. KNO3 and apparently what is Na2CO3 crystallized out. I concentrated the liquid by heating, and white mass what is likely Na2CO3 crashed out. I plan to process it further by cooling it down to 0 and see if it solidifies completely, and if, add just a bit to extract the KNO2, and eventually dry it in desiccator.

What I was afraid if this reaction could have generated any cyanides, according to my previous post. I haven't smelled any almonds, but if I try to sniff it too much, I begin to imagine that I notice a trace of it, but this happens every time. I suppose it should be detectable, unless there's that genetic thing? I smell normal and bitter almonds(BzH), though.

Σldritch - 5-11-2020 at 06:52

Nice to see someone actually giving this a try! Some this i learned while doing this when making nitrite for azide (which worked further proving this synthesis works):

1. The smell is from some volatile carbon compound presumably formed from the starch. You get it from pyrolysing starch too, which is probably the source. It can not be cyanide because we use oxidizer in a huge excess and knowing the smell of cyanide i can say this is not it though i can see the similarity. I like to clean up the filtrate with activated carbon, it gets rid of the smell and most of the color (it is slightly yellow as a solid).

2. I strongly recommend using Potassium Nitrate and Calcium Hydroxide or something equivalent such as Calcium Nitrate and Potassium Hydroxide to get good separation. While you might be able to get out Sodium Carbonate easily with just sodium you can not separate the Sodium Nitrate and Nitrite. There tends to be a lot of nitrate left over, which is worth saving. Mixing in potassium might improve it but at the cost of carbonate separation.

3. Crude yield tends to be 10% with mixing Calcium Nitrate and Potassium Nitrate lazily. Grinding does not seem to improve it much. Crude yield is 30% when dissolving the salts and then mixing and boiling it down until ignition. Takes more time but the better yield is absolutely worth it. Obviously use metal container for this.

4. When doing it as suggested in (3). Do not use starch, use sucrose or you will have a horrible, caustic, goopy mess to work with.

[Edited on 5-11-2020 by Σldritch]

Fyndium - 5-11-2020 at 08:37

Ah, thanks for the input. I thought you were long gone, as this is a few years old topic.

Is the ratio of starch/sucrose indeed 1mol per 12mol of nitrate and hydroxide? I used 100g of KNO3, 48 grams of NaOH and 20g of starch (16g would've been the original, but the amount was so insignificant I upped it just in case).

I was originally gonna use calcium hydroxide, but I thought that since you got it working, I should follow is as exactly as I can, because I urgently need the nitrite. Good thing is, I'm gonna react it with nitric acid, so any nitrate or carbonate residue does not matter, they will just be either consumed or be a spectator in the reaction.

If I need more, I'll use your instructions, use sucrose and calcium hydroxide and probably I'll blend them together with mini blender as I've done prior with different similar reactions.

I'm just currently cooling the solution to 0C. After filtration I'll put it in desiccator.

Σldritch - 5-11-2020 at 14:42

Im still here but as i no longer have year round access to my lab i can usually not add much to the discussion. To this one, however, i can.

The stochiometry is correct. Starch is basicly carbon hydrate and theoretically we want the carbon to end up as carbonate which means two nitrate molecules for every carbon atom. Without the base the acidic Carbon Dioxide probably displaces much of the nitrogen dioxide. But really you might have success using more carbon, i have never had a reason to try because the Potassium Hydroxide is the most valuble to me. Theoretically using more carbon means using more base to keep the acidity down but since so much nitrate always is left over this may not reflect reality. I guess a lot of the carbon is oxidized to carbon monoxide which is much less acidic leaving base and nitrate. From the crude product this seems to be the case, which is another reason to use Calcium, the hydroxide salts are easy to separate.

If you make the mix from Calcium Hydroxide and Potassium Nitrate you may want to add something to slow down the reaction. Normally water does it. I think the reaction running too hot is the cause of it going reddish. It forms a foam of Nitrogen Dioxide which can be smelled.

Another thing is, i never fractionally crystallize the crude nitrite but use it as is. With the amount of Alkali Hydroxides and the very soluble nitrite itself it is hard to do from water. You may have more luck with methanol or ethanol-water.

Fyndium - 5-11-2020 at 15:00

Could it be beneficial to actually use less nitrate compared to base and carbon, or does this kill the reaction kinetics?

I did not check if the stuff was red hot, I just left it cool when it stopped under a steel lid and came back when it was cold and rock hard. I ignited the reaction by simply heating the can, and definitely not too much to make it glow, more like how I heat my liquid reactions. It started very easy, foamed and bubbled a little, and then the big smoking started, lasted maybe 10 seconds or more, and ceased. The reaction seemed going extremely smooth, like it was buffered by something.

I first filtered the liquor after cooling a little when dissolving the rock, removing all solid impurities and crashed out stuff, cooled it to 0C, decanted off the formed crystals of likely KNO3, and concentrated it by boiling, and during boil, white mass crashed out.

When I cooled down this with the crashed white mass, a lot more of sharp crystals formed, and a liquid was left. I drained this liquid in another petri dish, and scraped the crystalline mass off from the clearly separate layer of white hard stuff, that is likely carbonate of something, and collected this crystal mass, which formed a watery slush in another dish, and placed them both in a desiccator with lots of CaCl2. I'll trickle some HNO3 on them samples when they are in a more solid form. My interest in optimizing this reaction due to other ventures is little, and I only need the nitrite in catalytic amounts, hence I'm well happy with tens of grams of it, and if this nitr-o-pot method works to produce at least that per run, it probably lasts more than my need for it. But I'll write on any notes I come up with.

Σldritch - 6-11-2020 at 11:27

Maybe it could be beneficial. But depends on what you are optimizing for, if it is yield from nitrate you can probably increase it as there tens to be a lot left over. There is risk of overreduction, and kinetics will probably be faster not slower. Faster kinetics means higher temperature, which may not be a problem. However what may be a problem with increased rate is that it should become more sensitive to particle size to some degree, this effect should get stronger with increased proportion of reductant, too.

I suspect the best way to increase yield is to increase the amount of base making sure to capture all the nitrogen oxides. Calcium Oxide/Hydroxide can be very cheap, problem is just getting it fine enough, which is why is why i recommend using Calcium Nitrate/Potassium Hydroxide. Another way is to replace the carbon which has the problem of forming carbon monoxide making it harder to get the amount of reductant correct because different amounts of carbon monoxide is formed depending on rate and physical conditions changing the stochiometry.

When i first did this i tried using sulfur. It seems to be able to reduce nitrate at a lower temperature with one caveat, it loves to explode - even with excess oxidizer and base present - think some solvent could make it work but nitrite is so soluble it is probably not worth it (the solvent would probably have to be 200 C + so no distilling away an organic solvent without fire/explosion hazard). I would not recommend it but if someone is mad enough this is probably the route to very high yields.

I should probably have documented this when i did it but i did not bother because interest was so low anyway. If not you maybe someone else will optimize it further, verification helps either way.

Fyndium - 6-11-2020 at 12:04

Calcium hydroxide is sold in hardware store 50c a kg, so the base is the least of the issues.

I'm not sure which method is best or better than this, but this is the first one that I found documented and working, because I already struggle with the endless scale of economies, where I have to make every reagent's pre- and even pre-precursors, and if I have to test and try all different methods, no time in the world is enough.

Calcium hydroxide probably allows for better separation, because what's insoluble is always easier to process than soluble, unless the sol curve is very steep. Potato starch residues should have very limited solubility in water in cold temps.

The stuff is half-dried by now, I'll keep it in desiccator for one or more days to get it completely dry, and then measure the collected weight. I presume this is far from pure, though.

Fyndium - 8-11-2020 at 05:42

I got it dried. Desiccator dry weight is 58g. I suppose the product is not very pure because the yield is so high. Maximum theoretical yield would be about 85g, hence the starting point is 68%.

Like I previously said, I first dissolved it in boiling hot water, filtered, cooled to 0C, filtered again from KNO3 and carbonates, and concentrated about half, cooled, decanted what was likely more Na2CO3 and KNO3 off, and then evaporated. If there is more than 50% of KNO2 I'm highly happy.

Gonna test it soon so I'll tell if it works. Can't determine yields though because I will be using just a catalytic amount.

Fantasma4500 - 28-2-2022 at 10:09

hey i would just like to put out a warning for method similar to which was posted
im trying to find the best and most efficient method for nitrite as i have a way of purifying it and quantifying how much the specific yield for each method happens to be
one of these bags i mixed up happened to be..
20g NaNO3 10g NaOH 4g D-glucose

as from flour 165g - 3.3g
NaNO3 85g 1020 20g
NaOH 40g 480g 9.6g
which i acquired from the original post of this thread
scaled it down to fit the sizes im doing for test
initially i used way too much flour and too little NaOH- of course this turned into a flare, only had about 10g of mixture however. so i fixed the ratios with my bag of nitrate glucose and NaOH and put, 2-3g into my crucible, put it on the gas burner, covered it up with some aluminium foil just for good luck and stood back, eventually left the room as 10 seconds had passed
then all of a sudden a loud explosion came about
https://gyazo.com/5a22662db19e4ff9ed53b7e6d9d24db2

the aluminium foil was blown off, but not torn to shreds and there was barely any smoke
now im having flashbacks to playing around with a mixture that explodes when being melted down, yellow powder i believe it goes by, its a mixture of KNO3 K2CO3 and S
could i just have created a more OTC friendly version?

"Yellow powder is potassium nitrate, potassium carbonate and sulfur melted together in a 3:2:1 ratio. "
and i was using 20-10-4 NaNO3 NaOH d-glucose
with the loud explosions im gonna have to let this one method go for producing nitrite, i might give it a go in aqeous solution- not sure if i dare yet to try with flour instead

LuckyWinner - 10-3-2022 at 11:14

why is nobody talking about this?

'Thermal decomposition of sodium nitrate under isothermal conditions at around 600 °C is sequential reaction, which is NaNO3 → NaNO2 → Na2O.'

1.how long do you place the NaNO3 into a furnace, till no more oxygen bubbles escape?
2.how do you prevent further decomposition of the created NaNO2?
3.it may be better to use KNO3 → KNO2 since its easier to separate unreacted KNO3 from KNO2?

Quote: Originally posted by Antiswat

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

Fantasma4500 - 10-3-2022 at 18:05

KNO3 thermal decomposition? im in the nightmarish midst of experimenting, turns out my bucket of NaNO3 is contaminated with SODIUM CHLORATE
this gave me some funny results for about a week, including loud explosions with NaOH and glucose

KNO3 thermal decomposition, i can share some already now here, the nitrite test i prepared with about a teaspoon of FeSO4 (bought, nice grade) and 20% H2SO4
keeping KNO3 molten at 550*C for about an hour i got a very vague positive, but a positive the none the less, so it does decompose at that temperature, quite slowly- but generally they write its 650-850*C, some places faultily mentioned 300something but thats sodium nitrate

i never quantified how much NaNO2 is formed by my quite careless decomposition reaction- but i might get to collect some real data eventually, probably not a whole lot, but indeed some. you might wanna reach a high temperature and hold it there for a bit, and ignore the whole logic of science that it will decompose when reaching x temperature

only positive with KNO3 thermal decomp is that KNO3 isnt very soluble compared to KNO2 so you can do a crude seperation if you have great yields. its a furnace reaction youre looking at though.
i also looked into sodium/potassium oxide formation and it was at +850*C, so barely feasible unless-- unless you have a furnace. which i might get to once i have enough of messing around with hotplates, or they get enough of me messing around with them. if anyone wanna take a step ahead of me its quite doable, dampening some sand with sodium silicate, drying it out, preferably over fire. ive tested different compositions of this material and found just straight up sodium silicate/sand to be the best, calcium silicate mixes just gave crumbly texture, CaSiO4 is easily made by CaCl2 + NaSiO4 in case anyone wanna try some with that.

one last note, i dont think KNO2 is that soluble in ethanol after all, sometimes they list 4g/100mL as "soluble", my latest research didnt imply ethanol as a way of purification of nitrites.

Nitrate reduction via starch

BAV Chem - 12-3-2022 at 08:22

Over the past few weeks of playing around with this reaction I came up with the following procedure:

100g of finely powdered sodium nitrate and 25g of starch are thoroughly mixed together and placed in a metal can of about 1L in size. The can is heated from below (in this case using an electric hotplate) until the mix ignites. With a slight pop the can starts to put out clouds of white smoke and some sparks and in about 10s the reaction mix quickly turns from a white powder to a grey melt. After the reaction the contents are allowed to cool and dissolved in water and the solution is filtered to remove suspended carbon and other insolubles. Now the yellow solution is boiled down until it starts to bump horribly or something precipitates, both indicating it's saturated. On cooling in an ice bath a white solid precipitates (likes to supersaturate) which is filtered off. This precipitate is mostly sodium carbonate (from decomposition) and some remaining nitrate. Then some dilute nitric acid is added to the solution with strong stirring until a slight odor of NO₂ is detectable and the PH is around 9. This is done to neutralize traces of NaOH. Once again the solution is boiled to saturation, chilled and filtered. The filtrate is then simply evaporated to dryness (boiling will only cause it to bump horribly and make a mess).

Most of the nitrate seems to be consumed using this method. The product obtained is around 50% pure, yield is around 30%. I also came up with a way of measuring its purity.

For this 200mg of product and 1g of urea are dissolved in 30ml of water and 5ml of dilute nitric acid (~20%) is added all at once with strong stirring. The amount of gas generated is measured. This reaction takes some minutes so it may be heated a little but allow the gasses to cool and contract again. In theory 200mg of nitrite should generate 130ml of gas.

I'm uncertain if this test is any good so some kind of confirmation would be appreciated. It seems to be somewhat accurate (+-5%) and I already used the nitrite to make 5-ATZ. Also this test is easily thrown off by carbonate impurities so ideally a control run without any urea is carried out (no heating needed!) to correct for carbonate.

[Edited on 13-3-2022 by BAV Chem]

Fantasma4500 - 13-3-2022 at 17:35

as for testing nitrite, the best shot ive found at it so far is to first up take an amount of sodium nitrite and dissolve in water, then dilute that down until it with 20% H2SO4 and FeSO4 doesnt bubble

i took 1g NaNO2 and added to 100mL water, 1mL of this was reacted with 1mL 20% H2SO4 / FeSO4
this solution is 1% NaNO2 or 1000ppm NaNO2
however it doesnt really form visible bubbles- so 1% NaNO2 is around the cutting line for bubbles to be seen

so if you take 1g of sample NaNO2 and dilute that down until you hit the point where it stops bubbling when reacted with the nitrite test- or maybe go by how far it has to go before the solution stops being transparent- this will however depend on how much solution you have
best bet to eliminate the carbonate would be to use Ca(OH)2
it can end up forming some alkali hydroxide and that can then end up reacting with CO2 in the air to yet again form carbonate.

edit: using flour/Ca(OH)2/KNO3 i just had a minor breakthrough, hoping to write out something useful soon on getting quite pure nitrite and a collection of methods ive attempted.

[Edited on 14-3-2022 by Antiswat]

Keras - 14-3-2022 at 10:07

Reading this thread, I was wondering if potassium nitrate could be reduced to potassium nitrite using a mild reducing agent such as sodium thiosulphate.
So, not taking care to weight/measure anything, I dissolved a bit of potassium nitrate in water (less than a gram of KNO₃ into 10 ml of distilled water), then added some crystals of sodium thiosulphate. I was hoping to see either a cloud of colloidal sulphur form, but nothing happened.

I put the beaker in the fridge, and when it was cold I added a dozen drops of HCl in the hope of seeing the blue tinge of nitrous acid. However, the solution turned cloudy with sulphur, so definitely sodium thiosulphate was oxidised at this point. Since I couldn't see any blue colour (in any case, the nitrous acid would've been far too dilute), I added a few drops of benzyl alcohol (that’s all I had at hand) in the hope that it would be nitrated by the acid.

So here I am: the upper solution is almost clear. All the colloidal sulphur has fallen to the bottom of the beaker at this point. At the bottom, I also have a ‘pearl’ of benzyl alcohol, which is partly covered by something yellow. Is it sulphur trapped at the interface between the two liquids? Is it nitrobenzyl alcohol? I have no clue.

unionised - 14-3-2022 at 11:43

"Thiosulfates are stable only in neutral or alkaline solutions, but not in acidic solutions, due to disproportionation to sulfite and sulfur, the sulfite being dehydrated to sulfur dioxide:

S
2O2−
3 + 2 H+ → SO2 + "S" + H2O"

From
https://en.wikipedia.org/wiki/Thiosulfate

So there's no evidence of an oxidation.

Keras - 14-3-2022 at 12:42

Quote: Originally posted by unionised

"Thiosulfates are stable only in neutral or alkaline solutions, but not in acidic solutions, due to disproportionation to sulfite and sulfur, the sulfite being dehydrated to sulfur dioxide:

S
2O2−
3 + 2 H+ → SO2 + "S" + H2O"

From
https://en.wikipedia.org/wiki/Thiosulfate

So there's no evidence of an oxidation.

Ah, thanks, I didn't know that (I didn't even bother to check on Wikipedia). So you’re right. It’s just thiosulphate disproportionating. Some stronger reductant might be in order. I suppose triphenylphosphine would work, but it’s not cheap. Sodium dithionite?

clearly_not_atara - 14-3-2022 at 13:52

You might want to check the electrode potentials. I suspect that reducing nitrate to nitrite without overreduction to NO/N2O is not so easy. Nitrite is only stable in alkali, nitrate ion is not very reactive, so you have a catch-22 situation.

Maybe rextalizng kno3 is a more worthwhile endeavor? Chlorate is an unwelcome guest.

Fantasma4500 - 15-3-2022 at 07:39

@keras
dithionite is very difficult to handle

as for electroreduction its done with pure platinum (deposited platinum doesnt work well for some reason, it has to be solid) and the concentrations/amperage is very low
quite impractical. if you wanna use electricity just get a high voltage powersupply and do oswald, add in airpump and blow it directly into NaOH solution, this will acquire you half and half nitrate/nitrite however.

ELECTROREDUCTION
THE PRODUCTION OF NITEITE FROM NITRATE.
The reduction of nitrate to nitrite can be accomplished satisfactorily, and the process is the subject of a recent patent.1
It has been shown (Miiller and "Weber)2 that in a divided cell, smooth platinum or copper cathodes reduce nitrate to nitrite and ammonia, but platinised platinum gives much ammonia and little nitrite. A spongy copper or silver cathode was found to give the best results. With a current density of 0.25 amps, per dm3 and a concentration of 2.3 grams of sodium nitrate per litre, a current efficiency of 90 per cent, was obtained. The current efficiency with an amalgamated copper cathode was found to diminish when 50 per cent, of the nitrate had been changed.
Considerable care is evidently needed to prevent the formation of ammonia, since it has been shown by W. H. Easton3

Keras - 15-3-2022 at 22:21

Quote: Originally posted by Antiswat

@keras
dithionite is very difficult to handle

How so?

Quote: Originally posted by Antiswat

as for electroreduction its done with pure platinum […] smooth platinum or copper cathodes[…]

Does copper work? Between solid platinum and ‘spongy’ copper, I doubt anyone would hesitate.

Keras - 15-3-2022 at 23:08

Still, I’m surprised that the sodium dithionite method doesn’t work. If I follow the CRC Handbook, we have:

NO₃⁻ + H₂O + 2 e⁻ ⇋ NO₂⁻ + 2 OH⁻ e = 0.01 V
2 SO₃⁻ + 2 H₂O + 2 e⁻ ⇋ S₂O₄²⁻ + 4 OH⁻ e = -1.12 V

So logically, this reaction:

NO₃⁻ + 2 OH⁻ + S₂O₄²⁻ → NO₂⁻ + 2 SO₃⁻

should have e = 1.13 V, which is a pretty high K and should be almost quantitative, no?

[Edited on 16-3-2022 by Keras]

Fantasma4500 - 16-3-2022 at 08:37

@keras iirc chemplayer had a video on it, its on bitchute
he mentions that it heats up with contact with air because its so strongly reductive that it exothermically reacts with oxygen in air
i would consider anything towards P2O5 difficult to handle.
one other problem is that nitrite can also react so you dont want it to be too powerful
im hearing lead is really great for this reaction as it stops at nitrite- my experiences with it is that its very tedious to be stirring
ive tried with sawdust from bandsaw and in wide excess it gives some success, maybe coffeegrinding it would give better results, otherwise metal workshops usually have a beltgrinder, they have a collection tray underneath where one may find mainly steel dust with some Fe3O4- Fe3O4 reacts very fast with HCl so a quick minute rinse in warm HCl should do- decantation and flushing with EtOH would yield a great reagent

CaSO3 is listed as typical reagent for this reaction, i tried with calcium sulfamate but that didnt do much in my attempt
sulfite is doable as you may simply dump SO2 from a sulfur candle into Ca(OH)2- which should be doable from CaCl2 and NaOH, CaCl2 is commonly sold as air dessicant granules

Keras - 16-3-2022 at 12:45

Alright, I tried potassium nitrate + sodium hydroxide + sodium thiosulphate.

There’s a slight cloud forming when you combine potassium nitrate and sodium hydroxide in solution, but it clears up almost immediately.
So far, I’m faced with nothing much to report. Pretty everything has dissolved, and I didn't notice anything special happening. The reaction, however, should be undetectable with the naked eye.

What is the best way to know if nitrite has really formed? A test that would evidence nitrites and not nitrates. Evaporate the solution and examine the salts? Try to acidify it to form nitrous acid? Anything else?

S.C. Wack - 16-3-2022 at 16:11

Nitrite can be determined quantitatively by titration of a solution with permanganate in warm dilute sulfuric acid: 2KMnO4 + 5KNO2 + 3H2SO4 -> 5KNO3 + 3H2O + 2MnSO4 + K2SO4

[Edited on 17-3-2022 by S.C. Wack]

Keras - 16-3-2022 at 22:26

Thanks!

I put my hands on Vogel's Qualitative Inorganic Analysis, and there’s at least two experiments I could follow with what I have at home: the HCl method to produce the blue nitrous acid, and oxidation of KI.

Your method is also mentioned in there.

Fantasma4500 - 17-3-2022 at 12:00

i would suggest adding in an alcohol to absorb the NOx- rather have a bit of vasodilation than NOx.

i came across sodium dithionite- OTC. i wasnt yet aware that this is OTC so ill be trying to see if this does something interesting for nitrite
any other interesting methods are welcome.

Lionel Spanner - 17-3-2022 at 13:52

I may well be very late to the party, but has anyone tried the method described in US patent 729,515, namely heating sodium nitrate with calcium hydroxide and graphite, giving sodium nitrite, calcium carbonate, and water vapour?

The author reported that repeating the process a second time reduced any unreacted nitrate to negligible quantities.

Attachment: US792515.pdf (139kB)
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BAV Chem - 18-3-2022 at 00:10

Quote: Originally posted by Lionel Spanner

I may well be very late to the party, but has anyone tried the method described in US patent 729,515, namely heating sodium nitrate with calcium hydroxide and graphite, giving sodium nitrite, calcium carbonate, and water vapour?

The author reported that repeating the process a second time reduced any unreacted nitrate to negligible quantities.

This looks like another very doable variation of this process as the reaction with normal charcoal is too exothermic, leading to a lot of decomposition.

On another note, what is the calcium hydroxide even there for in this reaction?
Some include it while others leave it out completely. I had some success with starch and no hydroxide and got a 30% yield (see a few posts above).

[Edited on 18-3-2022 by BAV Chem]

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