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Author: Subject: Very OTC Sodium NItrite
AJKOER
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[*] posted on 5-10-2017 at 14:11


Photolysis run:


20171005_141947-480x640.jpg - 46kB
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[*] posted on 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?




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[*] posted on 6-10-2017 at 03:35


or stop taste testing the ever clear for potency!
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[*] posted on 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]
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[*] posted on 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]




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AJKOER
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[*] posted on 6-10-2017 at 08:02


Pure KNO3 is sold as stump remover aid in stores with home garden sections (Home Depot,...).
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[*] posted on 6-10-2017 at 13:47


Updated picture following photolysis in sunlight for 6 hours:

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

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

20171006_174244.jpg - 344kB

[Edited on 7-10-2017 by AJKOER]
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[*] posted on 9-10-2017 at 09:17


Tried preparing Isopropyl nitrite from the nitrite i made. Yield was 30%.
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[*] posted on 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.




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[*] posted on 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.




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[*] posted on 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]
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[*] posted on 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]
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[*] posted on 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]
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[*] posted on 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.
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[*] posted on 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]
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[*] posted on 14-10-2017 at 10:56


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



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

Nonsense.
https://en.wikipedia.org/wiki/Devarda%27s_alloy
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[*] posted on 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




~25 drops = 1mL @dH2O viscocity - STP
Truth is ever growing - but without context theres barely any such.

https://en.wikipedia.org/wiki/Solubility_table
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[*] posted on 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.:(
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[*] posted on 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.
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