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Author: Subject: Preparation of ionic nitrites
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[*] posted on 7-7-2010 at 01:14


The lead reduction method is useless crap IMHO. I once read in an old chemistry book that several runs are required to obtain nitrite of acceptable purity. That is, melt nitrate with lead, extract the fused mass, remove nitrate by crystallization, evaporate, melt with lead again... repeat at least three times. This might be acceptable on a industrial scale but not for the amateur experimenter. I have tried this method twice now and each time there was a HUGE amount of nitrate crystallizing from the liquid upon cooling. Nitrite has such a high solubilty that as a rule of thumb, if anything crystallizes from the solution then the reduction was not even close to completion. I never cared to dry and weight the remaining nitrate but I roughly estimate that at least 50% of the nitrate remained unreacted.
The lead/nitrate melt was stirred excessively in the beginning but the mixture quickly solidified and made stirring impossible.
As a sidenote, reduction of thiocyanate with several metals yields cyanide, and this was once proposed by Erlenmeyer as a industrial process. It was noted that lead performs rather shitty compared to other metals due to its low melting point. The blobs of molten metal has a much smaller surface and intimate mixing is not possible to the same extent as with powdered zinc or iron. I believe the same effect comes into play when nitrate is reduced with lead. If higher yields and better purity is the goal, then we need a procedure where the reactants can be mixed prior to melting.

I suggest doing some research into alternative reducing agents such as carboxylic acid salts, iron powder or copper(I)oxide. Dont waste your time with the lead reduction. Using powdered lead might remedy the situation somewhat but lead shot/plumbing lead is definetly a dead end.
A promising experiment using calcium formate was described here:

http://www.versuchschemie.de/topic,13865,-Herstellung+von+Ni...

The reaction is vigorous but not explosive. Intimate mixing of the reactants is required to obtain a good yield. Using Ba/Ca formate/oxalate makes sure that oxidation products are insoluble and nitrite of high purity is obtained by simple filtration.

Iron formate/iron oxalate are also worth a try.

Powdered iron also reduces nitrate, producing some ferrate which can be decomposed into insoluble iron hydroxides by boiling.

[Edited on 7-7-2010 by Taoiseach]
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[*] posted on 7-7-2010 at 10:19


I did some more experimenting. A solution containing dissolved KNO3 and about 5-10% HCl was made and aluminum foil added. The smell of nitrogen dioxide became obvious immediately above the solution, but a brown color was not vissible. A variation was tried with more concentrated HCl solution and a brownish color developed in the solution and bubbles appeared with a strong burning smell, but still no brown gas. Next, solid HH4NO3, as much as would dissolve, was added into HCl (5-10%) solution. A small strip of Foil was added, and this time some brown gas developed over the solution, but more excessive generation of NO2 remained elusive. Hydrogen could still be ignited above this last solution, indicating that
2Al + 6 HCl --> AlCl3 (hydrate) + 3H2 was still the primary reaction

I think Nitrogen dioxide oxidizes hydrogen at ambient temperatures, so this probably limits much NO2 from developing.
Nitric oxide might be safe from reaction with hydrogen. I am unsure how to get rid of the hydrogen though. Nitric oxide would react with oxygen in water to form nitric acid, so using oxygen to oxidize all the hydrogen (through the nitric oxide that acts like a catalyst) would not work.
Perhaps by using dry chlorine (passing it through baked/powdered CaCl2), the hydrogen will be oxidized to HCl gas and nitrosyl chloride will be left over. If excess chlorine is not used, and the gas allowed to sit for a minute, so that all the chlorine is able to be reduced, this could then be bubbled into pure anhydrous acetone to possibly obtain nitroso-acetone ONCH2COCH3. Or it could be bubbled into a Na2CO3 solution to obtain NaNO2, NaCl also forming. These are just ideas.

If I could get the generation of hydrogen to be more diffuse, this might prevent any HNO2 or NO2 that forms from immediately getting reduced. This would require aqueous metal ions that would react with acid to give off hydrogen, as the reaction on solid metal is probably too concentrated.
Perhaps Cu+ and H2S would form CuS , H+ and 1/2 H2, driven by precipitation of insoluble CuS ?

I remember reading that reaction of dilute HNO3 on Cu, actually produced traces of NH4NO3.

[Edited on 7-7-2010 by Anders Hoveland]
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[*] posted on 7-7-2010 at 10:49


""The lead/nitrate melt was stirred excessively in the beginning but the mixture quickly solidified and made stirring impossible.""""

Then your issue is too low of a heat source. You must make sure that your Nitrate/Nitrite is melted as well else little is going to happen. Keep in mind that you ARE dealing with a duel phase reaction here and without rapid stirring and making sure to add small portions at a time you get Nitrite and Nitrate encased in Litharge.

Keep in mind these melting points and see why NaNO3 is so much easier to reduce the KNO3. It sounds to me that you tried to reduce KNO3 without using enough heat. NaNO3 proceed pretty rapidly and with ease but its harder to work up so its a trade off of sorts. Lead dont boil till 1749 °C so heat that shit up as high as you can get it and it will go along just fine.

327.46°C Pb
334°C KNO3
308°C NaNO3

440.02°C KNO2
271°C NaNO2

I honestly feel if someone can't get the Pb method to work its because there just not doing it right.

Quote:
I believe the same effect comes into play when nitrate is reduced with lead. If higher yields and better purity is the goal, then we need a procedure where the reactants can be mixed prior to melting.


Try ballmilling Nitrate with Pb shot prior to the melt then. I performed an experiment a while back because I read an abstract that mentioned the reduction performed this way successfully but I never received the full paper so I went ahead and added excess Pb shot to damped nitrate and Course silica in ballmill and let it run for a couple days. I added the silica in hopes that it will expose more surface area but I never quantified the results because they seemed to show a marked reduction in NOx production on addition to H2SO4. I still think that route warrents attention but I have not gotten around to experimenting with it more and completing it with a melt could be exactly what it needs. It was to messy for me to workup hence the reason I never really attacked it from all angles.

@anders

I will have to look at my notes but I did a few simular experiments a while back IIRC and if you use NaOH/Al NH3 is produced in abundence so there is a good chance your not getting your desired results because you are reducing the nitrate all the way to amine and locking it away as Ammonium chloride.

[Edited on 7-7-2010 by Sedit]





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[*] posted on 8-7-2010 at 14:07


@Sedit

>Then your issue is too low of a heat source. You must make sure that your Nitrate/Nitrite is melted as well else little is going to happen. Keep in mind
>that you ARE dealing with a duel phase reaction here and without rapid stirring and making sure to add small portions at a time you get Nitrite and
>Nitrate encased in Litharge.

The oxidation product of lead is PbO and its not going to melt at any reasonable temperature. The mixture will solidify towards the end of the reaction and there's nothing you can do about it - except using a huge excess of nitrate.

>Lead dont boil till 1749 °C so heat that shit up as high as you can get it and it will go along just fine.

No it wont. The nitrite will decompose and leave you with useless sodium oxides/hydroxide.

>I honestly feel if someone can't get the Pb method to work its because there just not doing it right.

Nobody has done any sort of analysis/estimation of purity so far. I believe that what has been produced so far by this method is nothing but a bunch of nitrate with *some* nitrite in it. Adding HCl says almost nothing about purity, as even very small amounts of nitrite will produce red fumes.

>Try ballmilling Nitrate with Pb shot prior to the melt then. I performed an experiment a while back because I read an abstract that mentioned the
>reduction performed this way successfully but I never received the full paper so I went ahead and added excess Pb shot to damped nitrate and
>Course silica in ballmill and let it run for a couple days. I added the silica in hopes that it will expose more surface area but I never quantified the
>results because they seemed to show a marked reduction in NOx production on addition to H2SO4. I still think that route warrents attention but I
>have not gotten around to experimenting with it more and completing it with a melt could be exactly what it needs. It was to messy for me to workup
>hence the reason I never really attacked it from all angles.

This is a very interesting method and much more promising than the high-temperature reduction.
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[*] posted on 8-7-2010 at 15:11


The solidity of NaNO3 melt is only semi riggid at completion and small portions on addition take care of that so even though there is a lumpy mass at the end much of it is nitrite. Honestly when it comes to the reduction of KNO3 however I really do not like this method at all. I have to have yeilds somewhere around either here or The Vespiary for the Pb reduction. If I never posted them or can't find them i'll just have to perform the reduction again. I could use some nitrite anyway as I want to replenish depleting supplys of chemicals. What about adding SiO2 to the lead prior to the addition to help breakup the melt?

The ball milling is just not up to scale for me due to the small ballmill I have so its for nothing more then an experimental. If there was a way to ballmill and melt at the same time it would get my seal of approval over the range of nitrate cations.

Quote:
No it wont. The nitrite will decompose and leave you with useless sodium oxides/hydroxide
.
Ihave allowed the melt to sustain melted for extended period of time but I have never see a marked decrease although it would be interesting to find a means to yeild pure hydoxide to see how much decomposition takes place. The Na should decompouse in the mix as its being oxidised but it does not appear to do so.

Adding HCl is not a great sign of nitrite and H2SO4 seems to provide a better color from the mixture when dripped in.

As a side note a side product of the oxidation iv been messing around with using HNO3 and EtOH appears in all shapes and formes to be a mixture of nitros acid and acetic acid it takes on the distinct blue color of nitros acid but smells strongly of AcOH. Not quite GAA status but not really wet either. Its honestly quite a problem for me at the moment. Large amounts of NOx is produced in the first stage of the oxidation.





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[*] posted on 8-7-2010 at 23:36


Quote: Originally posted by Taoiseach  
I suggest doing some research into alternative reducing agents such as carboxylic acid salts, iron powder or copper(I)oxide.


Be careful what reductant you use, sodium acetate and potassium nitrate will explode violently when melted together according to the Ber. article posted in the cyanides thread on the nitrite method of NaCN. The same reference noted a delayed explosion in a melt consisting of a molecular mixture of sodium formate and NaNO2.

There is also some old method on forming it by heating nitrate alone: the KNO3 is glowed strongly, then the mass is solubilized in hot water. After 24 hours, the liquid is poured off from precipitated KNO3. Then neutralized with dilute acetic acid, and mixed with the double volume of ethanol. After several hours, three layers form, the middle layer contains a yellowish oily liquid, from which the KNO2 is obtained by evaporating over H2SO4, N.W. Fischer (Pogg. Ann. 150 [1848] 116).
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[*] posted on 9-7-2010 at 03:33


>There is also some old method on forming it by heating
>nitrate alone: the KNO3 is glowed strongly, then the mass
>is solubilized in
>hot water. After 24 hours, the liquid is poured off from
>precipitated KNO3. Then neutralized with dilute acetic acid,
>and mixed with
>the double volume of ethanol. After several hours, three
>layers form, the middle layer contains a yellowish oily liquid,
>from which
>the KNO2 is obtained by evaporating over H2SO4, N.W.
>Fischer (Pogg. Ann. 150 [1848] 116).

KNO3/NaNO3 doesn't decompose cleanly; there is an equilibrium formed between nitrate decomposing into nitrite and nitrite decomposing into various oxides. That must be the reason they add acetic acid - it converts the oxides/hydroxides into acetate which is decently soluble in EtOH whereas nitrite is not.

>Be careful what reductant you use, sodium acetate and
>potassium nitrate will explode violently when melted
>together according to the >Ber. article posted in the
>cyanides thread on the nitrite method of NaCN. The same
>reference noted a delayed explosion in a melt
>consisting of a molecular mixture of sodium formate and
>NaNO2.

I recently had a go at the Ca formate reduction and its not explosive tough exothermic and a bit violent. I tried again with an equal volume of Fe dust added. It made the reaction more controllable and also improved thermal conduction. When the reaction set it the whole pot started glowing dark red for a few seconds; lots of fumes were given off and some hissing but no flame. The mass was left to cool, extracted with hot water, filtered to remove CaCO3 and precipated with EtOH. The nitrite formed a heavy yellow oil with EtOH; this was evaporated to dryness. Btw cyanides also form this "oil" (probably some sorta gel) with EtOH.
Yield is poor tough because the solubility of nitrites and cyanides in EtOH increases drastically with the amount of water present.

>The ball milling is just not up to scale for me due to the
>small ballmill I have so its for nothing more then an
>experimental. If
>there was a way to ballmill and melt at the same time it
>would get my seal of approval over the range of nitrate
>cations.

Lead sponge can be precipated from lead acetate by adding zinc metal strips or zinc powder. After torough washing and drying it should be possible to intimately mix it with the nitrate prior to melting.





[Edited on 9-7-2010 by Taoiseach]
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[*] posted on 9-7-2010 at 08:50


Quote: Originally posted by Taoiseach  
KNO3/NaNO3 doesn't decompose cleanly; there is an equilibrium formed between nitrate decomposing into nitrite and nitrite decomposing into various oxides. That must be the reason they add acetic acid - it converts the oxides/hydroxides into acetate which is decently soluble in EtOH whereas nitrite is not.


Sodium and potassium nitrate are very few of the alkali compounds that readily yield their oxides on thermal decomposition. It would probably also be a good idea not to do this in glass or platinum (who uses it anyway?).

Quote:
Lead sponge can be precipated from lead acetate by adding zinc metal strips or zinc powder. After torough washing and drying it should be possible to intimately mix it with the nitrate prior to melting.


Concerning a method where spongy Cu-powder made from CuSO4-soln. and zinc dust, is melted together with KNO3 is described below. Then there is also something by O.L. Erdmann (J. pr. Ch. 97 [1866] 387 footnote) where KNO3 with an excess mass of iron filings is molten in a cast iron crucible and heated to moderate glow.

Attachment: J. chem. Soc. Lond. 36, 595.pdf (176kB)
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[*] posted on 21-7-2010 at 11:03


Hi folks!

I looked through the thread a couple of months ago, and based on the comments I had not even tried the molten lead method.. It seemed so messy and troublesome.. I like to reproduce hard to reproduce processes @ home:) but grinding lead, filtering lead, melting lead, three things that I would avoid at any cost, I just stay away from anything that involves lead, and I suggest this attitude to all of you!

Nitrosyl sulphuric acid(NOHSO4, chamber crystals) replaces NaNO2 in most of the cases, so make that instead of XNO2, It also stores well in a well stoppered glass bottle, and relatively easy to make, (easier than any XNO2 shit)

Take an aspirator and connect it to the end of a setup that consists of a heated quartz tube, in which you burn sulphur, the quartz tube is connected to a bubbler full of H2SO4 the escaping SO2 from the bubbler is led to cc HNO3, it takes some time and the reaction mixture requires cooling but NOHSO4 is not that hard to make afterall. After about a hour yelow cristals starts to precipitate if the solution is properly cooled. One drop of the resulting yellow liquid makes a shitload of N2O3 if it is added to water.

I found a book on www.archive.org search for: The Aromatic Diazo-compounds And Their Technical Applications, its a nice book with preparative examples, including the use of nitrosyl sulphuric acid.
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[*] posted on 21-7-2010 at 11:52


Thats is interesting Jimmy no doubt about it deserves more attention.

For some people Bisulfite's might be a better means at SO2 production then a sulfur burner even though in the long run sulfur would be cheeper. I assume the conentrated H2SO4 is a means of drying the SO2 but that should pose no problem at all.

The concentrated HNO3 could be made with H2SO4 and KNO3 then frozen and cold filtered to remove as much KHSO4 as possible so I see much potential for this being highly over the counter.

Its there any risk involved if the nitric acid used is red from NOx contamination?


[Edit]
This statement in a patent on the synthesis of nitrosyl chloride seems to contridict your statement that addition of the nitrosylsulphuric acid to water generates the desired gas.

Quote:

Solutions containing more than 40% of nitrosylsulphuric acid crystallize when the water content exceeds 10%. To overcome this disadvantage, it is then necessary either to reduce the content of nitrosylsulphuric acid, which is reflected by a decrease in the productivity, or to heat the solution to a temperature of the order of 50 to 100° C., which is expensive.

The presence of water causes, in the more or less long term, hydrolysis of the nitrosylsulphuric acid, which correspondingly decreases the yield of nitrosyl chloride. Hydrolysis becomes more significant as the temperature rises.


However even though it seems to suggest water will crystalize the nitrosylsulphuric acid it seems that it does so slowly and does not 100% hinder the formation if this patent is to be trusted.

Reference: http://www.freepatentsonline.com/6238638.html



What about reacting KNO3 in excess H2SO4 and then chilling the solution significantly and slowly adding Sodium Metabisulfite to generate SO2 insitu.

H2SO4 + Na2S2O5 -->> 2 SO2 + Na2SO4 + H2O

H2SO4 + KNO3 --> KHSO4 + HNO3

2HNO3 + 3SO2 + 2H2O --> H2SO4 +2NO


Yes I know the Stoichiometry leaves alot to be desired at this point but im just museing here so bear with me. I dont know much about this compound right now and am speaking as I learn.

[Edited on 21-7-2010 by Sedit]





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[*] posted on 21-7-2010 at 18:15


My scale has had abit of a meltdown due to a H2SO4 accident a few weeks back so it will be a little bit before I can give anything conclusive but I thought I would share this.

I took KNO3 and dissolved it in a test tube of 97% H2SO4 making sure to have excess H2SO4 there. I then c chilled this in an Ice/NaCl bath to lower the temperature down as much as I could with ease. There was a slight reaction as the addition was slowly made and at this point is became a thick syrupy consistancy. Unaware if anything at all had happened I let it sit in the ice back for a few hours as I went about my own things. I came back to a clear'ish solution with needles of more then likely bisulfites crystalising out. I decided to drip H2O in at which point large amounts of light yellow gas as formed and quickly oxidised to dark NO2.

This looks very promising to me but alot more test need to be done to determine if the gas was indeed N2O3.

[EDIT]

In the morning there is a thin layer of yellow liquid, resting on top o all the precipitated salts that where in the reaction, that resembles all photos I have seen of nitrosylsulphuric acid. Addition of H2O indeed generates a large amount of gas production that by all means appears to be N2O3.

Is there any solvent I could use to isolate this compound from the salts to run better, purer test with it?

[Edited on 22-7-2010 by Sedit]





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[*] posted on 22-7-2010 at 08:27


SO2 generation by means of K2S2O5 is not that good idea IMHO, if you drip ccHCl to solid K2S2O5 then the KCl cristallizes on the surface of the K2S2O5 crystals because of the endotherm reaction, you can heat the flask of course but a cake will form and it will prevent agitation, magnetic stir bar is useless here. If you dissolve the K2S2O5 in water and drip cc HCl into that, the SO2 will remain in the water, you have to heat the solution to push out the SO2, but it means more H2O vapour in the resulting gas, and your bubbler with H2SO4 will quickly become uneffective. So burn sulphur with an aspirator, I also tried to burn sulphur in a direct air flow, by means of an aquarium pump, but is also poses many problems, aspirator is the way to go!

You have to use dry SO2 here, or you will continously hydrolize the NOHSO4 formed during its production.
I do not have experience with H2O+NOHSO4 system, but I noticed that, in the book the H2SO4 was always used in large excess to the NaNO2, this is probably due to the water formed in the reaction, which has to be 'absorbed' by the excess of H2SO4!

It should be noted however when NOHSO4 added to water it makes a very exotherm and vigorous reaction. I made some N2O3 volcano by injecting 1ml NOHSO4 in H2SO4 to water, I bet If you could see that you would hardly believe that NOHSO4+H2O mix even exists in any proportion(it also new to me:).

In brauer preparative inorg. chem fuming nitric is used, and I also used fuming nitric without any problem.

You probably start with 95% HNO3, in the initial stage of the reaction the NOHSO4 formed will react with the 5% water present to give H2SO4 and N2O3, as the reaction proceeds the H2O will be consumed, and H2SO4 wont form anymore till the end of the reaction, depeding on the moisture content of your SO2 of course. You can follow the temp. of the reaction mixture by the colour, if it becomes intense orange you have to cool it till it go back to yellow.

I do not really know how could you isolate the NOHSO4 from H2SO4+NaNO2. Prepare relatively conentrated NOHSO4 from fuming HNO3+SO2 and filtrate the crystals on a glass frit, but you do not have to isolate it to use it.
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[*] posted on 22-7-2010 at 09:21


Dripping H2SO4 onto slightly damp Metabisulfite is working wounders for generating SO2 with ease. Theres good odds that this will also remove a large portion of the water from the reaction.

Iv been running small scale test tubes since you posted this and im pretty confident this can be done with Bisulfite as the SO2 source and quite possible insitu. The substance that appeared over night was light golden liquid and fumed vigerously with water generating a blue solution if only a small amount of water was used. A light yellow gas was formed but on heating the solution formed dense brown NOx fumes.

When given a chance I want to run it the right way as a control to be sure further experiments are what I am looking at.

[Edited on 22-7-2010 by Sedit]





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[*] posted on 23-7-2010 at 10:12


Sedit Yes the bisulphite method works, it is in the book:)

Today I made isopropyl nitrite by NOHSO4 and it worked well.

I took about 100ml izopropyl alcohol and diluted it with 20ml water, I chilled the mix to 5°C started to agitate while slowly dripped the NOHSO4+H2SO4 into the mixture (I dunno what concentration, I sucked it up from the white NOHSO4 crystals) every drop made a hissing noise but no nitrogen oxides formed, all the gasses absorbed by the alcohol.
After the addition of about 10ml NOHSO4 solution I poured the whole mixture into cold water, and a yellow layer separated, as usually.. I smelled it got headache etc. so the method works, but needs to be optimised ofc.
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[*] posted on 12-2-2011 at 00:28


Quote: Originally posted by kmno4  

NaNO2 content is easy to determinate using KMnO4/H2SO4.
(add slowly [with stirring], NaNO2 solution to KMnO4/H2SO4 sol. untill it becomes colourless)

I've been looking into decent spot tests to differentiate between nitrites and nitrates. How does this one work? Does the MnO4- ion oxidise nitrites but not nitrates?
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[*] posted on 13-2-2011 at 08:27


Quote: Originally posted by Polverone  
I tested a slightly modified version of Muspratt's nitrite preparation method this afternoon. 100 g of KNO3 were mixed with 12.1 g of charcoal and the whole thing ball milled for about an hour. I tested a little bit of the mix and found that it would burn without an external heat source. I poured the mix into a stainless steel dish and ignited it.

The reaction wasn't terribly fast, due to the great excess of oxidizer, but it was fairly vigorous. There was a lot of bubbling and splashing of the molten salt since my dish was barely large enough to hold the charge of powder. When burning finished I placed a sheet of copper over the top of the dish and waited for it to cool.

I had different results. 10g KNO3 was grinded and mixed with 1.2g activated charcoal with a mortar and pestle until a fine gray powder was obtained. I tried igniting the mixture with a lighter but all that happen was the powder being flamed crackled and turned black. I had to blowtorch the mix and keep heat on it to keep the reaction going. Balls of white molten salt rolled around on the mix and ultimately I ended up with a dirty gray coloured solid clump. This is the same colour that the KNO3 turned when I melted it by itself in a boiling tube. This is even after I recrystallised it so I don't really know whats going on.
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[*] posted on 13-2-2011 at 09:12


What was the source of the KNO3? How did you recrystallize it?

(I guess I'm wondering if it really is KNO3. People have used stump removers that were not KNO3, thinking that they were KNO3.)

It sounds as if you ground KNO3 and charcoal together in a mortar. Please tell me you didn't do that.

[Edited on 13-2-2011 by entropy51]
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[*] posted on 13-2-2011 at 09:44


Quote: Originally posted by entropy51  
What was the source of the KNO3? How did you recrystallize it?

It sounds as if you ground KNO3 and charcoal together in a mortar. Please tell me you didn't do that.

My friend gave me about 20g of KNO3 that he got on ebay. I recrystallised by adding boiling water until the KNO3 was fully dissolved then cooled to around 5C and collected the crystals.

I did ground the nitrate with the charcoal in the mortar. It was outdoors and I was ready to deal with it if it ignited. This isn't very different to ball milling is it?
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Rosco Bodine
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[*] posted on 26-8-2011 at 04:32


Quote: Originally posted by JohnWW  
Quote: Originally posted by Nicodem  
Quote: Originally posted by Taoiseach  
Details are in French pat 388,563 which I was unable to find tough.

You probably mistyped it or something as it is right where it is supposed to be: FR388563 (also available in German version as DE203751)

When one tries to save PDFs of those French and German patents of 1908, v3.espace.com throws up popup windows that immediately cause Firefox to crash. Has anyone been able to download them, somehow? If so, please post them here are attached files.


Here are the patents attached FR388563 and DE203751

Check the archive for the old E&W forum as there were numerous additional patent processes listed there.
I'll try to dig those up again and repost them here.

Attachment: DE203751 Nitrite from Nitrate and Formate.pdf (59kB)
This file has been downloaded 659 times

Attachment: FR388563 Nitrite from Nitrate and Formate.pdf (71kB)
This file has been downloaded 599 times

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terraxus
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[*] posted on 30-8-2011 at 10:58


why not just heat NaNO3?
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Alastair
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[*] posted on 7-10-2011 at 12:33


Quote:
Quote: Originally posted by madcedar  


Sodium nitrite is made in two steps, firstly ethanol and nitric acid is used to make ethyl nitrite, then ethyl nitrite is reacted with sodium hydroxide to make sodium nitrite.

2C2H5OH + N2O3 = 2C2H5N02 + H2O
C2H5ONO + NaOH = C2H5OH + NaNO2




Why not just react N2O3 with NaOH directly?
And btw this method isnt bad for any impurities, you can purify gas from excess NO or NO2 by putting it into your freezer (one becomes solid, the other stays gas while N2O3 remains in liquid form)
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[*] posted on 8-10-2011 at 09:44


Can reaction of (molten) Potassium nitrate with Zinc make Potassium nitrite ??

If yes, then i think it would be like :

KNO3 + Zn = KNO2 + ZnO :D

ZnO is insoluble in water, while nitrite is soluble




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niertap
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[*] posted on 16-4-2012 at 18:41
KNO3->[RED]->KNO2 effective reductions


Has anyone tried the nitrate to nitrite reduction using iron filings? I've heard heating nitrates with lead works well, but heating them with iron filings seems like a better and quicker method. Would it work? If not what would be formed?

[Edited on 4-17-2012 by Polverone]




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[*] posted on 16-4-2012 at 21:07


There are no experimental details of using iron, hence asking if anyone has tried and not posted here about it. Any way tried the reduction with lead you do have to stir frequently and is superior because of lower mp than if you were to use iron. But if iron is all you have go for it. Post back whatever findings you have, would love to hear the outcome.

[Edited on 4-17-2012 by Polverone]




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[*] posted on 16-4-2012 at 21:27


I saw a little bit about it in a thread about the creation of K2FeO4, but nothing concrete.

I've been thinking about that and from user experiences it seemed like K2ferrate, iron oxides, and or nitrates could form. I'm assuming higher heats cause two nitrites to react with an atom of iron to form the ferrate; lower heats possibly just causing the nitrate -> nitrite reduction, producing Fe2O3.

The completely thermal approach appears to the reaction seems unpredictable, so I am currently trying an aqueous approach. Transition metals seem to be very good at catalyzing oxido/reducto reactions.

[Edited on 4-17-2012 by Polverone]




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