potassium nitrite synthesis [Archive] - The Explosives and Weapons Forum

kingspaz

July 27th, 2002, 07:39 PM

for the past few weeks i've been working on DDNP as a primary. everything can be made or aquired easily apart from a nitrite salt. i have posted this new topic for people to just throw ideas nomatter how extraordinary or insane (may provoke thought) they sound for the preparation of HNO2 or a salt of it such as KNO2. i am particularly interested in preparing it from KNO3 as this is readily available. i have tried using Al and KNO3 which seems to work but the yields are really poor. also is there a method to prepare NO and NO2 from one reaction?
any help will be appreciated.

megalomania

July 27th, 2002, 09:59 PM

While nitrites have never interested me, I do recall that nitrates can be decomposed into nitrites by heating. I doubt it may be that simple, but then again it may not be too far off. You may have to mix the nitrates with something to liberate the oxygen.

mongo blongo

July 27th, 2002, 10:42 PM

Yes heating with elemental lead to remove the oxygen. It will give PbO and a metallic nitrite (if you used a metallic nitrate).

kingspaz

July 28th, 2002, 08:12 AM

i am aware of both of those processes as i recall nbk a while ago saying if you left KNO3 molten long enough it would gradually decompose. also the lead process, why lead? is it due to the low melting point? other than that i cannot see any reason why lead would be used since there are many much more reactive metals which could be used more effectively. also is there a way to generate NO and NO2 simultaneously as this may prove easier than synthesising a nitrite.

mongo blongo

July 28th, 2002, 11:01 AM

Lead is a good choice because the melting point is low and the reaction is easy to control. There are some more reactive metals but using something like Al or Mg with say KNO3 is a bad idea. Also Pb is very easy to find and quite cheap.

kingspaz

July 28th, 2002, 04:39 PM

mongo, have you had any success with the process? if so can you please describe the procedure you used, ratios of KNO3 and Pb aswell as heating time....i think heating time is the crucial thing.

[ July 28, 2002, 05:13 PM: Message edited by: kingspaz ]

xoo1246

July 28th, 2002, 06:00 PM

Aren't you thinking of sodium nitrite?(the more noble the metal the more metal oxide you get)
Then you separate the nitrite from PbO by dissolving it in distilled water I assume.
Edit: From The Hive
<blockquote>quote:<hr />On a laboratory scale, the melt-phase reduction of alkali nitrates (preferably NaNO3 is used) with metallic lead is suitable for the production of nitrites:

NaNO3 + Pb ==> NaNO2 + PbO

One melts 1 mol (85 g) sodium nitrate in an iron crucible of approx. 15 cm diameter and 1 mole (207 g) of metallic lead (cut into small pieces) is added slowly with good agitation. After all the lead is added, the mixture is agitated for a further 30-45 minutes until all the lead is oxidized. Careful agitation is important for the achievement of a good yield. After cooling the reaction mixture, the solid cake left is cut up into smaller pieces and placed in a larger iron container. The formed sodium nitrite is extracted with 200ml of warm water, followed by two 75ml portions of water, and the insolubles is filtered off (Caution: Lead oxide is very poisonous!). To remove remaining lead from the aqueous solution, carbon dioxide is bubbled through for one minute (not longer!) and the insolubles again filtered off. After neutralization with dilute nitric acid, the solution is concentrated under vacuum. The first crop of crystals which precipitates consist largely of sodium nitrate. Now the solution is evaporated to dryness, and the still molten residue (consisting of sodium nitrite contaminated with some sodium nitrate) is washed briefly with ethanol. Care must be taken not to heat the sodium nitrite much higher than its melting point (271°C) as the compound decomposes at 320°C, with evolution of nitrogen oxides. <hr /></blockquote>

[ July 28, 2002, 05:19 PM: Message edited by: xoo1246 ]

Polverone

July 29th, 2002, 12:50 AM

I tried a number of times to prepare potassium and sodium nitrite from metallic lead and the nitrates before I gave up and just ordered the NaNO2. The nitrites of both K and Na (IIRC) decompose at nearly the same temperature as the nitrates decompose to form the nitrite. So it was very tough for me, since I wasn't able to keep the temperature very controlled. One could (I suppose) keep the temperature at a reasonable level by repeatedly withdrawing the heat source until the liquid begins to solidify, then replace it, so that the material is always held near its melting point. But I didn't have enough patience for that, especially since I found a place that would sell me NaNO2 for $3.00 a pound. I'd just e-mail you the name of the dealer but you're in the UK so it'd do you no good :(

When I made my attempts with lead I heated the material strongly and stirred with a spoon. It looked like the lead was oxidizing nicely but when I filtered my reactants I saw that most of the lead had remained unchanged. The material I obtained gave only the faintest haze of red gas when I added HCl to it, and pure NaNO2 gives you copious red-brown clouds upon addition of HCl. I made some attempts with lead and some without, and the ones without it actually seemed to work better, but I can't guarantee that since the heating conditions were so uncontrolled.

Gleaned from the Hive:

US patent 792,515
2NaNO3 + 2NaOH + C = 2NaNO2 + Na2CO3 + H2O
2NaNO3 + Ca(OH)2 + C = 2NaNO2 + CaCO3 + H2O
(CaCO3 is insoluble in H2O)

US patent 670,021
NaNO3 + CaO + SO2 = NaNO2 + CaSO4
KNO3 + K2SO3 = KNO2 + K2SO4

If you get frustrated with the lead-reduction procedure those patent methods might be worth trying.

kingspaz

July 29th, 2002, 07:03 AM

hey thanks guys!
i'll look at those patents and if neither of the processes are workable i'll have an attempt at the lead process as xoo has posted.
thanks again!

edit: i tried one synth from the patent following this equation:
2KNO3 + Ca(OH)2 + C ---> 2KNO2 + CaCO3 + H2O
i used graphite as a carbon source as the patent recomends it. also it is pure carbon which charcoal is not. the reaction seems to have worked as upon addition of the carbon a scentless gas was produced. i think this was water vapour. water vapour continued to be evolved for around 2 minutes. i then allowed it to cool and dissolved what i assume to be KNO2 out with a little water then passed through a filter. i then boiled the water off. i tested it by adding it to a little picramic acid in weak H2SO4. the picramic acid didn't dissolve which was gay so i added the KNO2 anyways. i got a couple of little flecks or orange brown precipitate and a funny smell (decomposed HNO2?). for dissolving the picramic acid would concentrated H2SO4 be a good idea then add this to water containing KNO2?
thanks for the help.

[ July 29, 2002, 04:26 PM: Message edited by: kingspaz ]

Polverone

July 30th, 2002, 02:18 AM

You assume you got KNO2? Well I would test it first, myself. Dump a bit of the solid into HCl. No red fumes, no KNO2. I'd be interested if this method worked out well for you, since I like to keep nuggets of knowledge around for the day when I might not be able to buy certain things.

kingspaz

July 30th, 2002, 10:24 AM

yes, i forgot to add into the edit on my last post 'i know i shouldn't assume things'. the problem is i can't get HCl and i have run out of H2SO4 for the time being. i will however make up a large batch using this process and i will test it when i have some H2SO4. thanks for the help!

andreas

July 30th, 2002, 10:53 AM

As a means of temp. controll use a bath of molten lead as the heat sorce. It would keep the lead molten but wouldn't overheat the reaction.

Polverone

July 30th, 2002, 04:11 PM

Do you have any acids? Citric, acetic, oxalic, phosphoric? Any of these should be plenty strong enough. I would dilute the H2SO4 if you use it for testing, since concentrated H2SO4 will produce fumes to some extent even with nitrates.

kingspaz

July 30th, 2002, 04:46 PM

hey i do have some citric lying round actually. i can't remember why i bought it. thanks polverone i'll use that to test. the art place doesn't get the graphite sticks in for 2 weeks though :(....assholes...
i do have 3 though so thats about 30g of graphite which is enough for a decent amount of KNO2. thanks for the help. i may not get this done for a few days though because of the rain.

megalomania

July 31st, 2002, 01:24 AM

Could you tell us what patent number that is?

Polverone

July 31st, 2002, 01:57 AM

Yes I could. In fact, I already did <img border="0" title="" alt="[Wink]" src="wink.gif" /> . Look above the equations in my earlier post.

kingspaz

August 4th, 2002, 10:50 AM

right, the great rain that has persisted the last 6 days, day and night has stopped! so today i'm going to make some KNO2 via the last method i tried. using graphite as a carbon source. also before i tested the KNO2 i made with the Al method:
2Al + 3KNO3 ---> Al2O3 + 3KNO2
i added about 10-15g of the crude product to a concentrated citric acid solution. the stuff fizzed a fair bit and the solution turned orange and began to fume vigourously. the fumes had an orange colour and did not leave the flask but stayed at the bottom with the liquid.
about half the crude KNO2 i added did not react or dissolve so i assume that would be KNO3 impurities. but as i have leanrt i shouldn't assume <img border="0" title="" alt="[Wink]" src="wink.gif" />
i'd just like to know whether an orangy colour is acceptable in this chemical test as i've never had experience with this.

nbk2000

August 4th, 2002, 03:33 PM

The orange gas sure sounds like NOx to me.

kingspaz

August 6th, 2002, 10:49 AM

right,
i've just purified the KNO2 from the slag left over from the carbon method. i have then tested it using a strong citric acid solution as i did above. the results where disapointing as when i added the 'KNO2' nothing happend. it just sank to the bottom and began to dissolve which tells me its most likely KNO3!
back to the my old Al method i think...

Microtek

November 6th, 2002, 01:43 PM

I would like to recommend a procedure for nitrite production that I have just discovered ( it's not my idea, I just read a book ).
The method is derived from Griess's original procedure for DDNP synthesis and can be used directly if desired.
Griess prepared DDNP by bubbling nitrous gas through an alcohol solution of picramic acid at low temp. The "nitrous gas" is an equimolar mixture of NO and NO2 and acts almost exactly as nitrous anhydride N2O3. This gas was generated by reducing HNO3 of density 1.30-1.35 g/mL with starch. Now I have been generating NOx by dissolving Cu in HNO3 which is rather impractical, so I tried out this new method and this is how it went:
20 mL of 50-55 % HNO3 was poured on 2-3 teaspoonfuls of starch in a flask. Nothing happened at first but after shaking gently and heating until the mix turned brownish, bubbles started forming and heat was removed as gas evolution accelerated. From this point, huge volumes of brown gas was evolved over the course of about 5-10 minutes. The evolution was very steady.
( additional info in a while )

kingspaz

November 6th, 2002, 07:06 PM

thats an interesting idea. i suspect the equilibrium is similar to this:
NO2 <=> N2O4
so:
NO + NO2 <=> N2O3
do you have any information as to exactly how the starch reduced the HNO3? if this was bubbles through KOH then you should get KNO2 solution (KNO2 purified by evapouration) so HNO2 can be made as needed as KNO2 is storable.

Marvin

November 6th, 2002, 08:53 PM

I'm not happy with the current methods of estimating nitrite. I think reacting it with an ammonium compound, with some catalytic TM ions around to make sure it goes to completion and measureing the amount of nitrogen gas produced would work well.

Reacting nitric acid with copper or starch is supposidly a good method for small amounts, there is no definate single reaction going on, and the ratio of NO2 to NO in the product depends on the concentration of the acid as well as the reducer, the more dilute the acid, the more NO. For very dilute nitric acid you eventually get hydrogen with active metals. Get the ratio right and its quite possible to produce nitrite in which there is no chemically detectable amount of nitrate formed. As it reacts and the nitric acid dilutes the ratio changes to excess NO, which is wasteful but doesnt contaminate the product.

Id think nitrous acid should be acidic enough to displace carbonate, which is a lot cheaper than potassium hydroxide, if you need K salts. If you cool an equimolar amount of NO and NO2 enough you get solid N2O3 crystals. Its questionable whether this formation is crucial to formation of nitrite in solution though.

Microtek

November 7th, 2002, 04:00 PM

I made use of this easy way of producing NO/NO2 by bubbling the gas through a solution of 4.80 g NaOH in demineralized water for about 2-3 minutes. I then evaporated the water until the crystals were completely dry ( by heating and passing dry air over the soln ) and weighed the material. A total of 8.1 g was produced. I did a qualitative test for nitrite by adding 50 % sulfuric acid to a soln of the material and got nice foaming with a brown gas.

Marvin

November 8th, 2002, 10:07 PM

Its nice it works, but what you got might be contaminated with a lot of unreacted NaOH. Theres a few niggling questions I havnt been able to find answers to. Will a solution of nitrite in water absorb oxygen from the air.... and would a mixture of nitrate and hydroxide absorb NO with the production of nitrite......

If the latter is true, then we could make maximum use of nitric acid if we choose to use that method, and if the former is true then multiple recrystalisations might not actually be increasing the purity of the nitrite.

I think the 'best' method of determining quantity of nitrite might simply involve titration in mild acid with permangonate. Titrations are spectaculaly dull, but do give very reliable results.

Microtek

November 9th, 2002, 08:20 AM

I'm not sure what would happen if pure NO was bubbled through NaOH soln, but it is quite easy to determine if NO2 is present in the gas stream simply by looking at the colour. Air could then be added to the stream if too little NO2 was present.
Or two gas generators could be set up, one with a high concentration of acid, the other with a low concentration. The gas streams would be merged, and bubbled through the NaOH soln. As the concentrations dropped the low conc generator could be shut down while the high conc one progressed to the NO/NO2 production concentration. Once it was diluted further, it could be used as the low conc generator etc. etc.

PS. How would you use this titration? What indication would there be that stochiometric amounts had been added?

[ November 09, 2002, 07:23 AM: Message edited by: Microtek ]

Marvin

November 10th, 2002, 10:44 PM

If a solution of nitrate and hydroxide would produce nitrite from NO, this would be a big improvement. Firstly the yeild of nitrite from nitrous gasses would be doubled, and secondly its much easier to use dilute nitric acid, which only produces NO without adding air.

I'll refersh the idea of titration for those who dont know. You have an indicator in solution that tells you if the chemical you are testing for is still present, then you add a carefully made concentration of a chemical that destroys it, a bit at a time until its all gone. From the equation of the destruction, and the amount of destroyer used, you can work out the amount of chemical your interested in knowing the amount of.

This is a bastardisation of a very wide ranging method, but it will serve. In the case of nitrite with permanganate, we produce nitrate, and depending on the ph, probably either manganese dioxide, manganous oxide or manganese chloride. Since the permanganate is deep purple itself, we dont need an indicator, we just add it, preferably from a buiret, but a gradiated cyclinder will do, until we get a faint perminant purple colour in the solution.

Normally, we'd do a double titration to work out the total amount of undecomposed permanganate used from a reagent that doesnt decompose, but this isnt required. What we do need is a method of weighing quite accuratly the permanganate to form the known concentration solution. Any sort of balence will do. I suspect the product in mild acid would be Mn (II). Permanganate stains can be removed with conc citric acid solution. I apologise if this is an overly simplified answer Microtek, but others might find it useful.

I suspect nothing would happen if NO was bubbled through hydroxide on its own, hyponitrates can be made, but require rather more subtle methods. My idea that it would reduce nitrate is based on NO2 + NO producing only tiny amounts of nitrate if balenced correctly, if the production of nitrate was irriversable, I would expect a fairly sizable amount of nitrate formed.

At the moment I'm lookiing at decomposition of TM metal nitrates to NO and NO2 and the metal oxide. The quantities are often right for pure nitrite production, owing to the valence of oxygen relative to nitrate. They decompose at temperatures reachable with pyrex containers. The nitrates themselves are often impossible to prepare free of water, but if a mixture of sodium nitrate and the metal chloride is used the anhydrous TM nitrate is formed and decomposed in situ to leave the metal oxide and sodium chloride. Any chlorine/HCl formed would contaminate the product, but this should not be large, and easily removable if needed. This could work, this could definatly work.

Microtek

November 11th, 2002, 05:03 AM

Just for the record that is an over simplification, but I don't take offense quite so easily so no apology needed.
I wasn't asking what a titration is but how you would indicate the transition from nitrite to nitrate. I wasn't thinking that MnO4- would be reduced to colorless Mn(II)2+, silly me.
Anyway, I tried this and it works brilliantly though you need to acidify the soln very carefully, otherwise the reduction will be along these lines:

2[MnO4]- + 3NO2- + H2O --> 2MnO2 + 3NO3- + 2OH-

Where it should have been

2[MnO4]- + 5NO2- + 6H+ --> 2Mn(II)2+ + 5NO3- +3H2O

So of course you can use both reactions, but one or the other and if the first one is used pH must not be too high or the reduction will proceed too slowly.

BTW, what do you mean by TM? Transition metal? If that is the case then I'm pretty sure it will work. At least I think I read somewhere that Cu(NO3)2 decomposes to give NO2 + NO + CuO. And when you prepare the coppernitrate you get a lot of NOx as well.

frogfot

August 23rd, 2003, 04:54 PM

Tested the KNO3/Pb method. Had 10 g KNO3 and 20 g lead in a beaker, heated with gas burner for 30 min and stirred with glass rod. Whole thing melted and turned to yellow-brown (a sign of PbO). Not all lead was consumed. When i left beaker to cool, it cracked when salt solidified (duh!), thats why theyre using iron cup.. It was hard to separate all Pb and PbO so I didn't weight them..

Anyway, nitrite was extracted with hot water. On test with HCl sample turned yellow and bubbled slightly, NOx smell was present :) The clear solution was left for some hours and it became turbid, most probably PbCO3 forming in contact with airs CO2..

I dont understand following:
carbon dioxide is bubbled through for one minute (not longer!)

What is the reason not to bubble it longer? The dissociation constants show that HNO2 (4E-4) is much stronger than H2CO3 (4,31E-7), so evolution of HNO2 couldn't be the reason..

Btw, another less user friendly alternative to CO2 would be H2S, this would give much purer product, though this wouldn't be nescessary for most purposes of nitrite...

Trinitrotoluene

August 23rd, 2003, 05:47 PM

A while ago I tested the KNO3/Pb method. I used 41grams of Pb metal from fishing weights, and 20 grams of KNO3 from stump remover. It was placed into a copper crucible and heated under a hotplate. I did manage to get the lead of melt and it stayed at a molten state for 5 minutes, but the hotplate wasn't hot enough. My goal was to obtain PbO. The nitrate/nitrite was washed away with water, I only received a fairly thin layer of PbO covering the Pb.

DBSP

August 23rd, 2003, 06:03 PM

During my attempt to synthesise CTMTNA I tried making sodium nitrite by reducing it with lead. And I got a fair ammount of PbO, I might still have the PbO left in the filter paper, check that later. I heated the NaNO3 and lead together in a melting pot used for casting lead over a propane burner. And I was able to maintain a fairly high temp because of the gas burner. As soon as the NaNO3 had melted(when melted alone) it started releasing its oxygen(noticed small bubbles at the hottest places).

mongo blongo

August 23rd, 2003, 10:10 PM

I have tried NaNO3/Pb many times. I would definitely advise heating the nitrite extract with Pb twice or more (to react any remaining NaNO3). I can tell you this because the amount of PbO collected was less than half than expected after simple stehiometric calculations.( Not using nitrite weights because of the uncertainties of the product). This was after only the first reaction by the way.
I tested my nitrite presence with Iron II Sulphate and distilled acetic acid (gives brown solid suspension which is dissipated with heating).
I do admit that I haven't actually used my nitrite in any synthesis except Cyclotrimethylenetrinitrosamine in which the result was very convincing. I never tried it with DDNP though.

frogfot

August 29th, 2003, 09:05 AM

70 g Pb/51 g KNO3 gave 43 g unreacted KNO3 after 40 min of heating.. and occasional stirring. I think that continious stirring is very important to drive reaction faster. Next experiment will be on bigger scale with a motor stirrer..

Mongo, whats your usual yield of nitrite?

Btw, what if one heats KNO3 with some widely available catalyst? For example, would iron(III) or manganese(IV) decompose nitrite too?

mongo blongo

August 29th, 2003, 10:41 AM

Can't remember any yields sorry dude. It was some time ago. How did you know how much KNO3 reacted? It looks exactly the same as nitrite.

frogfot

August 29th, 2003, 03:51 PM

Well, I evaporated extracted liquid to 50 ml and precipitated KNO3 at 0*C (it cann't be nitrite because of its solubility..). Remaining 10 ml was evaporated to dryness, got about 3 g solid, from which <1,3 g is KNO3 (calced from its solubility).

tiac03

November 26th, 2003, 02:25 AM

the only place i read about getting nitrite is from the book "CIA explosive-preparations" since i havn't tried any of the methods in it i can't tell you if it actually works. (or if anything they say is safe to do) i downloaded it off of kazaa. (pdf file someone scanned the manual and put it into pdf format)

(since the manual is about improvised explosives they use easy to find containers.)
an exerpt:

Mix 12g of lead and 4G of potassium nitrate (or sodium nitrate). place in an iron pipe (with an endcap) and heat with a blow torch or hot coals for an hour or more.

remove container from heat, allow to cool and then chip out the yellow solid (with screw driver).
put in a mason jar add 1/2 cups (120 ml) of methyl alcohol to it. (orange brown solid, cream color in alcohol)

heat it in a pan of hot water untill it reacts (solution turns darker upon heating)

filter mixture through paper towel. solid left on paper towel is lead monoxide. wash it twice with 1/2 cup hot water. then let dry before using in explosives (such as lead picrate).

the liquid you filtered the lead monoxide out of (in a jar) is then put into a hot water bath untill all alcohol has evaporated. the remaining powder is nitrite.snowy liquid; some white powder appears but not much.

well that is a paraphrased version of what is written in the manual. i can't tell you if it will work or not.

and i guess they use lead (to answer your question) because you can then use the lead monoxide to make lead picrate (used as a primary explosive in detonators)

metal dragon

March 4th, 2004, 02:10 AM

While mindlessly flicking through a chemistry book I found something interesting. It was preparation of potassium nitrite. It was simple example as well potassium nitrate was put into a small boiling tube (test tube) and a Bunsen burner was used to heat this sample, O2 was given off, this was proven by the ignition of a glowing splint. The mixture that remained was in fact KNO2 . This procedure was taken from a series of books called “chemlab”.

Mr Cool

March 27th, 2004, 04:44 PM

Now, why didn't everyone else think of that five years ago?

Oh, wait, we did.

I have a little experiment to do with sodium metabisulphite, results posted when I have them...

kddflx

April 18th, 2004, 12:45 AM

here's an idea for reducing HNO3 to HNO2 using Hydrogeniodine from 'Van Nostrandd's Scientific Encyclopedia Fifth Edition':

under the heading Iodine, it reads: "Hydrogen Iodine, HI, is the least stable of the hydrogen-halides, and correlatively, the best reducing agent, readily reducing vanadic acid, nitrous oxide to ammonium, nitrous acid to nitric oxide, and HNO3 to nitrous acid."

nbk2000

April 19th, 2004, 01:18 PM

frogfot

September 24th, 2004, 04:59 AM

I've recently tested Pb/KNO3 synth with 99,95% lead and got some really strange results.. after about 12 min heating 207g Pb/101g KNO3 and stirring constantly the mix started to heat up on its own until a red glow! So I removed the gas burner and mix continiued glow for about a minute.. During reaction there separated some solids (hight mp, probably potassium oxides).
Those who tested this synth many times, is this normal? Is it some kind of runaway..?

On extraction I've got 55 g KNO2 (very basic on dissolution, probably contains alot of KOH because of potassium oxides..) and 215 g yellow powder of PbO. Why doesn't any source mention this runaway??.. :(

Bugger

September 24th, 2004, 06:00 AM

HI, better known as hydriodic acid, is also a schedule 1 meth precursor, making simple possession of it a felony under various drug control acts. Either use it for making crank or, better yet, leave it the fuck alone and find better ways of making HNO2.
I presume you are in the States. Has the U.S.A. gone that Fascist, outlawing even dear old HI? I cannot think of any other country where even use, let alone possession, of HI is illegal. And besides, it can be easily made as a byproduct of iodination of hydrocarbons, or direct reaction with gaseous H2 (made by action of an acid or alkali on a suitable metal), using iodine bought OTC as a disinfectant. If you live close to the Canadian or the Mexican border, how about setting up a laboratory in either of those countries, if you want to keep and use HI?

chochu3

September 27th, 2004, 06:19 AM

Heating a nitrate salt will take off an oxygen but after cooling most of it will suck up back the oxygen giving you the nitrate. I read somewhere in one my books (to lazy to reference) that nitric acid can be reduced to nitrite by reduction with hydrogen iodide. I myself haven't tried this but might be worth time.

September 27th, 2004, 08:26 AM

Jome skanish

September 27th, 2004, 10:50 AM

How about use one of these for preparation of KNO2:

Cr2O3+4KOH+3KNO3--->2K2CrO4+3KNO2+2H2O
or
2MnO2+2KOH+3KNO3--->2KMnO4+3KNO2+H2O

EDIT: These are reactions performed at high temperature, at least one of reactant has to be molten, which for KOH means a temperature of at least 360 degrees centrigrade. Just felt I had to mention this so no one would mistake them for being "wet" or someting...

The first one definately work. Its on frogfots page! The second I think I saw proposed here somewhere, but I dont know if anyone ever successfully performed it.
It seems quite easy to separate the KNO2 from remaining KNO3or KMnO4, KNO2 is a very soluble salt at low temperatures (73,6g/100ml H2O @ 0 degrees centrigrade), and that goes for only two of the possible contaminants. (solubility: KNO3 13,3g, KMnO4 2,83g, they should be easy to separate, but KOH and K2CrO4 is highly soluble, ~60 and 58,2g /100g H2O @ 0 degrees centrigrade.

My idea: First, separate out the low-solubility salts by fractional cristallisation (these are KNO3, KMnO4 or just KNO3 depending on which reaction you chose)
That should leave you with a solution of KNO2, KOH and K2CrO4 or KNO2 and KOH. Adding copper sulfate to the first solution (from the Cr2O3 reaction) should precipate out Cu(OH)2 and CuCrO4, leaving in solution K2SO4 and KNO2. Adding to the second solution (the MnO2 one) should precipate out Cu(OH)2 + a little CuMnO4, Cu(MnO4)2 and so on (manganates and permanganates of copper).
Now, from both experiments we are left with only K2SO4 and KNO2 in solution. K2SO4, as opposed to KNO2, has a very low solubility @ 0 degress centrigrade....

One would probably have to repeat some of the steps to get clean nitrite, but it should work, shouldn't it?

Bugger

September 27th, 2004, 06:30 PM

The problem with the second of those reactions, in which MnO2 is allegedly oxidized to KMnO4 on heating with molten KOH and KNO3 (which is partially reduced to KNO2), is that KMnO4 decomposes at less than 240ºC, below the melting-point of KOH which is 380ºC (and KNO3 melts at 333ºC). And, MnO2 decomposes to Mn2O3 at less than 230ºC. So this reaction is a defnite non-starter. Where did you get this spurious reaction from?

Chromate is much more resistant to heat, though, and nitrate just might be capable to oxidizing Cr2O3 to chromate(VI) while undergoing partial reduction to nitrite, under these conditions.

Jome skanish

September 27th, 2004, 06:59 PM

I dont know if this particular reaction is possible, but the more commonly quoted synthesis for KMnO4, (with MnO2, KOH and KClO3, like here: http://www.roguesci.org/theforum/showthread.php?t=934&highlight=potassium+permanganate ) also seems to be performed at similar temperatures. Whats interesting is that even though the temperature causes decomposing of the KMnO4, the processes that promotes its formation is faster, at least until some kind of equlibrium is reached

However, if one were producing nitrite by the Pb+XNO3 method, simply adding water and filtering out PbO would stil leave XNO3 and XOH as impurities. They could be removed by the exact purification-method described above!

meselfs

September 28th, 2004, 12:41 AM

I've always mixed a small amount of sugar with KNO3, without melting or even powdering them. The amount of sugar must be more then what is used for cheap rocket fuel...

Then ignite. It won't exactly burn, it'll smoke alot, boil, and produce sparks and heat.
If it fails to sustain itself, use more sugar. If it ends up with alot of black crud, use less sugar.

It's crude, but it works, it's always done the job for me.

chochu3

September 28th, 2004, 09:06 AM

The above post I should of metioned the reaction did proceed with lead and formed its oxidize and was stirred and let to bubble for an hour, but the remaining KNO2 did turn back into KNO3. Maybe the reason for this is because the contents was mostly in contact with air as a small wide rxn vessel was used.

nbk2000

September 28th, 2004, 02:11 PM

So burning sugar and nitrate forms nitrite? Are you sure it's not carbonate?

frogfot

September 29th, 2004, 11:44 AM

meselfs, I didn't get what this have to do with nitrite synth..

About how to separate K2CrO4 and KNO2, theoretically there are a simple way.. mp of the first one is 968*C and mp of second is 387*C (thats difference by 2,5 times). So, if only there was an easy way to filter those apart while KNO2 is still liquid..

This shure would be an easier nitrite synth than Pb/nitrate synth (and one would also increase yields of dichromate :) )

meselfs

September 29th, 2004, 01:43 PM

So burning sugar and nitrate forms nitrite? Are you sure it's not carbonate?
Definitely it contains carbonate, but it's mostly nitrite.

Jome skanish

September 29th, 2004, 05:32 PM

Frogfot: Melting it!?
How about adding Fe/CuSO4, filter out the precipate (mostly Cu/FeCrO4, but also some Cu/FeCO3 and Cu/Fe(OH):s.) then cool the solution to precipate the sulfate, which has low solubility at low temps.

frogfot

September 30th, 2004, 02:00 PM

Sounds nice, but I wanned to recover the chromate also to make dichromate.. although maby Cu/FeCrO4 have some good uses too.
It's kinda easy to melt nitrite while keeping chromate solid, even with a gasburner..

Jome skanish

October 1st, 2004, 11:29 AM

Oh, you have tried it! Is it an effective method, does it separate good?

If it is possible to add acetic acid before removing the nitrite (I dont know how acidic soln. has to be to release HNO2?), then it would be much easier to precipate out the formed potassium dicromate, and then use CuSO4 and after that fractional crystallisation to remove KOH,K2CO3 and KNO3 from the nitrite soln. The dicromate, in turn could simply be recrystallised at low temps (everything else is more soluble)

Now you have both quite pure K2Cr2O7 and KNO2 :D

frogfot

October 1st, 2004, 07:35 PM

I havn't tried to separate them.. yet. When melted it becomes like a goo of chromate crystalls soaked into nitrite.. so somekind of suction filtration is required.
Yp, one could recover nitrite after dichromate recovery, but solution releases some NO2, and even more if heated (gotta heat to concentrate it). Thats a bit nasty (and there will be some acetate left).
Neutralising the acid would give more work.. though maby it's worth a try when more time comes (school have just started..).

Jome skanish

October 2nd, 2004, 09:02 PM

I found some chart on solubilities in alcohol and water, it seems like NaNO2 is insoluble in etanol, but Na2CrO4 is slightly soluble (0,34g / 100 ml). If etanol is added to a water solution of the two, perhaps NaNO2 precipates like KNO3 and NaNO3 does when one is making black powder! But first one would have to remove to remove all remaining NaO/KOH, since its soluble in alcohol. Potassium cromate doesn't seem to be soluble etanol, so If potassium cromate is needed you'd have to precipate it out of Na-soln. with KCl or something. I cant find anything on KNO2, but I doubt it is soluble i etanol.

What can we make of this?
Na2CrO4 and NaNO2 is made by fusing together NaOH, NaNO3 and Cr2O3.
They are filtered from remaining Cr2O3.
Possible contaminants are NaOH, NaNO3 and Na2CO3. Some Ca(NO3)2 is added and all carbonates and hydroxides are precipated. These are filtered of.
By adding etanol everything but Na2CrO4 precipates, the solution is filtered once again. The filtrate is mostly NaNO2 and some NaNO3. NaNO3 can be removed too, but in most nitrite uses nitrates does not make any difference. I dont know how they can be separared if it's needed, their solubilites seems pretty "close".

The liquid that went though when the others separated out contains Na2CrO4. The etanol should probably be removed, but that's a piece of cake, just heating. Adding K-ions (by KCl or KNO3) one can crystallise out K2CrO4 at low temps, perhaps more than one time if product purity is important.

Only flaw with this procedure I can see is the difficulty to separate out the nitrate. As soon as I find cheap Cr2O3 in sweden I will attempt it!

EDIT: Perhaps stupid CaCrO4 precipates when adding Ca-nitrate. I dont know its solubility but it "sounds" like a low-solubility salt, anyone with a table?

frogfot

October 3rd, 2004, 07:38 AM

0,34g/100ml is quite low solubility.. for 34 g one would need 10 litre alcohol..

Heres some info on solubility of CaCrO4*2H2O in 100 ml water:
22,4 g@0*C, 16,3 g@20*C, 18,2 g@45*C

I have a bottle with about 500 ml of acidic nitrite containing liquid from last preparation of chromate (about month old), it's black and theres a visible amount of NOx inside the bottle. I would try to extract the nitrite if only there was no NOx..

Jome skanish

October 3rd, 2004, 09:23 AM

Perhaps IPA or acetone could be tested instead if it is more soluble in one them?
Now im just guessing. I guess I should try to actually do something instead...

Ca(OH)2 probably wont form in solutions with only a little NaOH, its precipation requires very high pH's. I tried to precipate it out of Ca(NO3)2 soln with ammonia, but it seems like ammonia is not alkaline enough. NaOH worked only if concentrated, so I guess adding Ca(NO3)2 to remove hydroxides and carbonates wont work :(

Dont you have a fan or something? (!) Or perhaps you mean you nitrite is decompsed?

nbk2000

October 4th, 2004, 02:06 PM

What about seperating the nitrite from the carbonate after sugar/KNO3 combustion/reduction?

The simplier, the better, as this would remove the need for a furnace for reducing large amounts (pounds), by simply adding the mix as a constant feed into a crucible where it's kept burning by the heat of that which has already burnt, jsut fast enough to keep the flame going, but not fast enough to generate a huge smoke plume.

meselfs

October 4th, 2004, 03:26 PM

Well, I can't say exactly how much carbonate forms, but I think it's mostly nitrite.

The idea is that you make your proportions so that you have slightly less sugar (or whatever fuel) to reduce the nitrate to nitrite, mix them well, and ignite.

The carbonate is less anionic then nitrite, thus nitrite is favored.

Any carbonate that does form can usually be ignored, I think. I've never had trouble with it. The only trouble is knowing how much nitrite you have so you can mix your reactants properly... but I always use excess nitrite. It's hard to get in a pure form in general no matter what you use. It even takes oxygen from air (slowly) to turn into nitrate.

bobo

October 5th, 2004, 11:51 AM

According to Shimizu, nitrite is often found in the ash of incomplete burning of nitrates. For instance, the author states that the mixture of sodium nitrate and shellac 10:2 produces liquid residue and strontium nitrate and shellac 10:2 leaves a solid ash. Both residues are mainly nitrite.

For potassium nitrate, the author states that potassium nitrate with an ordinary organic fuel (rosin and shellac are named as examples) burns mainly to nitrite, and it takes carbon and sulfur to get all the oxygen out effectively.

The numbers produced by Shimizu are 0.396 g oxygen from 1 g of KNO3 maximum versus about 0.158g for inferior fuels. 0.396g/g would indicate that K2O is quantitatively formed from the oxidant. 0.158 corresponds stoichiometrically with nearly complete reduction to nitrite.

The_Rsert

October 5th, 2004, 03:28 PM

Has anyone information of oxalic acid (?g in 100ml water at 0°C, etc.)?
Because, maybe you will get some oxalic acid, if you react KNO3 and sugar.
C12H22O11 + 9 O2 ----> 6 HOOC-COOH + 5 H2O
But I'm not sure of this.

nbk2000

October 5th, 2004, 08:44 PM

Well, with hexa, sugar/nitrate, and HCl, you've got HDN, so if the nitrite is somewhat ghetto purity, so what, as long as it works when you need it to, and can make it anytime. :)

bobo

October 7th, 2004, 11:19 AM

why would you use nitrite in HDN synthesis, if nitrates will do? Not that I would ever consider HDN unless as a precursor, but that is due to the cost of making hexamine here.

The_Rsert

October 14th, 2004, 10:09 AM

Has anyone information of making HNO2?

Marvin

October 14th, 2004, 12:13 PM

Rsert, How ignorent of chemistry do you need to be not to know you add HCl to a solution of sodium nitrite in an ice bath? This is very simple chemistry for diazotisation.

If you are wondering how to make nitrites from HNO2, the only real way is to lead the right mixture of NO and NO2 onto a base in solution. The mixture can be made correctly by diluting nitric acid to a specific point and then reacting it with starch.

Heating sodium nitrate with a manganese (II) salt to produce NO and NO2 might also be a useful method but these are generally less easy and require higher temperatures.

The_Rsert

October 16th, 2004, 07:44 AM

Yes, thats clear, but I need another way without nitrite/HCL?
Maybe reducing 20% HNO3 or somwhat like that i mean.
Without reducing nitrates before of course.
Sorry about my bad english.

Marvin

October 16th, 2004, 07:26 PM

Ok, I see.

I thought I had exact instructions for the starch process but these seems to be not true. The method recommended in Mellor is to react nitric acid with starch and feed the gasses into a freshly made slurry of sodium carbonate out of contact with air. The concentration of the acid and temperature needs to be regulated to leave a very small excess of NO which will pass through undissolved. Excess NO2 than required to form nitrite will produce nitrate as well as nitrite, a small excess of NO will form only very very tiny amounts of nitrate. The bubbling needs to be continued until the solution is acid, and then boiled out of contact with air until neutral. It can then be safely exposed to air to evaporate/concentrate in order to crystalise the nitrite. A guide to roughly how concentrated or hot the acid needs to be it doesnt say, you'll have to use the undissolved NO coming through as a guide if you want to try this.

This and the fused nitrate methods seem to be the only reasonable options so far.

Edit, just to clarify for anyone not following,

Reduction of very dilute nitric acid produces NO, reduction of concentrated nitric acid produces NO2, so at some stage of dilution you get 1:1 NO:NO2 you need to make nitrous acid. Temperature will also affect it.

The_Rsert

October 17th, 2004, 04:50 PM

Thanks for the little info, Marvin.
But that's very inefficient.
I will fuse some sulfurless BP in a beaker, than I will mix it with cold water, than I will filter it and than I will recrystallize the filtrate.
And than I will see how pure the KNO2 is...

Rosco Bodine

October 18th, 2004, 03:02 AM

There is another patent regarding the use of sulfur as
the reducing agent for making nitrites from nitrates ,
the temperature of the melt being low enough so that
the sulfur mixes and reacts with the melt , and does not
simply deflagrate on contacting the melt .

US572819

The reactions described for two nitrites are :

3 NaNO3 + 2 NaOH + S -----> 3 NaNO2 + Na2SO4 + H2O

3 KNO3 + 2 KOH + S -----> 3 KNO2 + K2SO4 + H2O

If it is not necessary to produce the Na or K nitrite alone as
a single pure compound , I speculate that it may be easier to
produce a mixed nitrites , from a combined mixture of the
Na and K nitrates and their hydroxides in some proportions
which may lower the melting point of the mixture still further ,
if this is found desirable to facilitate a controlled reduction .

According to another patent describing a eutectic mix of nitrates
for use in a fusible variation of black gunpowder , a mixture of
55 parts KNO3 and 45 parts NaNO3 is a eutectic .

GB715828

Magnesium nitrate or perhaps calcium nitrate , and their oxides
or hydroxides added to the system could also lower the melting
point of such mixtures . Perhaps the reaction requires a temperature
high enough for the melt that the sulfur will also be molten , but
I am uncertain about this factor . It may be that the reaction
is temperature dependant at a different rate for the different
salts and would have to be run at an optimum temperature so that
the different reactions run in reasonable harmony . Or the experimental
use of a eutectic melt may be no advantage . This is simply an
additional idea offered for whatever value it may have .

bobo

October 19th, 2004, 07:45 PM

do you get nitrite/nitrate mixed with polysulfides maybe? I think, this method is based upon sulfide formation and hence the impurity to be expected is sulfide. That may or may not be problematic, but the waste surely stinks to high heavens and even in a fumehood it may rise to unacceptable. It sure got some wrong attention when more than one dead bird was found on the roof near the same location.

Rosco Bodine

October 19th, 2004, 11:39 PM

There would likely be some sodium polysulfide formed in situ as
an intermediate , transition compound , then oxidized to a normal sulfate .
Also it may be that some sulfur dioxide or even sulfur trioxide may be
produced in small amounts , and escape the mixture as fumes , however
the reactions as stated above would predominate at a conservative
temperature . As for the dead birds , chemistry is not a pastime
for the foolish or careless , and has left many dead ducks for proof as well :eek:

Marvin

October 21st, 2004, 02:00 AM

Molten solutions of nitrates and sulphides can detonate, cf Yellow powder, so this doesnt sound like the safest method. If sulphur dioxide is acceptable, or even sulphites there is a method that uses it in aq. solution. The sulphate is precipitated as calcium sulphate but its not clear if this drives the reaction forward.

Overall,
NaNO3 + CaO + SO2 => CaSO4 + NaNO2

Its not 100% clear this is in solution, but it seems to be. So far nothing beats the simplicity and reliability of fusing nitrates with a 15% excess of granulated lead, everything Ive read recommends that process.

"As for the dead birds , chemistry is not a pastime for the foolish or careless "

Or birds it would seem.

Rosco Bodine

October 21st, 2004, 03:24 AM

The danger of explosion in such a mixture would be present
if all of sulfur was mixed together at once with the
oxidizer mixture and an uncontrolled exothermic reaction
proceeded faster and faster , as a "runaway" reaction .

But such a danger would not be present for sulfur added
gradually in small portions . Watching the temperature ,
and controlling it by the rate of addition of the sulfur ,
the reduction should proceed smoothly . It seems possible
that whether or not the sulfur takes fire on contacting the
melt could even be useful as a visual indicator of the temperature .
The sulfur previously added should have already been
converted to sulfate , so the reaction mixture itself is
not flamable and would not be set afire even from the
ignition of sulfur on contact with the surface of the melt
were it allowed to become too hot .

I have never tried this particular method , but using a
mixture of the nitrates and hydoxides which could provide
a low temperature melt should result in a smooth reduction
with little danger of explosion .

March 21st, 2005, 08:25 PM

I was just taking a look at my KNO2 and I was wondering if it's supposed to look like it does. Now from reading its msds, I see "Appearance: White or slightly yellow granules." So far so good. And it has a solubility of 300g in 100g of water, so that would explain the hygroscopicity (I think I spelled that one right). When I take some in my hand it sort of melts when squeezed between my fingers. I'm just wondering if it has too much water in it.

I got it from a friend who gave a me a bunch of his old chemicals and it's a pretty old lab bottle. I would hate to think I'm weighing out 10 grams when in reality it's a lot less and has a large water content.

This is what it looks like:
http://img.photobucket.com/albums/v78/wolfhound/KNO21.jpg
http://img.photobucket.com/albums/v78/wolfhound/KNO22.jpg

Is it good the way it is? Or should I let it sit in a container with some silica gel packets to desiccate? Any help would be appreciated. Thank you in advance.

March 21st, 2005, 08:25 PM