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Author: Subject: Nitric Acid Synthesis
chief
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[*] posted on 28-4-2010 at 04:59


14 EUR ? Hopefully not for just 1 liter ??
==> H2SO4 can be got ..., just not from any store ..., except maybe a "drug-store", where it cost's 5 $/liter ...

Only such stores nearly dont't exist any more, and not each will sell it ...

I myself could get the 30-liter-canister for maybe 20 bucks, but I don't need it ...
==> Anyhow: When somehow roasting the plaster (CaSO4) it should be quite easy to get a bigger amount of the acid, the setup could be let run by itself, with not much attention ...

This way there would be a cheap source for other acids as well: HCl is obvious, but the usual standard-fertilizers could be distilled with the H2SO4 by the bag ... : Would certainly give quite a bit of HNO3 or maybe H3PO4 or whatever ...; this way re-crystallization could be avoided, the more simple distillation would come into it's place ...

Even the tubes of some old TVs could be used as a flask for that, if gently enough heated ...
==> With good thermal insulation and maybe the heat-input via a hot-air-gun from below (so they don't break such a process would require a day for a 25kg-charge, and electricity for maybe 1 or 2 $ ...)

Besides: In Germany, and probably many other states, almost no real chemical can be had from stores any more ...; everything thinned with loads of water, mixed with a ton of other ingredients and packaged in colorful bottles ...
==> Almost only useless crap in the stores ..., and much overpriced ...
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hissingnoise
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[*] posted on 28-4-2010 at 06:01


~chief, I'd happily pay more - making H2SO4 seems such a total hassle.
BTW, how come you're using dollars in Germany?


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[*] posted on 28-4-2010 at 06:06


Would heating the bejeesus out of a metal container filled with KNO3 over a fire and then bubbling the gases through water be a reasonable efficient way to make 60% or better HNO3?

If I'm not mistaken, first the KNO3 would decompose into KNO2 and O2. Then the KNO2 would decompose into K2O and NO and O2. The NO and O2 would react and form NO2 which when bubbled through the water forms HNO3.

What do you think? Worth a try?

KNO3 is the only nitrate I can get in large quantities cheap enough to try this.
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[*] posted on 28-4-2010 at 07:10


Depends on the fertilizer; but if you have clean KNO3, then it's a standard-textbook-way to make HNO3 by distilling it with H2SO4 ...
==> Some excess of H2SO4 is added, I believe the double amount ...,
==> and pure HNO3 distills over ...

I did it a couple of times with NH4NO3, following an old patent from 1911/1913, worked well, only I had use for gallons of the HNO3 and then ordered it, for 65 EUR-ct/kg ... :D:
==> It's said to be easier with NH4NO3 than with NaNO3 or KNO3, because everything in the flask is liquid and therefore no problems with partial overheating etc. occur.

================

When I mention prices here I convert them to $, because I can't find the EUR-sign on the keyboard ...

================

And now: What do most fertilizers consist of ?
==> probably no Chloride
==> maybe some sulfate, which doesn't react with H2SO4
==> maybe some phosphate
==> and maybe some nitrate ...

Don't know if any H3PO4 would distill over ...
==> also certainly not the H2SO4 ...

So what comes through the cooler would be HNO3, maybe H3PO4 (can it be distilled ?) ...
==> Might be a clean way to separate the precious parts from the fertilizer ...

=====================

Whoever suggests paying in excess of 15 $ for H2SO4/1 liter
==> is not a amateur-chemist,
==> but a hobby-chemist ... :o :D

[Edited on 28-4-2010 by chief]
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[*] posted on 28-4-2010 at 07:25


A well-stocked garden centre might have fertiliser marked 12-0-43 which is essentially pure KNO3. . .
And if you have HNO3, why not boil it with sulphur to get H2SO4?

[edit] Actually these two acids are soooo dependant on each other, it's sickening!
The bitches of chemistry?


[Edited on 28-4-2010 by hissingnoise]
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[*] posted on 28-4-2010 at 07:39


If you boil HNO3 with sulfur, absorb the nitrogen oxides in cold water or better cold hydrogen peroxide, to regenerate nitric acid.
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[*] posted on 28-4-2010 at 08:41


Quote: Originally posted by hissingnoise  
A well-stocked garden centre might have fertiliser marked 12-0-43 which is essentially pure KNO3. . .
And if you have HNO3, why not boil it with sulphur to get H2SO4?

[edit] Actually these two acids are soooo dependant on each other, it's sickening!
The bitches of chemistry?


[Edited on 28-4-2010 by hissingnoise]


Never heared of that one ... ; what else can HNO3 be boiled with to give some acid ?
==> Anyhow the sulfur would be even harder to obtain than any of the previously mentioned ingredients ...

===============

I now wonder how the phase-diagram of HNO3/S/H2SO4 would behave under temperature ...
==> maybe 50%-HNO3 could be boiled with S to give H2SO4 with lower water-content, thereby generating higher-grade nitrating-acid by just boiling HNO3 with sulfur ??

Teh 69%-azetrope of HNO3/H2O surely is different in the presence of H2SO4, so that either in the boiling- or in the receiving- flask would be some higher-concentrated HNO3 possible ... ?

================

It also might cause thoughts of the reactins during the explsion of black-powder: The S-vapor might partially react with the NO2 of the nitrates to some H2SO4, which might play some role ... ... ?

[Edited on 28-4-2010 by chief]
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[*] posted on 28-4-2010 at 09:02


Quote:


It also might cause thoughts of the reactins during the explsion of black-powder: The S-vapor might partially react with the NO2 of the nitrates to some H2SO4, which might play some role?

The basis of the lead-chamber process, sans the charcoal. . .

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[*] posted on 28-4-2010 at 09:19


Quote: Originally posted by Jor  
If you boil HNO3 with sulfur, absorb the nitrogen oxides in cold water or better cold hydrogen peroxide, to regenerate nitric acid.


Instead of boiling HNO3 with sulfur, would heating the crap out of KNO3 work? That would produce oxides of nitrogen as well right?
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[*] posted on 28-4-2010 at 10:05


Thermal decomposition of KNO3 gives K2O, nitrogen and oxygen. . .
No NOx is formed!

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[*] posted on 28-4-2010 at 10:29


Well that sucks :(
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[*] posted on 28-4-2010 at 10:37


I know the nitrates of Ca and Cu do produce NO2 on strong heating. . .


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[*] posted on 28-4-2010 at 10:47


But I can't get these nitrates. (well not cheaply anyway)

I don't suppose one can make Ca or Cu nitrate from K nitrate?
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[*] posted on 28-4-2010 at 13:58


The nitrates of 2-valent Metals produce NOx upon thermal de-composition, those of 1-valent metals don't ...
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[*] posted on 28-4-2010 at 19:38


Heating NH4NO3 has to be done very carefully, as the salt is prone to unstable decomposition, esp. if contaminated; thus synthesis via NH4NO3 is not recommended pathway.

Quote: Originally posted by chief  
Depends on the fertilizer; but if you have clean KNO3, then it's a standard-textbook-way to make HNO3 by distilling it with H2SO4 ...
==> Some excess of H2SO4 is added, I believe the double amount ...,
==> and pure HNO3 distills over ...

I did it a couple of times with NH4NO3, following an old patent from 1911/1913, worked well, only I had use for gallons of the HNO3 and then ordered it, for 65 EUR-ct/kg ... :D:
==> It's said to be easier with NH4NO3 than with NaNO3 or KNO3, because everything in the flask is liquid and therefore no problems with partial overheating etc. occur.

================

When I mention prices here I convert them to $, because I can't find the EUR-sign on the keyboard ...

================

And now: What do most fertilizers consist of ?
==> probably no Chloride
==> maybe some sulfate, which doesn't react with H2SO4
==> maybe some phosphate
==> and maybe some nitrate ...

Don't know if any H3PO4 would distill over ...
==> also certainly not the H2SO4 ...

So what comes through the cooler would be HNO3, maybe H3PO4 (can it be distilled ?) ...
==> Might be a clean way to separate the precious parts from the fertilizer ...

=====================

Whoever suggests paying in excess of 15 $ for H2SO4/1 liter
==> is not a amateur-chemist,
==> but a hobby-chemist ... :o :D

[Edited on 28-4-2010 by chief]
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[*] posted on 28-4-2010 at 19:40


Solid amorphous sulfur is surprisingly has a quite inert aqueous chemistry, except in alkaline condition. I've never had any experience with this, but this pathway might not be as easy as it sounds.


Quote: Originally posted by hissingnoise  
A well-stocked garden centre might have fertiliser marked 12-0-43 which is essentially pure KNO3. . .
And if you have HNO3, why not boil it with sulphur to get H2SO4?

[edit] Actually these two acids are soooo dependant on each other, it's sickening!
The bitches of chemistry?


[Edited on 28-4-2010 by hissingnoise]
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[*] posted on 28-4-2010 at 19:43


The main sulfur reaction during the combustion of blackpowder is an oxidation by KNO3 (and air if it's in open space) to form mainly SO2.
Quote: Originally posted by chief  
Quote: Originally posted by hissingnoise  
A well-stocked garden centre might have fertiliser marked 12-0-43 which is essentially pure KNO3. . .
And if you have HNO3, why not boil it with sulphur to get H2SO4?

[edit] Actually these two acids are soooo dependant on each other, it's sickening!
The bitches of chemistry?


[Edited on 28-4-2010 by hissingnoise]


Never heared of that one ... ; what else can HNO3 be boiled with to give some acid ?
==> Anyhow the sulfur would be even harder to obtain than any of the previously mentioned ingredients ...

===============

I now wonder how the phase-diagram of HNO3/S/H2SO4 would behave under temperature ...
==> maybe 50%-HNO3 could be boiled with S to give H2SO4 with lower water-content, thereby generating higher-grade nitrating-acid by just boiling HNO3 with sulfur ??

Teh 69%-azetrope of HNO3/H2O surely is different in the presence of H2SO4, so that either in the boiling- or in the receiving- flask would be some higher-concentrated HNO3 possible ... ?

================

It also might cause thoughts of the reactins during the explsion of black-powder: The S-vapor might partially react with the NO2 of the nitrates to some H2SO4, which might play some role ... ... ?

[Edited on 28-4-2010 by chief]
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[*] posted on 28-4-2010 at 19:48


It is SAID, that most of transition metal nitrates liberate NOx upon heating, as well as alkaline earth metal but not alkali metal (except LiNO3, which is an anomaly).

Quote: Originally posted by chief  
The nitrates of 2-valent Metals produce NOx upon thermal de-composition, those of 1-valent metals don't ...
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[*] posted on 29-4-2010 at 06:28


I managed to find calcium nitrate. It's expensive though.. $45 for 50#. Geez!

How much HNO3 could I make from that? If I were to buy it, here's what I would probably try...

I would place the nitrate in a metal container. Then I'd make a hole in the lid for the Al tubing. The tubing would lead the gases into a 5 Gallon PP bucket filled with 2 gallons of cold water. I would put a lid on this bucket with a tiny hole in it so the NO2 doesn't readily escape but doesn't build up too much pressure either. And then start heating!

That's what I would do, unless someone has a better idea? But either way I'd have to think about it a little more cuz for $45... I dunno.. that's a little pricey!!
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[*] posted on 29-4-2010 at 06:59


If you can get sulfuric acid, mix fairly dilute H2SO4 and a solution of calcium nitrate, separate the precipitate of hydrated calcium sulfate, concentrate the solution somewhat, then distill. When the original dilute HNO3 is 10% to 15% in strength it's not too bad for filtering or decanting, in terms of the bulk of the CaSO4 - which will be the hydrate CaSO4 2H2O Wash the ppt, use the wash liquor to make the next batch of calcium nitrate solution; use the CaSO4 to make desiccant by heating it to 150-180 C.

Going through the decomposition to NOx and forming HNO3 from that is a lot of work and somewhat wasteful as the nitrate has other decomposition paths besides giving NO2.

BTW, that calcium nitrate may well be a hydrate, so there's both less 'NO3' than you'd think, and heating it will result in the formation of a solution or mush of the nitrate and steam coming off before the nitrate decomposes.

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[*] posted on 29-4-2010 at 07:19
Black powder - Potassium nitrate/sulphur


Quote: Originally posted by arsen  

It also might cause thoughts of the reactins during the explsion of black-powder: The S-vapor might partially react with the NO2 of the nitrates to some H2SO4, which might play some role ... ... ?


Extracted from :—

THE INITIATION, BURNING AND THERMAL DECOMPOSITION OF
GUNPOWDER
By J. D. BLACKWOOD AND F. P. BOWDEN, F.R.S.
Research Laboratory for the Physics and Chemistry of Surfaces,
Department of Physical Chemistry, University of Cambridge
Proceedings of the Royal Society (London)
Vol. 213. A. (8 July 1952) 285 1 19

(Received 3 November 1951-Revised 7 January 1952)

[Plates 3 to 6] {I do not have these plates. /djh/}

PART III. THERMAL DECOMPOSITION

It is well known that mixtures of potassium nitrate, charcoal and sulphur will react
exothermically with the production of a large volume of gas. These reactions
have formed the basis of extensive studies; one of the earliest and most
thorough is that of Noble & Abel. These workers examined the products but did
not formulate a reaction to explain the steps in the process. The more recent
contributions have been referred to earlier in the paper.

It appears that the presence of sulphur promotes reaction, and this has been
attributed by Hoffman (1929) to the formation of hydrogen sulphide at about
150oC from sulphur and organic matter present in the charcoal. This hydrogen
sulphide, he suggests, reacts exothermically with potassium nitrate above 280oC
to form potassium sulphate. It. is clear from the figures of Noble & Abel that the
amount of potassium sulphate formed is a function of the oxygen content of the
charcoal and not of the hydrogen content which only varies between 2 and 4 %.
It is also clear from experiment that the most reactive gunpowders are those
made from charcoal, the carbon content of which falls within quite narrow limits.
We shall see that reaction rate appears to be related to the organic materials
present in the charcoal as well as to the presence of sulphur, and the suggestion
put forward by Hoffman may require modification. Very little attention appears to
have been given to the preliminary reactions which appear to control the
behaviour of the gunpowder, and a study of these will be described in this
section.

The system potassium nitrate + sulphur

Experiments were carried out on mixtures of these materials. When they are
mixed together by grinding to a fine state, no reaction occurs. If they are heated,
gas evolution begins at a temperature of 250o C and is slow but continuous. As
the temperature is raised, the gas evolution is increased but is still of the same
form. Some typical results are shown in figure 15. The gases evolved during
heating are oxides of nitrogen, first nitric oxide and then nitrogen dioxide. If the
temperature is raised above the melting-point of potassium nitrate (334o C
sulphur dioxide can be detected quite readily although it cannot be detected at
lower temperatures. A small amount of potassium sulphate and potassium nitrite
is formed in the solid residue.

Attachment: Black Powder Blackwood & Bowden.pdf (327kB)
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[*] posted on 29-4-2010 at 07:37


Ok, that's it! I will no longer waste time trying to find other ways to make HNO3. I'll just accept the fact that distilling is simply the best way to go!

Ah, I feel kinda relieved actually :P
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[*] posted on 29-4-2010 at 09:07


Quote: Originally posted by arsen  
Heating NH4NO3 has to be done very carefully, as the salt is prone to unstable decomposition, esp. if contaminated; thus synthesis via NH4NO3 is not recommended pathway.


Question would be: _How_ manageable is it ... ?
==> I didn't find the old patent yet ...

... but I did it: It was a setup involving a sand-bath .... in a remote room ... where noone was ... : I let it run for a time, then disconnected the power-source and looked at it later ...
If I remember right I even watched the temperature ... electronically ...

The NH4NO3 should clearly not be directly heated by a flame or whatever ...
==> Also the NH4NO3 I used was lab-grade ...
===================

I believe someone _here_ pointed me to the patent, in a similar thread ...
==> It was quite detailed about the temperature-control etc. ; still have it on my old laptop, will post it sometime ...

It was a german/austrian patent from 1911(Austria) and 1913 (Germany), shortly before the HNO3-synthesis from coal was invented ...
==> It also discussed the advantages over the then-standard NaNO3-way ... ...

====================

But one of my favourites, which I never tried, would be the use of either Ca(NO3)2 or Ba(NO3)2 with H2SO4:
==> Both form insoluble sulfates with H2SO4, so the Ba/Ca just ppt. as heavy=soluble sulfates, and the HNO3 can be just filtered or distilled ...
==> Question only would be how well it works with water-free H2SO4 to make dry HNO3: Will the HNO3 the come easily out of the sulfate-crystallizate ?

[Edited on 29-4-2010 by chief]

[Edited on 29-4-2010 by chief]
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[*] posted on 29-4-2010 at 11:22


I learning chemistry in unversity and my university is a one man who working with glass he made for me retort.
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[*] posted on 2-5-2010 at 09:25


Here you can see my retort, its made from 1liter flash and glass pipe. I use it only for nitric acid.















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