Flower_18
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Potassium nitrate and potassium nitrite...
Hello,
I have a difficulty which I hope that some person could help me to clearify:
Potassium nitrate is decomposed in heat according to the formula KNO3 ---> KNO2 + O2. This reduction is said to take place at temperatures above
560 C. Ok. But now I read:
http://en.wikipedia.org/wiki/Potassium_nitrite
... that Potassium nitrite (in oxygen) in it's turn transformed into potassium nitrate at temperatures above 550 C. So would what happens if we heat
for example a portion of potassium nitrate in a small crucible constantly at let 's say 600 C. Would one end up with a mass which constantly absorbs
and rejects oxygen??? Because first the potassium nitrate would loose oxygen and become potassium nitrite but the latter in it's turn would absorb
oxygen and become potassium nitrate once again... I mean would these two phases co-exist? I don't get it. Hopefully some smart person in here could
help me to understand. Thank's a lot in advance!
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jock88
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Indeed, indeed. The old Moltox process.
http://www.ntis.gov/search/product.aspx?ABBR=DE83012847
Attachment: ESL-IE-86-06-78.pdf (1.1MB)
This file has been downloaded 671 times
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Zyklon-A
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There is some confusion on the subject of nitrates decomposing. Without a reducing agent, KNO3 doesn't let go of it's oxygen atom under
most conditions...
There is a equilibrium, that changes inconsistently at varying temperatures it would appear.
Also if you heat it up to much, it can give off nitrogen, nitrogen oxides, oxygen, or almost any mixture of those, according to one or more of the
following reactions:
2KNO3 ↔ 2KNO2 + O2.
4KNO3 ↔ 2K2O + 4NO + 3O2
4KNO3 ↔ 2K2O + 2N2 + 5O2.
[Note that most (all?) reactions are reversible.]
[EDIT] Read this: http://www.sciencemadness.org/talk/viewthread.php?tid=52#pid... Tons of useful information in there.
Also, Wikipedia says that KNO2 explodes at 537°C, so I don't know how that makes any sense. Also it apparently decomposes at 440°C.
[Edited on 26-2-2014 by Zyklonb]
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Flower_18
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Awesome man! Thank's a lot!  I have to dig into the moltox process...
[Edited on 26-2-2014 by Flower_18]
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Flower_18
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Quote: Originally posted by Zyklonb  | There is some confusion on the subject of nitrates decomposing. Without a reducing agent, KNO3 doesn't let go of it's oxygen atom under
most conditions...
There is a equilibrium, that changes inconsistently at varying temperatures it would appear.
Also if you heat it up to much, it can give off nitrogen, nitrogen oxides, oxygen, or almost any mixture of those, according to one or more of the
following reactions:
2KNO3 ↔ 2KNO2 + O2.
4KNO3 ↔ 2K2O + 4NO + 3O2
4KNO3 ↔ 2K2O + 2N2 + 5O2.
[Note that most (all?) reactions are reversible.]
The reactions giving rise topotassium oxide and eventually superoxide (explosive?) is as far as I am concerned only valable at temperatures above 800
C.
[EDIT] Read this: http://www.sciencemadness.org/talk/viewthread.php?tid=52#pid... Tons of useful information in there.
Also, Wikipedia says that KNO2 explodes at 537°C, so I don't know how that makes any sense. Also it apparently decomposes at 440°.
Thank's a lot! I will have a look into the thread. I do not think that potassium nitrite "exploded" above 537 C.
By the way one easy way to verify all this would simply be to observe the rise and fall of volume. As oxygen is constantly bound and set free.
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Oscilllator
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The following is a post of mine from the thread zyklonB linked that is particularly relevant to this thread:
| Quote: Originally posted by Oscilllator | A couple of my own observations regarding potassium nitrate -> nitrite through thermal decomposition:
Over several batches, approx 250g KNO3 was boiled in a stainless steel saucepan for about 15 minutes. The source of the heat was a coke-fired forge
that had previously proved itself capable of melting steel During the boiling, copious quantities of smoke was evolved.
When the molten salt was cooled down, it formed a distinctly green solid (greener than it looks in the photo). perhaps this is some kind of iron
compound from the stainless steel?
This was then dissolved in approx 500ml of water, heated by placing it on the forge. This brought the 500ml of water to the boil in about 10-15
seconds  . The resultant brown muddy liquid was filtered to obtain a clear
yellow solution that formed brown fumes upon addition of HCl. When the liquid was cooled down, large quantities of needle-shaped crystals formed that
proved to be potassium nitrate  .
This solution was then boiled down and a second crop of crystals were formed. They had a slightly different shape, and I am not certain that they are
KNO3 or KNO2, but I'm afraid the odds favour the KNO3. Irrespective of that, large quantities of KNO2 were formed in this process, as least enough
that the evolution of NO2 was so great that the test tube almost bubbled over.
It is interesting to note that wikipedia and a number of other sources say that KNO2 explodes when heated above around 700 degrees. I am 90% sure this
is not the case based on the fact that a rounded tablespoon of KNO3 was fully decomposed all the way to some black gunk that bubbled when water was
dropped on it. Presumably this black substance is K2O. At any rate, it was sitting on top of a red-hot piece of steel and didn't do anything that
could be described as an explosion .
I do have a hypothesis regarding the entire method of producing nitrites by thermal decomposition: Is it possible that as the nitrate decomposes the
formed nitrite also decomposes, so at no point in time is there a solution (if thats the right word) of 100% nitrate?
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I'm afraid you'll have to head over to the actual thread to see the pictures I took as they didn't survive the quotation process 
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Flower_18
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Oscillator wrote:
"I do have a hypothesis regarding the entire method of producing nitrites by thermal decomposition: Is it possible that as the nitrate decomposes the
formed nitrite also decomposes, so at no point in time is there a solution (if thats the right word) of 100% nitrate?"
It seem as you are 100% right! Heating potassium nitrate below 800 C would constantly form potassium nitrite which in it's turn is transformed into
potassium nitrate again, as the process is reversible. If you go over 800 C the process is more complex. Of course oxygen is needed in order to
transform the nitrite into nitrate again but heating potassium nitrate in a crucible or other contained outside (in the air) is perfectly enough as I
have understood.
So in practise: one would observe first bubbles (oxygen leaving potassium nitrate) and then followed by a decrease in volume. And after, the reverse.
Right...? I have to perform this. Thanx for the pictures also! I did go through the thread... ;o
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Zyklon-A
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Coincidentally, before you started this topic, I was experimenting with the reduction of KNO3 to KNO2. I used charcoal
(carbon) as a reducer, but I may try Al powder today. My results were basically similar to all the other people who tried it on the topic that I
linked....
I used 0.6 g carbon and 10 g KNO3, assuming that this reaction would occur: 2 KNO3 + C → 2 KNO2 +
CO2. However that wasn't able to burn by it self, so I added about .3 g more carbon, then it lit fine.
It burned very slowly, but hot enough to fuse the steel lid of my crucible to my crucible. After it cooled, I dropped a little HCl(aq)
on it, immediately, dense brown/orange NO2 was coming off. I also smelled chlorine, which I assume is do to the reduction of
HCl(aq).
[Edited on 26-2-2014 by Zyklonb]
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Flower_18
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I always thought that the reaction between C and KNO3 would give rise to:
K2CO3. Potash. So this reaction also gives us potassium nitrite (KNO2), very interesting indeed... Thanks for the comments.
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Zyklon-A
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Well, it can, in fact, it probably did make some potassium carbonate. For me it was only a test, I through it away after it was tested with
acid.
Gun powder makes K2CO3, and CO2.
[EDIT] Gun powder typically has a reaction similar to this:
10 KNO3 + 3 S + 8 C → 2 K2CO3 + 3 K2SO4 + 6 CO2 + 5 N2. It's hard
to give an exact equation, and there may be other reaction products as well.
[Edited on 26-2-2014 by Zyklonb]
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AJKOER
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In aqueous solutions, there are a couple of paths. In my opinion, reactions with Zn, Cd/Cu, and even Al foil (an amalgam also) are all reduction
reactions as a byproduct of the formation of an electrochemical cell (that is, a battery). The reaction proceeds in even dilute concentrations without
warming.
It is apparently well known and commonly employed by biochemist to use Zinc dust, but only sparingly as further reduction (to NH3, for example) can
ensue. The chemistry/biochemistry involves the formation of nitrate by bacteria, but perhaps in only very small doses. To test for the microbes, any
formed nitrate is first reduced to nitrite thereby permitting the use of a sensitive starch based test.
Source: See American Society for Microbe Biology library article "Nitrate and Nitrite Reduction Test Protocols" at http://www.microbelibrary.org/library/laboratory-test/3660-n... To quote from FIG. 5:
"Zinc dust will reduce nitrate to nitrite, but will not further reduce nitrite to nitrogen gas or other nitrogenous by-products when used sparingly."
The important word here is "sparingly", where an excess of Zinc will push the reduction further (see Aluminum discussion below).
-------------------------------------------------------------------------
Here is a commercial based method, using Cadmium (and Cadmium treated with Copper) that is complexed either by NH4Cl or the Sodium salt of EDTA. In a
neutral to basic solution the half cell reactions are given by:
[NO3]- + H2O + 2 e- ---> [NO2]- + 2 [OH]-
and for the oxidation of Cadmium by:
Cd + 1/2 O2 + H2O --> Cd(OH)2
or, if using EDTA:
[NO3]- + Cd + [EDTA]4- + H2O ---> [NO2]- + Cd[EDTA]2- + 2 [OH]-
The method is claimed to have near quantitative reduction of nitrate to nitrite.
Source: Journal of the Marine Biological Association of the United Kingdom / Volume 47 / Issue 01 / February 1967, pp 23-31. Link to full text: http://www.google.com/url?sa=t&rct=j&q=determination...
------------------------------------------------------------------------
Now, with respect to Aluminum (Al foil, more precisely) for those wishing to work with more friendly metals, one may be able to convert nitrates into
nitrites and beyond. Per this source (see equation [1] under Section 6.4, link:
https://docs.google.com/viewer?a=v&q=cache:Jz8OCxPNXSoJ:...), to quote:
"1. 3NO3- + 2Al + 3H2O → 3NO-2 + 2Al(OH)3
2. NO-2 + 2Al + 3H2O → NH3 + 2Al(OH)3 + OH-
3. 2NO-2 + 2Al + 4H2O → N2 + 2Al(OH)3 +2OH-
Nitrate reduction was found to be pH dependant. At pH values less than eight
no nitrate reduction took place. Above pH 10.5 nitrate was reduced upon addition of
the aluminium powder. Aluminium powder has been suggested for the denitrification
of sodium-based nuclear wastes, employing the nitrate to ammonia and ceramic
(NAC) process (Mattus et al., 1993, 1994)."
My personal experience with using Al foil on KNO3 in the presence of Na2CO3 (to address pH requirements) is that it apparently slowly works (at least
a day). However, make sure there is an excess of KNO3 at all times. I happen to remove some of the solution leaving some unreacted Aluminum. By the
next day, I was greeted with a massive gas buildup. When I attempted to vent, I was greeted with a rush of ammonia gas. No doubt a reduction of the
KNO3 had occurred all the way to NH3 and beyond to N2 as per Equation [2] and [3] above.
--------------------------------------------------------------------
For those who feel comfortable working with Pb, you may wish to try this path suggested in an old journal, "Journal of the Society of Chemical
Industry", Volume 27, based on commercial processes with some very insightful commentary on actual processes (see pages 484 to 485, link: http://books.google.com/books?pg=PA484&lpg=PA484&dq=... ). For example, an interesting mention of an aqueous method. To quote:
"The heating of a concentrated solution of lead nitrate with finely divided lead leads to the formation of insoluble basic lead nitrite, which
furnishes sodium nitrite by double decomposition with sodium carbonate. This reaction is, however, only of theoretical interest. "
For a source of the "finely divided lead" try the thermal decomposition of Lead oxalate, said to proceed as follows:
3 PbC2O4 --Heat--> 2 PbO + Pb + 4 CO2 + 2 CO
You maybe able to improve Pb yield by heating the Lead oxalate in a vertical tube reducing O2 exposure and venting the exit gas CO over hot PbO.
Also, for those who believe heating KNO3 is a good path to KNO2:
"Thermal decomposition.—The alkali nitrates when heated above their fusion point evolve oxygen and furnish nitrites, but this reaction is, of no
practical importance owing to the simultaneous occurrence of a further decomposition to oxide. "
[Edited on 27-2-2014 by AJKOER]
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Zyklon-A
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Very interesting that you bring this up.... I did this yesterday as well. I used sodium hydroxide to adjust the pH, the first time I used a pH of
~14. It reduced it to ammonia easily, and about 3 grams of Al was consumed within 10 minutes. One thing I noticed, was that the reaction did not occur
at room temperature, only with heating did it produce NH3.
Another strange phenomena, was that if you add the Al, at STP, and wait about 5 minutes, it would become unreactive even if you raise the temp.
Anyway, I plan on trying with Na2CO3 today instead of hydroxide. I will let it react overnight, and then post results.
[EDIT] Typo
[Edited on 26-2-2014 by Zyklonb]
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Praxichys
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I have a lot of food grade sodium nitrite, and I got it for cheap. I can bottle some up and sell it to you if you want. I'd go as cheap as $5 for half
a kilo, plus shipping.
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quantumcorespacealchemyst
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KNO3 + Ir + heat ----> Iridium Oxide and Potassium?
Do you think putting KNO3 with Iridium and heating it will be able to go to K2O and react with the Iridium?
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