AJKOER
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Inexpensive Route to Nitrates
I have been thinking about this synthesis. Some feedback please.
1. Create ClNO. Per Wikipedia (http://en.wikipedia.org/wiki/NOCl ):
"Another method of producing nitrosyl chloride is by direct union of the elements at 400 °C:
N2 + O2 + Cl2 → 2 ClNO "
I was thinking of running dilute Cl2 (from a chlorine generator or heating previously created Chlorine hydrate) mixed with air (but more Cl2 than
O2), and perhaps a little water vapor, through copper tubing ([EDIT] Iron vessel) heated over a charcoal fire.
2. Deposit the gaseous output into cold water. Reactions:
HNO3 (aq) + 3 HCl (aq) <---> NOCl (g) + Cl2 (g) + 2 H2O (l)
Per this reference (http://site.iugaza.edu.ps/bqeshta/files/2010/02/94398_16.pdf ), some possible reactions to avoid like excess HNO3 as:
2ClNO + 4HNO3 → 6NO2 + Cl2 + 2H2O
Also, avoid neutralizing a solution and further treating with any unreacted chlorine as:
NaNO3 + Cl2 → ClNO + NaClO2
which in the presence of a hypochlorite (from, for example, the action of chlorine on a carbonate) and an acid (chlorine and water forms HCl + HOCl)
could form the explosive and toxic ClO2 (temperature dependent, see comments by Garage Chemist and BromicAcid at http://www.sciencemadness.org/talk/viewthread.php?tid=3314 on the aqueous reaction ).
At the cessation of gas generation, one can neutralize the final solution with say, NaHCO3 to form NaCl and NaNO3.
Be cautious with even dilute solutions and certainly do not attempt this synthesis without appropriate safety gear as ClNO gas is both extremely
corrosive and toxic (this is a valid criticism of this process).
Now, some possible side reactions:
Cu + Cl2 --> CuCl2
Cu + 2 ClNO --> CuCl2 + 2 NO
2 ClNO --Thermal Decomposition--> Cl2 + 2 NO
2 NO + O2 --> 2 NO2
2 NO2 + H2O <--> HNO2 + HNO3
Cl2 + H2O <--> HCl + HOCl
ClNO + H2O --> HNO2 + HCl
HNO2 + HOCl --> HCl + HNO3
3 HNO2 (aq) ---Hot or Conc ---> HNO3 + 2 NO + H2O
4 HNO3 + Cu → Cu(NO3)2 + 2 NO2 +2 H2O
Cu + 2 N2O4 → Cu(NO3)2 + 2 NO
2 Cu(NO3)2 --Thermal decomposition--> 2 CuO + 4 NO2 + O2
Cu + CuCl2 → 2 CuCl
6 CuCl + 3/2 O2 + 3 H2O → 2 Cu3Cl2(OH)4 + CuCl2
and others,...., but I don't believe these reactions will significantly interfere with the formation of HNO3 and nitrates there from.
Optionally, one can separate out the chloride by adding Lead or Silver acetate (whose preparations have been addressed on this forum), ([EDIT] or
extract NO using Copper metal, add O2 and water to form HNO3).
[Edited on 5-2-2013 by AJKOER]
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Adas
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I have no idea what you are trying to achieve O.o
Rest In Pieces!
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barley81
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That would be pretty expensive if it works at all. If you really want to get nitric acid from nearly nothing, it would be cheaper to use the
Birkeland-Eyde process. Also, where did the equation Cl2 + NaNO3 -> NaClO2 + ClNO come from? AFAIK that doesn't happen.
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AJKOER
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Quote: Originally posted by barley81  | | That would be pretty expensive if it works at all. If you really want to get nitric acid from nearly nothing, it would be cheaper to use the
Birkeland-Eyde process. Also, where did the equation Cl2 + NaNO3 -> NaClO2 + ClNO come from? AFAIK that doesn't happen. |
Thanks for the comment.
I have since added a link with, perhaps, a personal observation by Garage Chemist on the creation of ClO2, and comments by BromicAcid with cited
references.
I mention this apparently not too well known reaction both for its obscurity and that, in a rationale attempt at safety by immediately passing the
NOCl and Cl2 mixture into a baking soda solution (not recommended) one may, indeed, develop a false sense of safety.
[Edited on 4-2-2013 by AJKOER]
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blogfast25
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Generally speaking no one does. Move on. He fancies himself as a bit of an inventor, about to crack one of the Holy Grails of chemistry. Personally
I'd like to see him crack his first test tube 
[Edited on 4-2-2013 by blogfast25]
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AJKOER
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| Quote: Originally posted by barley81 | | That would be pretty expensive if it works at all. If you really want to get nitric acid from nearly nothing, it would be cheaper to use the
Birkeland-Eyde process... |
Per Wikipedia (http://en.wikipedia.org/wiki/Birkeland-Eyde_process ) on the Birkeland-Eyde process:
"An electrical arc was formed between two coaxial electrodes, and through the use of a strong magnetic field, was spread out into a thin disc. The
plasma temperature in the disc was in excess of 3000°C. Air was blown through this arc, causing some of the nitrogen to react with oxygen forming
nitric oxide. By carefully controlling the energy of the arc and the velocity of the air stream, yields of up to 4% nitric oxide were obtained. The
process is extremely energy intensive. Birkeland used a nearby hydroelectric power station for the electricity as this process demanded about 15
MWh/Ton of nitric acid. The same reaction is carried out by lightning, providing a natural source for converting atmospheric nitrogen to soluble
nitrates."
My process appears to be less energy intensive and, I suspect, over 4% yield.
Now, the expense argument is a point, certainly not commercially competitive process. However, one may be able to inexpensively generate Chlorine (in
compliance with local laws) from home chemicals (there is at least one forum on this). Then, using just air, copper tubing, charcoal and baking soda,
not much of an added expense to nitrates.
I am assuming that one can also, with more work, inexpensively extract NO from a dilute product solution by reacting with Cu (from the tubing). Add
air to the NO gas to create NO2, dissolve in water and treat with more air for HNO3.
[Edited on 4-2-2013 by AJKOER]
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AJKOER
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Blogfast:
Thanks for your indirect compliment on the process.
If there was any issue, believe me Blogfast would let us all know (and not too politely either as you may have surmised).
Blogfast, as you seem to enjoy heating things up (thermites and the like), I thought you might find this synthesis cooler.
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hissingnoise
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| Quote: | | Thanks for your indirect compliment on the process. |
You appear to get everything ass backwards . . . ?
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AJKOER
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Hissingnoise:
You speak too soon, I know Blogfast25, he is just a little frustrated now ( ), as
he has not actually found anything wrong with the synthesis. If you don't believe me, ask him yourself. Silence, in his case, is certainly a tacit
approval.
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AJKOER
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On further research, an iron vessel in place of copper to heat Cl2, O2 and N2 together may be appropriate. This reference (http://nitrogen.atomistry.com/nitrosyl_chloride.html ) notes that Nitric oxide decomposes ferric chloride at a high temperature via the reaction
below to form nitrosyl chloride:
FeCl3 + NO = FeCl2 + NOCl
Also, as per Wikipedia, "Anhydrous iron(III) chloride may be prepared by union of the elements:[7]
2 Fe(s) + 3 Cl2(g) → 2 FeCl3(s) "
So, any NO present could be converted into nitrosyl chloride via this path in an Iron vessel.
In addition, per the same reference, the preparation of NOCl from passing hydrogen chloride passed into liquid nitrogen trioxide (which evolves not
quite pure nitrosyl chloride) requires dry HCl:
2HCl + N2O3 = 2NOCl + H2O
However, this research (see http://www.pnas.org/content/early/2009/07/17/0904195106.full... ) strongly suggests a novel and important role of water vapor in the atmospheric
formation of NOCl and NO2Cl. As such, the question of dry Cl2, N2 and O2 gases, or with a little moisture, appears open.
The last reference also notes the importance of surface kinetics and the study uses SiO2 powder, which could be added to the Iron reaction vessel.
This is interesting as old Chem books mention the use of crushed glass in gaseous reactions, and also animal charcoal or bone-black as a catalyst.
[Edited on 5-2-2013 by AJKOER]
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AndersHoveland
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Another improvised way to make small ammounts of nitrates, though it is impractical, is to let ammonia stand with H2O2 in the presence of a base (NaOH
or Na2CO3). The base helps catalyze the decomposition of, and oxidation by the H2O2. The reaction is fairly slow. Some sodium nitrate will form. Under
different conditions the oxidation of NH4OH by H2O2 can form some ammonium nitrite, which will decompose to nitrogen gas if the solution is heated.
Typically however the H2O2 tends to oxidize the nitrite to nitrate much faster than the ammonia is oxidized. Indeed, if ammonia is mixed with H2O2
there is no obvious immediate reaction.
The slow reaction rate is probably due to either (or perhaps both) of the following equilibrium or decomposition pathway:
2 NH4OH <==> 2 H2O + 2 NH3
2 NH3 <==> NH4+ + NH2-
H2O2 + OH- <==> HOO- + H2O
HOO- + H2O2 --> OH- + H2O + O2Δ
O2Δ + 2 H2O <==> H2O3 + H2O
(this last equilibrium seems to occur through a simultaneous intermolecular resonance)
The decomposition of peroxide under alkaline conditions is actually more complex. O2Δ represents a long-lived singlet excited state of
diatomic oxygen. Dihydrogen trioxide is unstable and quickly decomposes, supposedly leading to transient superoxide ions. H2O3 is a stronger oxidizer
than H2O2.
The equilibrium constant of amide ions, NH2-, in aqueous solution is very low, a potential explanation for the very slow reaction rate.
Amide would be more vulnerable to oxidation than NH3.
| Quote: Originally posted by Formatik |
Under what conditions can ammonia be oxidized to NH4NO2 ?
The Ber. ref. from Hoppe-Seyler describes it. Namely, strong solutions of H2O2 with a few drops of NH4OH or solutions of ammonium carbonate (with or
without NaOH or Na2CO3) can be let to stand 24 hours without any nitrite formation occurring. But upon longer standing, even with a small amount of
hydroxide then nitrite forms. Nitrite also forms when a dilute solution of H2O2 is mixed with NH4OH and a little Na2CO3 and is evaporated over pure
conc. H2SO4 with a bell jar.
H2O2 forms (even in very dilute solutions) nitrite very rapidly, if the H2O2 solution is mixed with a few drops of NH4OH and a little NaOH or Na2CO3,
and this then boiled in a retort to a very small volume. They suggest this nitrite formation as a demonstration experiment because it is very quick to
do, and then after acidification of the colorless liquid with H2SO4, the HNO2 can be nicely be proven to be present.
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| Quote: Originally posted by AndersHoveland |
A mixture of hydrogen peroxide and ammonium hydroxide (in a 1:3 ratio) acts as a reactive oxidizer, which can attack organic compounds and elemental
carbon. The reaction rate is negligible at room temperature, but when heated to 60°C the reaction becomes vigorous and self-sustaining. Such
solutions are sometimes known as "base piranha". With a 1:1:5 volume ratio of NH4OH, H2O2, and H2O, respectively, the half-life times of peroxide were
4 hours at 50°C and 40 minutes at 80°C.
"Reaction of Ozone and H2O2 in NH4OH Solutions and Their Reaction with Silicon Wafers" Japanese Journal Applied Physics. 43 (2004) pp.
3335-3339. |
[Edited on 5-2-2013 by AndersHoveland]
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AJKOER
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AndersHoveland:
Those making NH4NO2 (from NH4OH, H2O2 & Na2CO3) should be aware of its apparently toxic nature (Wikipedia describes it as "acutely toxic").
Certainly avoid getting it on your skin, as someone might accidentally do upon mixing a 'cleaning' solution (strongly not recommended) from Oxiclean
(Na2CO3.H2O2) with hot water and Ammmonia.
Wikipedia (http://en.wikipedia.org/wiki/NH4NO2 ) further notes:
"Ammonium nitrite solution is stable at higher pH and lower temperature. If there is any decrease in pH lower than 7.0, It may lead to explosion. It
is desirable to maintain pH by adding Ammonia solution. The mole ratio of Ammonium Nitrite to Ammonia must be above 10% mole ratio.
NH4NO2 → N2 + 2 H2O "
Now, the comment of aqueous unstability (a potential for explosion) upon change in pH is probably accurate as I do recall some comments on factory
explosions.
What I would like to know is what is produced when, say Aluminum or Zinc or Copper, is attacked by the so called "base piranha"? The comment on the
life span of the solution implies to me that one may indeed be working with NH4NO2, which has a characteristic two hour half-life, decomposing between
60 to 70 C. Caution: the dry salt, itself, is considered a high explosive with limited applications due to thermal and shock sensitivity.
Now, on my speculated reactions (which could be unsafe if performed at room temperature or at all):
5 H2O + NH4NO2 + 3 Zn → 2 NH3 + H2O + 3 Zn(OH)2
2 NH4NO2 + 6 H2O + 4 SO2 → 3 H2SO4 + (NH4)2SO4 + 2 NH2OH
Note: the product, Hydroxylamine, if formed with a low temperature synthesis, has noted explosive tendencies as well (my speculation, this may be why
NH4NO2 seemingly explodes on lowering its pH).
(see http://en.wikipedia.org/wiki/Nitrous_acid and http://en.wikipedia.org/wiki/NH2OH )
I think it would also be interesting (and even more dangerous) to add, in small amounts, to an iced NH4NO2 solution, HOCl. This parallels the
recommended synthesis of NH2Cl which could also be formed by action on free NH3. Speculated successes/dangers:
NH4NO2 + HOCl ---Cold Dilute Solution Only--?--> NH4Cl + HNO3
Since, as was previously noted above:
HNO2 + HOCl --> HCl + HNO3
and with free NH3:
NH3 + HNO3 --> NH4NO3
Possible other side reactions which occur also in small quantities in the cold under appropriate safety precautions:
NH3 + HOCl <--> NH2Cl + H2O
NH2Cl + NH4NO2 + H2O --?-> NH4Cl + NH2OH.HNO2
NH4NO3 + NH2Cl + H2O --?--> NH4Cl + NH2OH.HNO3
but, I will confess, these reactions are most likely above my pay grade/propensity to be energetic/safety comfort level (as Chloramine has been known
to extenuate the poisonous properties of the combining compounds). Some help perhaps per comments per Wikipedia:
"Hydroxylammonium nitrate or hydroxylamine nitrate (HAN) is an inorganic compound with the chemical formula NH3OHNO3. It is a salt derived from
hydroxylamine and nitric acid. It is related to ammonium nitrate but has a higher oxygen content. In its pure form, it is a colourless hygroscopic
solid. It is used as a rocket propellant."
Also: "The compound is a salt with separated hydroxyammonium and nitrate ions.[1] Hydroxylammonium nitrate is unstable because it contains both a
reducing agent (hydroxylammonium cation) and an oxidizer (nitrate),[2] the situation being analogous to ammonium nitrate. It is usually handled as an
aqueous solution. The solution is corrosive and toxic, and may be carcinogenic. Solid HAN is unstable, particularly in the presence of trace amounts
of metal salts."
Here is a government report on Hydroxylamine nitrate discussing decomposition issues (especially in the presence of Iron when concentrated) and safety
(link: http://www.hss.energy.gov/healthsafety/wshp/chem_safety/docs... )
-------------------------------------------------------------------------
Back on topic, one of the inspirations for the recommended synthesis was a patent claim on the apparent speed (a few hours) and yield by the action of
polarized uv on O2 and Cl2 (and a note on the necessity to remove N2 due to nitrogen cross products). So if polarized uv can be argued to be a
successful commercial source of ClO2 by direct union (and apparently NOCl as well as an unwanted product), there may be some value in the straight
thermal combination approach. Now the process of heating a stream of corrosive gases to 400 C may not be a pleasant/safe/convenience laboratory
preparation or produce sizable quantities, but may still be doable outdoors with constructed equipment and heating source.
[Edited on 6-2-2013 by AJKOER]
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AJKOER
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Easy But More Perilous/Problematic Path to Heating Cl2, O2 and N2
-------------------------------------------------------------------------------------------
Combine a given volume of H2 with equal quantity of Cl2, 2/3 quantity of O2 (or more) and 1/3 of N2. With precautions, expose the mixture to sunlight
or an electric spark. Major issues to be addressed, the reaction will be explosive/expansive. The H2 dilution will, however, reduce the temperature
from the maximum temperature of about 2800° C (achieved with a pure stoichiometric 2 X 1 mixture) by at least 700 degrees (as observed by a hydrogen
flame in air, see http://en.wikipedia.org/wiki/Oxyhydrogen ), but still above the 400° C required formation temperature for NOCl.
Alternately, one could start by burning a hydrogen flame in air and change the air to appropriate Cl2, O2 and N2 mixture. [EDIT] Pre-mixing of the
gases and flow rate could be important.
Expected Reactions:
2 H2 + O2 --> 2 H2O
H2 + Cl2 --> 2 HCl
N2 + O2 + Cl2 --> 2 NOCl
----------------------------------
The increase in pressure (if the explosion path chosen) should also move the last reaction to the right. On net, and adding more Cl2 (to both sides):
3 H2 + N2 + 2 O2 + 3 Cl2 --> 2 NOCl + 2 Cl2 + 2 HCl + 2 H2O
And, with 2 more H2O added after the combustion (or explosion), we have:
2 NOCl + 2 Cl2 + 4 H2O <---> 2 HNO3 + 6 HCl
as required to form Nitric acid for nitrates.
Conceiable related support for this mode of synthesis relates to the problematic formation of NOx with the combustion of nitrogen containing fuels for
example, per Wikipedia (http://en.wikipedia.org/wiki/NOx ) to quote:
"The major source of NOx production from nitrogen-bearing fuels such as certain coals and oil, is the conversion of fuel bound nitrogen to NOx during
combustion.[7] During combustion, the nitrogen bound in the fuel is released as a free radical and ultimately forms free N2, or NO. Fuel NOx can
contribute as much as 50% of total emissions when combusting oil and as much as 80% when combusting coal."
So, assuming the temperature is sufficient to form nitrogen radicals, one could expect NOCl, as well as even some NO, NO2 and NO2Cl formation.
The major issue with this mode of synthesis is designing the experiment safely, and having the burning/detonation chamber releasing gases into a large
vessel for collection. If a detonation chamber is used, it could be designed as either expendable or not. The practicality of this combustion
synthesis remains to be proven, as does the direct thermal composition approach.
[Edited on 13-2-2013 by AJKOER]
[Edited on 13-2-2013 by AJKOER]
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Agecer
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| Quote: Originally posted by AJKOER | I have been thinking about this synthesis. Some feedback please.
1. Create ClNO. Per Wikipedia (http://en.wikipedia.org/wiki/NOCl ):
"Another method of producing nitrosyl chloride is by direct union of the elements at 400 °C:
N2 + O2 + Cl2 → 2 ClNO "
I was thinking of running dilute Cl2 (from a chlorine generator or heating previously created Chlorine hydrate) mixed with air (but more Cl2 than
O2), and perhaps a little water vapor, through copper tubing ([EDIT] Iron vessel) heated over a charcoal fire.
2. Deposit the gaseous output into cold water. Reactions:
HNO3 (aq) + 3 HCl (aq) <---> NOCl (g) + Cl2 (g) + 2 H2O (l)
Per this reference (http://site.iugaza.edu.ps/bqeshta/files/2010/02/94398_16.pdf ), some possible reactions to car dvd players avoid like excess HNO3 as:
2ClNO + 4HNO3 → 6NO2 + Cl2 + 2H2O
Also, avoid neutralizing a solution and further treating with any unreacted chlorine as:
NaNO3 + Cl2 → ClNO + NaClO2
which in the presence of a hypochlorite (from, for example, the action of chlorine on a carbonate) and an acid (chlorine and water forms HCl + HOCl)
could form the explosive and toxic ClO2 (temperature dependent, see comments by Garage Chemist and BromicAcid at http://www.sciencemadness.org/talk/viewthread.php?tid=3314 on the aqueous reaction ).
At the cessation of gas generation, one can neutralize the final solution with say, NaHCO3 to form NaCl and NaNO3.
Be cautious with even dilute solutions and certainly do not attempt this synthesis without appropriate safety gear as ClNO gas is both extremely
corrosive and toxic (this is a valid criticism of this process).
Now, some possible side reactions:
Cu + Cl2 --> CuCl2
Cu + 2 ClNO --> CuCl2 + 2 NO
2 ClNO --Thermal Decomposition--> Cl2 + 2 NO
2 NO + O2 --> 2 NO2
2 NO2 + H2O <--> HNO2 + HNO3
Cl2 + H2O <--> HCl + HOCl
ClNO + H2O --> HNO2 + HCl
HNO2 + HOCl --> HCl + HNO3
3 HNO2 (aq) ---Hot or Conc ---> HNO3 + 2 NO + H2O
4 HNO3 + Cu → Cu(NO3)2 + 2 NO2 +2 H2O
Cu + 2 N2O4 → Cu(NO3)2 + 2 NO
2 Cu(NO3)2 --Thermal decomposition--> 2 CuO + 4 NO2 + O2
Cu + CuCl2 → 2 CuCl
6 CuCl + 3/2 O2 + 3 H2O → 2 Cu3Cl2(OH)4 + CuCl2
and others,...., but I don't believe these reactions will significantly interfere with the formation of HNO3 and nitrates there from.
Optionally, one can separate out the chloride by adding Lead or Silver acetate (whose preparations have been addressed on this forum), ([EDIT] or
extract NO using Copper metal, add O2 and water to form HNO3).
[Edited on 5-2-2013 by AJKOER] |
It seems to be long process but if it get desired results than I have no hesitation in trying it
[Edited on 13-2-2013 by Agecer]
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AJKOER
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Agecer:
I think you efforts would be noteworthy.
I am continuing to think over various ways to create a safe, effective low temperature, relatively inexpensive and convenient synthesis (which appears
problematic) to nitrates.
Small scale my be the best 1st approach.
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ScienceSquirrel
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| Quote: Originally posted by AJKOER | Easy But More Perilous/Problematic Path to Heating Cl2, O2 and N2
-------------------------------------------------------------------------------------------
Combine a given volume of H2 with equal quantity of Cl2, 2/3 quantity of O2 (or more) and 1/3 of N2. With precautions, expose the mixture to sunlight
or an electric spark. Major issues to be addressed, the reaction will be explosive/expansive. The H2 dilution will, however, reduce the temperature
from the maximum temperature of about 2800° C (achieved with a pure stoichiometric 2 X 1 mixture) by at least 700 degrees (as observed by a hydrogen
flame in air, see http://en.wikipedia.org/wiki/Oxyhydrogen ), but still above the 400° C required formation temperature for NOCl.
Alternately, one could start by burning a hydrogen flame in air and change the air to appropriate Cl2, O2 and N2 mixture. [EDIT] Pre-mixing of the
gases and flow rate could be important.
Expected Reactions:
2 H2 + O2 --> 2 H2O
H2 + Cl2 --> 2 HCl
N2 + O2 + Cl2 --> 2 NOCl
----------------------------------
The increase in pressure (if the explosion path chosen) should also move the last reaction to the right. On net, and adding more Cl2 (to both sides):
3 H2 + N2 + 2 O2 + 3 Cl2 --> 2 NOCl + 2 Cl2 + 2 HCl + 2 H2O
And, with 2 more H2O added after the combustion (or explosion), we have:
2 NOCl + 2 Cl2 + 4 H2O <---> 2 HNO3 + 6 HCl
as required to form Nitric acid for nitrates.
Conceiable related support for this mode of synthesis relates to the problematic formation of NOx with the combustion of nitrogen containing fuels for
example, per Wikipedia (http://en.wikipedia.org/wiki/NOx ) to quote:
"The major source of NOx production from nitrogen-bearing fuels such as certain coals and oil, is the conversion of fuel bound nitrogen to NOx during
combustion.[7] During combustion, the nitrogen bound in the fuel is released as a free radical and ultimately forms free N2, or NO. Fuel NOx can
contribute as much as 50% of total emissions when combusting oil and as much as 80% when combusting coal."
So, assuming the temperature is sufficient to form nitrogen radicals, one could expect NOCl, as well as even some NO, NO2 and NO2Cl formation.
The major issue with this mode of synthesis is designing the experiment safely, and having the burning/detonation chamber releasing gases into a large
vessel for collection. If a detonation chamber is used, it could be designed as either expendable or not. The practicality of this combustion
synthesis remains to be proven, as does the direct thermal composition approach.
[Edited on 13-2-2013 by AJKOER]
[Edited on 13-2-2013 by AJKOER] |
This is not going to work as a practical synthesis.
It is possible that traces of nitrosyl chloride or nitrogen oxides may form depending on the temperature and pressure but trace is the word.
The main difficulty is the lack of reactivity of molecular nitrogen.
Natural nitrogen fixation occurs in lightning, certain bacteria. There are also several industrial processes.
One of the things that characterises them all is that they are all energy intensive.
The field has seen almost two centuries of research and there have been no major advances for over a century with the Haber Ostwald process dominating
the production of nitric acid completely.
The field has not been neglected either as a better route would have major implications for securing food supplies and would make the discoverers
immensely rich.
http://en.wikipedia.org/wiki/Nitrogen_fixation
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AJKOER
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ScienceSquirrel:
I see your point, but this synthesis is in concordance with a patent (see http://www.freshpatents.com/-dt20100204ptan20100025226.php ) that employs excited chlorine and oxygen produced by only polarized ultraviolet
radiation to create ClO2. With respect to yield, to quote: "Expected results are obtained until the chlorine dioxide concentration approaches about
8%" at which point an instability in the kinetics appears to occur.
The author also notes, which is directly applicable to NOCl formation, to quote: "In an embodiment, the reaction to produce chlorine dioxide is
carried out in a reaction space devoid of nitrogen. The presence of nitrogen does not prevent the formation of the chlorine dioxide, but nitrogenous
chlorine-containing compounds are potentially formed as by-products. This lowers the yield of chlorine dioxide and is, of course, undesired."
So the question is can direct thermal means produce excited oxygen and nitrogen as well in an acceptable yield? The literature cited 400 C for the
reaction:
N2 + O2 + Cl2---> 2 NOCl
So does a yield in the neigborhood of 8% also seems plausible? I suspect perhaps, yes. Clearly, the need for Cl2 and low yield do not makes this a
commerically competitive process, but it does benefit from available inputs. Also, given the nature of the products produced, low yield has safety
advantages.
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AndersHoveland
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| Quote: Originally posted by AJKOER |
Cu + 2 ClNO --> CuCl2 + 2 NO
Cu + 2 N2O4 → Cu(NO3)2 + 2 NO
ClNO + H2O --> HNO2 + HCl
HNO2 + HOCl --> HCl + HNO3
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Your expected reactions will indeed occur, but only in aqeuous solution.
Nitrosyl chloride can react with copper to form an addition complex, which will decompose in the presence of water. Anhydrous coppric nitrate is
actually a reactive oxidizer/nitrating agent in the absence of water. Apparently nitrosyl chloride does not completely hydrolyze in concentrated
solutions of hydrochloric acid. In concentrated acidic solution, hydrochloric acid can reduce nitric acid (the equilibrium being shifted by the pH).
Actually, this reaction does not readily proceed in the absence of addition chloride ions. The equation would be more accurately written as,
Cu + CuCl2 + 2Cl- → 2CuCl2-
A solution of copper(II) nitrate is unable dissolve a sliver of copper foil (I have tried this, letting it stand for two days, even adding some very
dilute nitric acid did not help). However, I did find a source that stated that "cupric nitrate solution reacts with finely divided copper to form
nitric oxide and basic copper nitrate."
Cupric nitrate has the formula Cu(NO3)2. Basic copper nitrate has the formula Cu(NO3)2∙3Cu(OH)2. Several other transition metals can form
similar insoluble basic salts.
[Edited on 14-2-2013 by AndersHoveland]
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Wilhelm Scream
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My calculations show that the production of NOCl by direct union of the elements isn't thermodynamically possible (HSC Gibbs calculations).
Wikipedia isn't always a trustworthy source. Also, even if it did work, this experiment would be one of the most dangerous things I have ever seen.
One small, single leak and suddenly your home is full of substantial quantities of hot chlorine and NOCl. I commend your originality, but it just
isn't worth it.
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AJKOER
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OK, some clarifications.
1. Another source other than Wkikpedia (see http://site.iugaza.edu.ps/bqeshta/files/2010/02/94398_16.pdf ) to quote:
"Also, nitrosyl chloride can be synthesized from its elements by heating nitrogen, oxygen and chlorine gas at 400°C:
N2 + O2 + Cl2 → 2ClNO "
However, this reaction, in my opinion and the author of the patent I cited on the direct union of Chlorine and Oxygen, is that Ozone is a likely
intermediary, leading to an active Oxygen radical and the formation of NO (and then NOCl) and some ClO2. Also, active Chlorine (along with CH and OH
radicals) is a suspected catalyst to Ozone formation. I believe it is also fair to state that the precise chemistry is not entirely understood at this
time.
2. Please note that I revised the suggested employment of Cu with Fe as per this reaction (see http://nitrogen.atomistry.com/nitrosyl_chloride.html ):
FeCl3 + NO = FeCl2 + NOCl
any Iron chloride formed would foster the formation of NOCl. Now, I do not have a reference with Copper behaving similarly. Also, both CuCl and FeCl3
form additive compounds with NOCl that decompose at high temperatures.
In a further embodiment to avoid a direct heating approach, I suggested a common practice of adding a fuel gas to effect a reaction (in this case H2).
I would further suggest the ignition via a light source as this could introduce Cl radicals, which are possibly beneficial to reaction pathways.
3. As to my listing of gases and aqueous reactions, I apology for the confusion. The original propsed synthesis was intended to be a direct gaseous
union (but not dry gases) followed by dissolving the NOCl and excess Cl2 in water to form a dilute Aqua Regia. The reason for including water vapor
was, as I referenced above, some proposed 'novel' chemistry of reaction pathways so there will be confusion nevertheless.
4. Gibbs energy calculations are suggestive as to a reaction success, but not definitive. This would be true here where free readical formations are a
suspected catalyst to varying reaction pathways.
5. While my speculation is a low yield of NOCl, please note that the direct thermal, or the fuel gas/photo ignition approach may produce still a
significant yield given the 80% formation of NOx in the burning of coal where the bounded nitrogen is activated at high temperatures.
Now, despite all of the research and feedback, the proposed synthesis has continued to evolved (as it comes to grasp with the chemistry) to facilate
the possibility of both higher yields and larger scale production. Else, one must contend with the scale and yield implied by the polarized uv
radiation synthesis.
[Edited on 14-2-2013 by AJKOER]
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AJKOER
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AndersHoveland:
With respect the oxidation of ammonia to nitrate, using ozone is another path. Here is a reference that notes adding NaBr increases the rate at which
O3 attacks NH3.
4 O3 + NH3 --Bromide--> HNO3 + 4 O2 + H2O
Link: http://www.eolss.net/Sample-Chapters/C07/E6-192-06-00.pdf
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