BlindedAchievement
Harmless
Posts: 37
Registered: 18-1-2009
Member Is Offline
Mood: Mutually Happy
|
|
What can oxidize ammonia?
I looked in my avaivable books and search engines to find anything that can oxidize ammonia to nitrogen oxide. All they gav me was something about an
astrobiological bacteria that can oxidize ammonia. What I need to know is if there is any agent that would work best, and how to use it. The chemicals
I have avaivable are listed in the Uses of chlorine thread in the beginnings section. I have some thought for potassium permanganate.
|
|
|
Xenoid
National Hazard
Posts: 775
Registered: 14-6-2007
Location: Springs Junction, New Zealand
Member Is Offline
Mood: Comfortably Numb
|
|
Ammonia is oxidised by oxygen in air to nitric oxide and water, using Pt gauze heated to 700 oC. This is the basis for the industrial production of
nitric acid.
4NH3 + 5O2 --> 4NO + 6H2O
The nitric oxide is allowed to cool and react with oxygen to produce nitrogen dioxide, this is reacted with water in towers to produce nitric acid.
Let me guess, you don't have Pt gauze in your list of chemicals!
As far as I am aware, most common oxidising agents tend to convert NH3 to nitrogen and water.
for example:
hot copper oxide; 3CuO + 2NH3 --> 3Cu + N2 + 3H2O
potassium permanganate; 2KMnO4 + 2NH3 --> 2KOH + 2MnO2 + 2H2O + N2
|
|
|
Mr. Wizard
International Hazard
Posts: 1042
Registered: 30-3-2003
Member Is Offline
Mood: No Mood
|
|
Ozone generators have the ability to remove ammonia from the air. http://en.wikipedia.org/wiki/Ozone
says it forms Ammonium Nitrate in the process.
|
|
|
hissingnoise
International Hazard
Posts: 3940
Registered: 26-12-2002
Member Is Offline
Mood: Pulverulescent!
|
|
I don't know how authoritative the wiki article is---I'd imagine ammonium ozonide would be the product there.
It's an oxidiser that could be used in propellants.
((ooops!) On checking it looks like I got it wrong.)
[Edited on 6-2-2009 by hissingnoise]
|
|
|
Mr. Wizard
International Hazard
Posts: 1042
Registered: 30-3-2003
Member Is Offline
Mood: No Mood
|
|
| Quote: | Originally posted by hissingnoise
I don't know how authoritative the wiki article is---I'd imagine ammonium ozonide would be the product there.
It's an oxidiser that could be used in propellants.
((ooops!) On checking it looks like I got it wrong.)
[Edited on 6-2-2009 by hissingnoise] |
Nobody ever admits they were wrong :-) I suspect you are actually right. Maybe it's an intermediate on the way to Ammonium Nitrate. Since it takes
two NH3 molecules to produce the required number of N atoms in NH4NO3 and 4 Ozone molecules, it's unlikely it's made by a single step. No doubt the
Wiki article is simplified.
|
|
|
Formatik
National Hazard
Posts: 927
Registered: 25-3-2008
Member Is Offline
Mood: equilibrium
|
|
NH3 oxidation
It is no doubt more complex than as seen on wiki. And I would take some of those reactions as the end result and being the products under very
specific conditions. NH3 gets oxidized to HNO2, then if enough O3 is present, further oxidized to HNO3; with an excess NH3 NH4NO2 and NH4NO3 form
(R.v. Helmholtz, F. Richarz, Wied. Ann. 40 [1890] 161/202, 167). With the reaction dry gases, then basically the only resulting product is NH4NO3,
even with a large NH3 excess, nitrite occurs only in traces, formation of NH2OH and H2O2 is never observed, so that the reaction is taken to be: 2NH3
+ 4 O3= NH4NO3 + H2O + 4 O2 (E. Dieckhoff, Habilitationsschr. Karlsruhe T.H. 1891, p. 59). If NH3 present in air in small amounts (25 mg/m3 of air),
then O3 present in small amounts (up to 30mg/m3 air) at regular temperature has no action on NH3, even though the same present amount of O3 would be
enough to cause complete oxidation of the NH3 (A. Erlandsen, L. Schwarz, Z. Hygg. Infektionskrankh. 67 [1910] 391/428, 404).
[Edited on 6-2-2009 by Formatik]
|
|
|
Nerro
National Hazard
Posts: 596
Registered: 29-9-2004
Location: Netherlands
Member Is Offline
Mood: Whatever...
|
|
Cr<sub>2</sub>O<sub>3</sub> will oxidize ammonia in air if you heat it to glowing. I did it as a freshman.
I believe the Cr<sub>2</sub>O<sub>3</sub> is reduced to CrO<sub>2</sub> - that's Cr(IV) - which is subsequently
oxidized by oxygen from the air.
#261501 +(11351)- [X]
the \"bishop\" came to our church today
he was a fucken impostor
never once moved diagonally
courtesy of bash
|
|
|
hissingnoise
International Hazard
Posts: 3940
Registered: 26-12-2002
Member Is Offline
Mood: Pulverulescent!
|
|
| Quote: | Originally posted by Mr. Wizard
Nobody ever admits they were wrong :-) I suspect you are actually right. |
I only admit to the most glaring mistakes, Mr. Wizard.
The wiki article, though, mentions an NH3 ozonide prep from bubbling ozone through a solution of Ca in liquid NH3.
The calcium doesn't react, apparently.
|
|
|
Xenoid
National Hazard
Posts: 775
Registered: 14-6-2007
Location: Springs Junction, New Zealand
Member Is Offline
Mood: Comfortably Numb
|
|
I came across this thesis on the catalytic oxidation of ammonia:
http://alexandria.tue.nl/extra2/200210267.pdf
Although primarily concerned with oxidation of ammonia to N2 for pollution control, it has some interesting data (see pages 7 & 8) on the use of
certain metals and oxides to oxidise ammonia to various combinations of N2, NO and N2O.
[Edited on 6-2-2009 by Xenoid]
|
|
|
AndersHoveland
Hazard to Other Members, due to repeated speculation and posting of untested highly dangerous procedures!
Posts: 1986
Registered: 2-3-2011
Member Is Offline
Mood: No Mood
|
|
What I am most interested in is what can oxidize NH3 to oxides of nitrogen at room temperature. Supposedly, hydrogen peroxide will react with ammonia
in the presence of a catalyst, but there seems to be virtually no information about this in the literature. (whether the catalyst is iron salts or
acetamide?)
I would like to share an interesting reaction from the literature of Marcellin Berthelot, dry ammonia gas reacts with the nitrogen dioxide and nitric
oxide, at room temperature,
(2)NO2 + (2)NO + (4)NH3 --> (2)NH4NO2 + (2)H2O + (2)N2
Solid ammonium nitrite inside a tube explodes if heated on a water bath to between 60-70degC. And the substance gradually decomposes at room
temperature, slower if cold, or faster in aqueous solutions, forming nitrogen gas.
While thinking about room temperature redox reactions with NH4NO3 we usually think about oxidizing it, but it is also possible to reduce it.
The reaction is slow, but gradually moves forward when nitric oxide is passed into a boiling aqueous solution of NH4NO3 for several hours. It should
be noted that nitric oxide alone is unable to to react with NH4ClO4 in the same way.
(14)NH4NO3 + (6)NO --> (24)H2O + (13)N2 + (8)HNO3
This is assuming excess NH4NO3, if there is too much NO, the HNO3 will be reduced to NO2.
[Edited on 3-10-2011 by AndersHoveland]
I'm not saying let's go kill all the stupid people...I'm just saying lets remove all the warning labels and let the problem sort itself out.
|
|
|
Formatik
National Hazard
Posts: 927
Registered: 25-3-2008
Member Is Offline
Mood: equilibrium
|
|
There is info on it, but you have not looked hard enough. The following below from Gmelin mentions the catalysts are NaOH, Na2CO3, Zn dust, Pt, or
Pd-H2.
Ammonium nitrite is also a sensitive explosive with about the power of tetryl: http://www.sciencemadness.org/talk/viewthread.php?action=pri...
Reaction between H2O2 and aq. NH3 can also give an azine when a ketone is present. This is the basis for the "Peroxide process" for production of
hydrazine.
[Edited on 5-10-2011 by Formatik]
|
|
|
AndersHoveland
Hazard to Other Members, due to repeated speculation and posting of untested highly dangerous procedures!
Posts: 1986
Registered: 2-3-2011
Member Is Offline
Mood: No Mood
|
|
What exactly is the chemistry between H2O2 and NH4OH ? Does acetone serve as a necessary catalyst for the reaction? When the two are mixed, there is
no observable reaction, and there are many references in the literature to mixes of the two chemicals, presumably without reaction.
Under what conditions can ammonia be oxidized to NH4NO2 ? Excess H2O2 would presumably oxidize the NH4NO2 to NH4NO3. Nitrite is a much more reactive
reducing agent than ammonia, so to obtain any nitrite one would think that a large excess of ammonia would have to be used.
Hydogen peroxide slowly decomposes in aqueous alkaline solution, so one would expect NH4OH/H2O2 solutions to gradually decompose, either with the
liberation of oxygen, or possibly the oxidation of ammonia.
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.
I'm not saying let's go kill all the stupid people...I'm just saying lets remove all the warning labels and let the problem sort itself out.
|
|
|
Formatik
National Hazard
Posts: 927
Registered: 25-3-2008
Member Is Offline
Mood: equilibrium
|
|
Quote: Originally posted by AndersHoveland  | | What exactly is the chemistry between H2O2 and NH4OH ? Does acetone serve as a necessary catalyst for the reaction? When the two are mixed, there is
no observable reaction, and there are many references in the literature to mixes of the two chemicals, presumably without reaction.
|
The Peroxide process was invented by Produits Chimiques Ugine-Kuhlman. The reaction is done from H2O2, 2-butanone, and NH3 in a 1:2:4 respective molar
ratio at atmospheric pressure and 50 C. The formation of azine is due to an intermediate that can oxidize the ammonia. The Pechiney-Ugine-Kuhlmann process.
| Quote: | | 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.
Attachment: Peroxide-process.pdf (554kB)
This file has been downloaded 14868 times
Attachment: Hoppe-Seyler.pdf (435kB)
This file has been downloaded 850 times
|
|
|
AndersHoveland
Hazard to Other Members, due to repeated speculation and posting of untested highly dangerous procedures!
Posts: 1986
Registered: 2-3-2011
Member Is Offline
Mood: No Mood
|
|
Thanks very much, Formatik.
If I can now be allowed to speculate a little into the reaction.
So it appears that H2O2 can much more readily attack imines, R2C=NH, than it can plain ammonia? I thought I came across a book that stated that there
is no reaction between H2O2 and imines without the use of some acetamide, which serves as a catalyst.
"A process developed by Produits Chimiques Ugine Kuhlmana (PcUK), and practiced by Atofina (France) and Mitsubish Gas (Japan) involves the oxidation
of ammonia by hydogen peroxide in the presence of butanone (MEK) and another component that apparently serves as an oxygen-transfer
agent. The reaction is carried out... at at 50degC. The ratio of H2O2/MEK:NH3 used is 1:2:4. Hydrogen peroxide is activated by acetamide and
disodium hydrogen phosphate (117). The mechanism of this reaction involves an activation of the ammonia and hydrogen peroxide because these compounds
do not themselves react (118-121). It appears that acetamide functions as an oxygen transfer agent, possibly as the iminoperacetic acid, HOOC(=NH)CH3,
which then oxidizes the transient Schiff base formed between MEK and ammonia to give give the oxaziridine, with regeneration of acetamide."
(117) U.S. Pat. 3,962,878 (aug. 3, 1976), J.P. Schirmann, J. Combroux, and S. Y. Delavarenne
(118) J.P. Schirmann and S. Y. Delavarenne, Tetrahedtron :ett. 635 (1972)
(119) E. G. E. Hawkins, J. Chem. Soc. C, 2663 (1969)
(120) E. Schmitz, Chem. Ber. 97, 2521 (1964)
(121) Can. Pat. 2,017,458 (Nov. 24, 1990), J.P. Schirmann, J. P. Pleuvry, and P. Tellier (to Atochem)
D. Todd, in R. Adams, ed.,Organic Reactions, Vol. 4, John Wiley & Sons, Inc., New York, 1948, Chapt. 8. "Hydrazine and its Derivitives"
One wonders why nitrite forms rather than nitrate, since the H2O2 should readily oxidize the nitrite. One possibility may be that nitrite is initially
the predominant product, but as nitrate begins to be formed, it alters the equilibrium, allowing formation of more nitrogen dioxide rather than nitric
oxide.
(2) NO2[-] + (2) H[+]aq <==> H2O + NO + NO2
NO2[-] + NO3[-] + (2) H[+]aq <==> H2O + (2)NO2
Nitrogen dioxide is known to be able to attack ammonia.
(2)NO2 + (2)NO + (4)NH3 --> (2)NH4NO2 + (2)H2O + (2)N2
In any case, solutions of ammonium nitrite decompose on heating to 60 to 70degC.
Perhaps the very slow spontaneous reaction of aqueous ammonia with H2O2 takes place the small equilibrium with amide anions, NH2[-].
NH2[-] + H2O2 --> NH2* + OH* + OH[-]
Even in NH4OH, there is a very slight equilibrium with amide ions.
(2)NH3 <==> NH2[-] + NH4[+]
In the base catalysed decomposition of hydrogen peroxide, the mechanisms is presumably
H2O2 + OH[-] --> HOO[-] + H2O
HOO[-] + H2O2 --> HOOOH + OH[-]
HOOOH --> HOOO[-] + H[+]aq
HOOO[-] --> OH[-] + O2
Is the dihydrogen trioxide intermediate reactive enough to oxidize ammonia? H2O3, also known as "trioxone" quickly decomposes into water and singlet
[excited] oxygen. The reaction is actually reversible, but much less favorable.
(2)H2O + O2* <==> H2O + H2O3
One question I would have though is what the dominant reaction in the spontaneous decomposition of NH4OH-H2O2 solutions. Is the reaction mainly the
base catalysed decomposition of H2O2 into O2 ? Or is the main reaction the formation of ammonium nitrite? What proportion of the final product
contains ammonium nitrate rather than nitrite?
And why does the reaction proceed so much faster when heated? Relative to the energy required to break most chemical bonds, the boiling point of the
solution is not really much hoter than room temperature. Why would this make so much difference?
[Edited on 5-10-2011 by AndersHoveland]
I'm not saying let's go kill all the stupid people...I'm just saying lets remove all the warning labels and let the problem sort itself out.
|
|
|
AndersHoveland
Hazard to Other Members, due to repeated speculation and posting of untested highly dangerous procedures!
Posts: 1986
Registered: 2-3-2011
Member Is Offline
Mood: No Mood
|
|
Can diatomic nitrogen be alkylated to form diazo salts?
Possibly heat benzene fluorosulfonate with nitrogen gas?
C6H5-OSO2F + N2 --> C6H5N[+]ΞN [-]OSO2F
I'm not saying let's go kill all the stupid people...I'm just saying lets remove all the warning labels and let the problem sort itself out.
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
NH3/H2O2 Photolysis at pH 9.3 to Nitrite/Nitrate
I came across an interesting source noting the photolysis based oxidation of NH3 with H2O2 at pH 9.3 with the formation of some nitrite and nitrate.
Source: "Removal of Ammonia by OH Radical in Aqueous Phase" by Lihuang, Liangle,.., Department of Environmental Science and Engineering, Fudan
University, Shanghai, China, 2008. To quote:
"Results show that the •OH, generated by H2O2 photolysis, could oxidize NH3 to NO2- and further to NO3-. Removal efficiencies of ammonia were low
and were affected by initial pH value and ammonia concentration."
However, the study was carried out at normal water temperature (being an environmental study) of 20 C. Question: would raising the temperature to say
50 C to 60 C increase nitrite/nitrate yield?
A possible reaction chain is given as, to quote:
"H2O2 --hv->2•OH (2)
NH3 + •OH --> •NH2 + H2O (3)
•NH2 + H2O2 --> •NHOH + H2O (4)
•NH2 + •OH --> NH2OH (5)"
Futher comments include:
"That the absorption location and peak of 260-290nm band were identical to that of •NHOH acquired by Simic and Hayon (28) through pulse radiolysis
study confirmed that this absorption band was mainly originated from the reaction between •NH2 and H2O2.
Once formed, •NHOH would decay in no time due to its unstable nature in water. However, its decay pathway had hardly ever been reported in previous
literature. In this study, the decay of •NHOH was found to produce species with absorption band centered in 300 nm. As far as the absorption
spectrum was considered, this species was NH2O2-, the product of the reaction between •NHOH and HO2- (29). NH2O2 - was able to dissociate to NO2-.
NO2- would gradually be oxidized to NO3- by •OH or H2O2.
When attacked by •OH, ammonia would be oxidized to •NH2. Then, •NH2 would be rapidly oxidized to•NHOH and further to NH2O2-. Afterward, the
unstable NH2O2- splits to NO2-, which could be oxidized to NO3-. These reaction processes were fairly comprehensible since they were in good
accordance with the results in Section 3.1. The concentration of NO2- first increased and then decayed with irradiation time while the concentration
of NO3- ascended monotonously with irradiation time."
Full paper available at http://www.google.com/url?sa=t&rct=j&q=nh3%20%2B%20h...
Recently, before coming across this source, I performed the following experiment: NH3/H2O2/Na2CO3.H2O2/Cu. When this solution was placed in sunlight
and allowed to warm on a hot day, I noticed an unusual result unlike prior trials absence the application of some sunlight and mild heating. Within 12
hours, a significant gas evolution occurred. Speculation, this may be due to the creation of gaseous N2 from the decomposition reaction of a created
nitrite:
NH4NO2 --> N2 (g) + 2 H2O
Note, the amount of nitrogen gas formed can be measured (via a water displacement, for example) and would indicate corresponding molar amount of
ammonium nitrite creation.
I plan on repeating the experiment and others may wish to study the attached reference and, with precautions, repeat as well on a small scale.
------------------------------------------------------
For those that have some doubts on the formation of any NH4NO2 in the reaction mixture NH3/H2O2/Na2CO3.H2O2/Cu, please see "Kinetics and Mechanism of
Copper Dissolution In Aqueous Ammonia" fully available after signing on to ones Facebook account at http://www.academia.edu/292096/Kinetics_and_Mechanism_of_Cop... where the author cites a rate for Cu dissolution as a function of available O2
and NH3 and a side reaction involving ammonium nitrite formation. To cite some of the author's reactions:
2 Cu + 4 NH3 + 1/2 O2 + H2O --> 2 [Cu(NH3)2]OH
2 [Cu(NH3)2]OH + 4 NH3 (aq) + 1/2 O2 + H2O --> 2 [Cu(NH3)4](OH)2
Cu + [Cu(NH3)4](OH)2 <---> 2 [Cu(NH3)2]OH
And, with respect to the above topic, a relevant side reaction:
2 NH3 (aq) + 3 O2 + 2 OH- --> 2 NO2- + 4 H2O
which, per my discussion/research cited above, NH4NO2 (and NH4NO3) formation may be enhanced via photolysis and mild heating. The application of heat
leading to an increased presence of ammonium nitrite is not surprising given that NH4NO2 is an endothermic compound (and also characteristically
unstable and even explosive, see its MSDS available at this link: http://images.mpbio.com/docs/msds/aust/en/205055-EN-AUST.pdf ).
[Edited on 12-8-2013 by AJKOER]
[Edited on 12-8-2013 by AJKOER]
|
|
|
|
![[*]](images/xpblue/default_icon.gif){kind=icon}
