AJKOER
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New Preparations Paths for Copper Ammonium Nitrate
I noticed recently that dissolving Cu in aqueous ammonia in the presence of H2O2 and NaCl appears to form the insoluble Copper Ammonium chloride,
[Cu(NH3)4(H2O)2]Cl2. As a reference on the solubility issues of the latter, please see http://www.google.com/url?sa=t&rct=j&q=&esrc=s&a...
So, if one were to dissolve Cu in aqueous ammonia and H2O2 in the presence of say KNO3 would one form Copper Ammonium nitrate, [Cu(NH3)4(H2O)2](NO3)2
?
Having attempted this experiment (now for 12 hours), I can report a possible affirmative. The dissolving of the copper (my Cu source was a pre-1982 US
penny consisting of 95% Cu and 5% Zn) in an excess of dilute aqueous NH3 and H2O2 appears to proceed as usual with the creation of a blue-greenish
solution for about 10 hours, then there was a gas evolution (O2 or N2 given the sudden nature from the decomposition of any formed HNO2 or nitrite,
see "Kinetics and Mechanism of Copper Dissolution In Aqueous Ammonia" available online) and the reaction appears to halt. I added more NH3 and in the
next 2 hours the top of the solution became completely clear. Also, below the clear solution were now layers. One was a thin band of bluish-green, and
a bottom thicker layer of more intense blue.
This bottom layer has the characteristic color of the tetra-amine complex [Cu(NH3)4]2+, and given the evident more limited solubility of the layer, I
would suspect, indeed, the creation of Copper Ammonium nitrate.
Note, as a point of interest, there is actually nothing sacred about the tetraamminediaquacopper cation as depending on the ammonia concentration in
solution, one could have anywhere from [Cu(NH3)(H2O)5]2+ (the greenish-blue thin layer) to [Cu(NH3)5(H2O)]2+ (a royal blue complex), with the latter
occurring in very concentrated ammonia solutions. Reference: https://docs.google.com/viewer?a=v&q=cache:IjHK0vuBZhcJ:...
Next, the CAN layer will have to dried, a tedious and expensive part of the usual synthesis of this salt.
However, per this link http://www.dtic.mil/dtic/tr/fulltext/u2/629884.pdf , I am interested in determining new paths that follow the apparently unwanted natural ways
that the dry CAN salt seems to form (like possibly exposing dry Basic copper carbonate to ammonia and Nitric acid fumes from the sublimation of mildly
heated NH4NO3).
[EDIT] At 19 hours, I cooled the solution and the clear top became cloudy. This resembled my original room temperature KNO3 in water solution. The
source of the Potassium nitrate: Tree stump remover.
Here is a picture which hopefully displays the layering:
[Edited on 7-9-2013 by AJKOER]
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DubaiAmateurRocketry
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Interesting.
It looks more like a cupper II nitrate IV ammonia hydrate.
This looks like an excellent explosive.
[Edited on 7-9-2013 by DubaiAmateurRocketry]
[Edited on 7-9-2013 by DubaiAmateurRocketry]
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Fantasma4500
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it isnt, its one of the weak ones; but however it is partially useful..
what would be more interesting would be to isolate the anhydrous TACN instead of the hydrated..
a guy here on youtube showed how he made it with..
NH4NO3
CuO
Hexamine
then im not sure if he added anything else?
procedure was to boil it all together, heat it so the NH4NO3 melted and at one point the melt would have a darker colour, or well dark blue colour..
by what i know that means the Cu++ has bound to water, i might be wrong, anyhow copper ions have bound with water to make it bluish
there is a video on it linked to the thread also
but what im just wondering about is... how can you possibly extract the TACN (as in ANHYDROUS) it could have really neat properties, considering that
with water in the molecules its still useful as secondary
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Krakermanworks
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I tried copper, amonia and kno3 had a funny result the crystals were light blue. Would the turnout be different if I added h202, thanks .
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AndersHoveland
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Do you need the ammonia? Is the reaction much slower without it? (just to get the copper to dissolve to begin with)
Might be more practical to dissolve copper, obtaining the chloride, then precipitate out the carbonate, then work from there.
Another possibility to consider, Cu(OH)2 can dissolve if the pH is high enough. Coppric oxide is slightly amphoteric, a concentrated NaOH solution
will dissolve it. Perhaps if one put the copper in NaOH and gradually added H2O2. a problem here would be that high pH catalyzes the decomposition of
H2O2
[Edited on 7-9-2013 by AndersHoveland]
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AJKOER
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Quote: Originally posted by Krakermanworks  | I tried copper, amonia and kno3 had a funny result the crystals were light blue. Would the turnout be different if I added h202, thanks .
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In a word, yes.
Per the reference on the kinetics of the reaction I gave above, oxygen is essential for a basically electrochemical oxidation of the copper metal with
the NH3 reputedly acting as a catalyst through the formation of copper ammonium complexes. So, H2O2 (even dilute) is a definite plus for moving the
reaction forward.
Now, for further clarification on a possibly important point, the reaction pictured above did take place over night (that is, mostly in the dark). I
recently provided a reference on the photolysis of ammonia and H2O2 which I suspect can change aspects of the above reaction including yield and
speed. As such, I will repeat the reaction in sunlight to see if I can ascertain any obvious differences.
[Edited on 8-9-2013 by AJKOER]
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Krakermanworks
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This is what it looks like
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Krakermanworks
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Oh and thanks heaps I will try with different concentrations of h202, what
concentrations would you suggest? Thanks it'll be a great help
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AJKOER
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OK, a few words of advice from working on this reaction.
1. Dilute H2O2 and NH3 have been fine in obtaining a royal blue solution.
2. Once one has obtained the deep blue solution, do not add more H2O2 (the result on one occasion was a surprising gold colored solution of nano-sized
copper), or even more NH3 (loss of the royal blue).
3. As a variation on the synthesis, I first prepared aqueous (NH4)2SO4 as follows:
MgSO4.7H2O + 2 NH3 + 2 H2O --> Mg(OH)2 (s) + (NH4)2SO4 +
7 H2O
After filtering out the Mg(OH)2, I added KNO3:
(NH4)2SO4 + 2 KNO3 --> K2SO4 + 2 NH4NO3
I then further added KNO3 to the solution (to increase the Potassium ion concentration) and cooled to remove hopefully all the K2SO4. I then reacted
this mixture in place of the KNO3 in the original preparation (that is, added NH3, H2O2 and Cu). The result was a surprising more vigorous reaction
than with just KNO3. I repeated the whole procedure except the cooling to remove the K2SO4, only NOT to obtain a royal blue solution (apparently the
sulfate does complex interfering with the formation of CAN as I originally suspected, hence the need for removal).
4. Finally, I will address some of the suspected chemistry on my recommended new aqueous path as it involves a potential safety issue. Today, I
experienced a loud retort when I opened a plastic reaction vessel which I had compressed to some extent initially to allow for gas expansion. With
several observations of suspicious rapid and significant solution gas evolution, I now am convinced that the side reactions discussed in "Kinetics and
Mechanism of Copper Dissolution In Aqueous Ammonia" involving the formation of Nitrous acid, HNO2, and NH4NO2 are quite real and significant in this
preparation. Shared properties of HNO2 (Wikipedia http://en.wikipedia.org/wiki/HNO2 ) and NH4NO2 (see Wikipedia http://en.wikipedia.org/wiki/NH4NO2 ) include rapid and significant Nitrogen, or NO and NO2 gases in the case of HNO2, evolution with change in
pH, concentration and temperature. A significant statement with respect to this preparation of CAN is that HNO2 is only stable in cold, dilute aqueous
solution, so hot concentrated solutions may not render similar results.
As a speculation, a key reaction sequence may be the reaction of KNO3 with the side product (Ammonium nitrite) of the electrochemical oxidation of the
metal copper in the presence of ammonia and oxygen:
NH4NO2 + KNO3 <--> KNO2 + NH4NO3
which interacts with the major electrochemical and chemical equilibriums reactions of Cu, NH3 and O2 as cited per the reference as follows:
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
So, upon addition of KNO3 forming some NH4NO3 (and KNO2) per the reaction above and further cooling:
[Cu(NH3)4](OH)2 + 2 NH4NO3 --> [Cu(NH3)4](NO3)2 (s) + 2 NH3 + 2 H2O
At this time, I would not also completely rule out the formation of any Copper ammonium nitrite, which may present some new unsuspected energetic
issues (for example, NH4NO3 is a relatively safe commercial HE, while NH4NO2 is an impractical HE with toxicity issues).
Also, I would also be concerned on heating an acidified form of the solution just prepared relating to known stability issues with hot aqueous NH4NO3
in the presence of metallic impurities (including Copper, Tin and Nickel see http://www.google.com/url?sa=t&rct=j&q=ammonium%20ni... ).
[Edited on 10-9-2013 by AJKOER]
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woelen
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There are some plain errors in your post.
1) You never can have nano-copper particles from copper(II), NH3, and H2O2. The solid you obtain most likely is some peroxide of copper. This is
unstable and slowly looses oxygen and in the presence of NH3 forms the deep blue ammonia complex again.
2) NH4NO2 is not only impractical, it hardly can exist. Toxicity issues are not the real issue (nitrite can be toxic, but if handled carefully it is
no real concern, it even is used as a food additive for curing meat). The real issue is that NH4NO2 is very unstable and preparation of this compound
is not possible from aqueous solution. You getdecomposition to water and N2. I tried this personally and wrote about that on sciencemadness some time
ago (search for ammonium nitrite in my posts).
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AJKOER
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Woelen:
Thanks for reading my post.
With respect to your 1st point I would completely agree if I have not performed the reaction myself. My speculation on what occurred was based on the
known (and personally witnessed previously by destroying Cu(NH3)4 complex with, I believe, an oxychloride) decomposition of tetraamminediaquacopper
cation to form nano-sized particles (it appears to a gold colored solution as the Copper(?) particles are too small to be seen by the human eye) with
dilute H2O2. Also, upon addition of dilute ammonia, the solution became blue-green and not royal blue.
On your second point, it is not speculation that HNO2/NH4NO2 is formed (as, however, a side product). Please review my cited reference "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...
My speculation (so noted) is that NH4NO2/HNO2 apparently do form in concordance with the cited reference, and are created/decomposed being a visible
intermediary in my reaction. My several observations are that within the first 1O minutes of the reaction, you can personally observed the
decomposition of any formed HNO2/NH4NO2 (well, at least if one uses cheap household ammonia which foams in this reaction and actually creates a foam
column as decomposition occurs). To be honest, the first time I saw this gas evolution, I thought all my H2O2 have suddenly decided to decompose on me
forming O2, and not N2 or something else. However, based on the cited reference, the stopping and then restarting of the this particular reaction
mixture with small additions of dilute H2O2 in the presence of excess ammonia (so H2O2 is, in fact, consumed), and my near explosive gas evolution
experience today, that I felt first hand (which I seriously doubt was from unconsumed and relatively small initial dilute dose of H2O2, but more
likely from the larger quantity of NH3 acting as a nitrogen source), it is not likely O2 in my opinion.
Now as pundits have been debating the precise mechanics of the NH3, O2 and Cu reaction for decades (which may or may not be completely accepted as of
today), I am not so bold to believe that I am precisely correct in my suggested reaction path with the addition of KNO3 to the mix.
However, I did bring up the possible chemistry because of what I believe is a real danger (gas evolution issue) to someone repeating my preparation
and also prospective issues upon drying and experimenting with the salt (given the possible presence of a dangerous unstable nitrite added to the mix,
which I hope is completely not the case).
Anyone repeating my experiment, do so in small quantities and take notice of my comments.
[Edited on 10-9-2013 by AJKOER]
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AndersHoveland
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What about dissolving Cu with NH4NO3 and a little H2O2 ? The reaction would likely be slow, but it could be a potentially easy route. then dry the
hygroscopic copper complex with some acetone
H2O2 can oxidize NH4OH, but the reaction tends to be very slow under normal conditions. If you just mix the two together there will
not be any immediate reaction. (I believe the reaction proceeds through base-catalyzed decomposition of the peroxide, generating transient free
radicals, if you were wondering, or possibly through the extremely small equilibrium of NH2- ions in solution)
Also, nitrite can act as a reducing agent, it is even slowly oxidized on exposure to air, so if there is excess H2O2 in the solution, any nitrite will
be oxidized faster than it is formed.
If you want to make nitrite through this reaction, you should do some research into the optimal reaction conditions.
[Edited on 10-9-2013 by AndersHoveland]
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AJKOER
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AndersHoveland:
Thanks for contributing.
Now, on your question/comment: "What about dissolving Cu with NH4NO3 and a little H2O2 ? The reaction would likely be slow, but it could be a
potentially easy route. then dry the hygroscopic copper complex with some acetone", yes I was attempting this by preparing NH4NO3 by the reaction of
KNO3 and Ammonium sulfate. Here is a Wikipedia comment on the preparation of Ammonium nitrate (link: http://en.wikipedia.org/wiki/NH4NO3 ):
"Ammonium nitrate can also be made via metathesis reactions:
(NH4)2SO4 + 2 NaNO3 → 2 NH4NO3 + Na2SO4 "
where I am using KNO3 in place of NaNO3. However, my CAN preparation reaction appeared to be not slow as being an electrochemical reaction in an ionic
solution.
With respect to your comment "H2O2 can oxidize NH4OH, but the reaction tends to be very slow under normal conditions. If you just mix the two together
there will not be any immediate reaction. (I believe the reaction proceeds through base-catalyzed decomposition of the peroxide, generating transient
free radicals..", you sound very similar to my reference citing the oxidation of NH3 with O2 in the presence of Copper ammonium hydroxide acting as
the catalyst.
Now, on your other question, I refer you to a prior thread on the "Electrolysis of Ammonia" (link: http://www.sciencemadness.org/talk/viewthread.php?tid=18912#... ) where I cited some research (source: Journal Chemical Society, London, Volume
88, Part 2, page 242 at http://books.google.com/books?pg=PA242&lpg=PA242&dq=... ) noting formation of Ammonium nitrite and nitrate, and equilibrium shifts between
them occurring as a function of reaction conditions. For example to quote a reference:
"Electrolytic Oxidation of Ammonia to Nitrites. Erich Muller and Fritz Spitzer (Ber., 1905, 38, 778—782. Compare Traube and Biltz, Abstr., 1904, ii,
727).—In the presence of a small amount of sodium hydroxide, ammonia may be oxidised electrolytically to nitrite even in the absence of copper
compounds.
In the presence of copper hydroxide and sufficient alkali, the oxidation of ammonia to nitrite does not cease suddenly when the nitrite concentration
has reached a certain value, but appears to proceed quite independently of the nitrite concentration. In these experiments, the oxidation was allowed
to proceed for a comparatively short time only, so that the amount of alkali present was not greatly reduced. The formation of nitrite is intimately
connected with the amount of alkali present, and when no sodium 'hydroxide is present, but only ammonia, nitrite, and copper hydroxide, it is found
that the nitrite is transformed into nitrate more rapidly than the ammonia into nitrite, and thus the concentration of the nitrite tends to decrease.
Nitrogen is also formed during the oxidation. J. J. S."
The source also notes, to quote:
"In continuation of the previous experiments, the influence of changing the concentration of the free alkali or ammonia on the rate of the
electrolytic oxidation of ammonia has been investigated. In presence of much ammonia, the amount of nitrite can be increased to about 11 per cent,
before oxidation to nitrate begins, whilst from an 11 per cent, nitrite solution to which ammonia, sodium hydroxide, and copper hydroxide had been
added a solution containing as much as 17 per cent, nitrite was obtained on hydrolysis."
In the current context, I doubt if there is any excess of my dilute H2O2, but upon reviewing that research further maybe one of us will express an
opinion on reaction conditions so that one does not have to be concerned with any nasty nitrite acting as a primer.
With respect to a possibly successful CAN formation reaction, the following comments per the above cited thread appear worthy of repeating. To quote
a reference (a 2011 study titled "Copper-Mediated Non-Enzymatic Formation of Nitrite from Ammonia and Hydrogen peroxide at Alkaline pH") cited there:
"The reaction is mediated by copper (II) as it fails to occur in absence of copper", and that the best order of addition of reactants is Cu then
aqueous NH3 and finally H2O2. The author also notes the need for excess ammonia, to quote: "as it is needed to maintain: (i) solubility of copper;
(ii) optimal alkalinity for expression of reducing potential of hydrogen peroxide; (iii) adequate concentration of free ammonia; and (iv) conversion
of nitrous acid to ammonium nitrite."
[Edited on 10-9-2013 by AJKOER]
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AndersHoveland
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| Quote: Originally posted by AJKOER | | The formation of nitrite is intimately connected with the amount of alkali present, and when no sodium 'hydroxide is present, but only ammonia,
nitrite, and copper hydroxide, it is found that the nitrite is transformed into nitrate more rapidly than the ammonia into nitrite, and thus the
concentration of the nitrite tends to decrease. |
This could well be. The oxidation of aqueous sodium nitrite solutions may proceed through an equilibrium, where it is actually the species nitric
oxide being oxidized. Having a large excess of hydroxide ions could prevent this equilibrium from shifting.
2 H2O <==> H3O+ + OH-
NO2- + H3O+ <==> HNO2 + H2O
2 HNO2 <==> H2O + NO2 + NO
so it appears that NaOH greatly helps increase yields
(read somewhere that sodium carbonate also works)
[Edited on 11-9-2013 by AndersHoveland]
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AJKOER
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Yes, but that quoted sentence is perhaps more pertinent/meaningful in the current context when rewritten in the negative as follows:
The formation of nitrite is intimately connected with the amount of alkali present, and when sodium 'hydroxide is present, with ammonia, nitrite, and
copper hydroxide, it is found that the nitrite is transformed into nitrate more slowly than the ammonia into nitrite, and thus the concentration of
the nitrite tends to increase.
Also, rewriting your reactions to show a net:
2 H2O <==> H3O+ + OH-
NO2- + H3O+ <==> HNO2 + H2O
HNO2 <==> 1/2 H2O + 1/2 NO2 + 1/2 NO
---------------------------------------------------
Net:
1/2 H2O + NO2- <==> OH- + 1/2 NO2 + 1/2 NO
So, increasing the presence of OH- moves the above equilibrium to the left (NO2- concentration increases) in agreement with the observational based
author's statement.
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AndersHoveland
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I wonder... what if one passed chloramine gas into this ammonia solution containing nitrite, after most of the base was neutralized with a little
acid?
The chloramine would instantly react with the excess ammonia to form hydrazine, and if there was nitrite floating around in there at the right pH, the
hydrazine would be oxidized by nitrite faster than being oxidized by chloramine, and hydrazoic acid could be formed in solution...
I am not sure how fast this would be oxidized by chloramine. I think chloramine can directly oxidize ammonia, whereas it would have to first hydrolyze
into hypochlorous acid to oxidize hydrazoic acid, the reaction rate would be much lower.
(chloramine gas made in separate container by simply mixing ammonia and hypochlorite bleach in optimal proportions)
[Edited on 12-9-2013 by AndersHoveland]
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AJKOER
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OK, interestingly but a bit off topic. However, looking back on my prior threads I did find the following comment by myself (see http://www.sciencemadness.org/talk/viewthread.php?tid=23807&... ) that is more on topic:
"Also, those forming Ammonium nitrite via NH3 and H2O2 in the presence of a lot of NaOH (or Na2CO3, which are cited catalysts) may also, by NaNO2
formation, be able to avoid the decomposition issue associated with NH4NO2. A source cites the reaction (see http://books.google.com/books?id=6OcDAAAAMBAJ&pg=PA277&a... ):
NH4NO2 + NaOH --> NaNO2 + H2O + NH3 (g) "
My point today being that substituting [Cu(NH4)4(H2O)2](OH)2 for NaOH may allow one to form some unwanted [Cu(NH4)4(H2O)2](NO2)2.
However, to address your chloramine comment, here is what I said in that same thread:
"On second thought, I now have stronger reservations on the use of NaOCl or HOCl in the possible presence of NH3. The interaction of any formed NH2Cl
and NaNO2 may proceed as follows (speculation):
NaNO2 + NH2Cl --?--> NH2NO2 + NaCl
where the possible formation of nitramide, or a voilent decomposition, is viewed as problematic. If this reaction is successful, replacing NaNO2 with
AgNO2 to produce an insoluble AgCl may be a preferrable procedure. [EDIT][EDIT] See Franklyn's prior (from 2007) comments/research on this very idea
titled "nitramide from chloramine and nitrite" and also the response by AndersHoveland at http://www.sciencemadness.org/talk/viewthread.php?tid=6042 . Also, a source (Canis, C., Rev. Chim. Minerale, 1964, 1, 521) notes that "Nitramide
is quite unstable and various reactions in which it is formed are violent. Attempts to prepare it by interaction of various nitrates and sulfamates
showed that the reactions became explosive at specific temperatures." Also with alkalies (source: Thiele, J. et al., Ber., 1894, 27, 1909) "A drop of
conc. alkali solution added to solid nitramide causes a flame and explosive decomposition" and also "Nitramide decomposes explosively on contact with
conc. sulfuric acid."
Caution: Chloramine is known to attenuated the posionous properties of other compounds (for example, with CH3OH), so exercise appropriate safety in
this chloramine/nitrite combination."
where my last comment relates to the ability of NH2Cl to magnify the poisonous nature of certain compounds, possibly including NH4NO2. Now Ammonium
nitrite is described by Wikipedia as "acutely toxic to both humans and aquatic organisms.[1]" link: http://en.wikipedia.org/wiki/NH4NO2 , which may seem strange to some (probably due to the properties and use of NaNO2), however, I do believe
oral, or inhalation of the dust, may present a serious risk. The explanation as to why is a point that a biochemist may wish to comment on
(interestingly, yesterday I read up on why chlorates are very toxic on ingestion by oxidizing Iron and disrupting its ability to absorb and transport
oxygen in the blood, but on skin contact with an aqueous solution, per my speculation, chlorates may be largely decomposed before absorption by
blood).
[Edited on 12-9-2013 by AJKOER]
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AndersHoveland
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I doubt that would be a problem. If chloramine were so vulnerable to nucleophilic substitution, it would immediately react with any base, forming
NH2OH, which would immediately get destructively oxidized. In fact, chloramine is not stable for long in alkaline solution, but the fact that
it can even be formed in ammonia shows the nucleophilic substitution is not immediate. So any such interaction would be insignificant.
If anything, the chloramine would oxidize some of the nitrite to nitrate.
NH2Cl + H2O <==> NH3 + HOCl
HOCl + HNO2 --> HCl(aq) + HNO3(aq)
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Culpable Cuprate
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| Quote: Originally posted by AJKOER | Woelen:
Thanks for reading my post.
With respect to your 1st point I would completely agree if I have not performed the reaction myself. My speculation on what occurred was based on the
known (and personally witnessed previously by destroying Cu(NH3)4 complex with, I believe, an oxychloride) decomposition of tetraamminediaquacopper
cation to form nano-sized particles (it appears to a gold colored solution as the Copper(?) particles are too small to be seen by the human eye) with
dilute H2O2. Also, upon addition of dilute ammonia, the solution became blue-green and not royal blue.
On your second point, it is not speculation that HNO2/NH4NO2 is formed (as, however, a side product). Please review my cited reference "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...
My speculation (so noted) is that NH4NO2/HNO2 apparently do form in concordance with the cited reference, and are created/decomposed being a visible
intermediary in my reaction. My several observations are that within the first 1O minutes of the reaction, you can personally observed the
decomposition of any formed HNO2/NH4NO2 (well, at least if one uses cheap household ammonia which foams in this reaction and actually creates a foam
column as decomposition occurs). To be honest, the first time I saw this gas evolution, I thought all my H2O2 have suddenly decided to decompose on me
forming O2, and not N2 or something else. However, based on the cited reference, the stopping and then restarting of the this particular reaction
mixture with small additions of dilute H2O2 in the presence of excess ammonia (so H2O2 is, in fact, consumed), and my near explosive gas evolution
experience today, that I felt first hand (which I seriously doubt was from unconsumed and relatively small initial dilute dose of H2O2, but more
likely from the larger quantity of NH3 acting as a nitrogen source), it is not likely O2 in my opinion.
Now as pundits have been debating the precise mechanics of the NH3, O2 and Cu reaction for decades (which may or may not be completely accepted as of
today), I am not so bold to believe that I am precisely correct in my suggested reaction path with the addition of KNO3 to the mix.
However, I did bring up the possible chemistry because of what I believe is a real danger (gas evolution issue) to someone repeating my preparation
and also prospective issues upon drying and experimenting with the salt (given the possible presence of a dangerous unstable nitrite added to the mix,
which I hope is completely not the case).
Anyone repeating my experiment, do so in small quantities and take notice of my comments.
[Edited on 10-9-2013 by AJKOER] |
I was just attempting essentially the same reaction hoping to get CAN ([Cu(NH3)4(H2O)]SO4+KNO3+H2O2+NH4OH). I added 'too much' H2O2 and obtained the
golden coloured solution, with a fine brown precipitate. Is it possible that copper azide could be formed during this reaction?
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Tellurium
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I would think, that Copper peroxide or Copper Oxide has formed. Probably the last.
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DraconicAcid
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The title compound isn't "copper ammonium nitrate", but tetraamminecopper(II) nitrate.
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
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