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[*] posted on 21-2-2012 at 05:47
Electrolysis of Ammonia


I'm just Curious...i want to electrolyse a mixture of ammonia and hydrogen peroxide...but as H2O2 is precious in my lab i won't like to waste it on something that might be unprofitable.
What are the products? Does NH4+ Get oxidized/reduced? I'm particularly interested in the NH4+ Ion...




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weiming1998
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[*] posted on 21-2-2012 at 06:14


I would think that the act of combining ammonia and hydrogen peroxide will give you ammonium nitrite according to the equation: 2NH3(aq)+2H2O2(Aq)===>NH4NO2+H2O, which the nitrite decomposes (explosive in acidic solutions!) by NH4NO2====>N2+2H2O. As for electrolysis, I don't know, but that will probably decompose before you can electrolyze a good amount of it.
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[*] posted on 21-2-2012 at 06:45


Whan ammonia is added to 3% hydrogen peroxide fizzing is observed. This is probably the result of catalytic decomposition from the alkalinity of the solution.



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Kola
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[*] posted on 21-2-2012 at 07:03


Quote: Originally posted by weiming1998  
I would think that the act of combining ammonia and hydrogen peroxide will give you ammonium nitrite according to the equation: 2NH3(aq)+2H2O2(Aq)===>NH4NO2+H2O, which the nitrite decomposes (explosive in acidic solutions!) by NH4NO2====>N2+2H2O. As for electrolysis, I don't know, but that will probably decompose before you can electrolyze a good amount of it.

does this mean that NH4+ is very stable in Aqueous solution?
Maybe i'll still try the electrolysis...at least some non spontaneous reactions might occur




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[*] posted on 21-2-2012 at 07:14


Quote: Originally posted by LanthanumK  
Whan ammonia is added to 3% hydrogen peroxide fizzing is observed. This is probably the result of catalytic decomposition from the alkalinity of the solution.



...Hint...basic piranha...
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[*] posted on 21-2-2012 at 08:31


Your basic piranha will OMNOMNOM your electrodes like charm. And it might cause an accident if the ingredients are mixed in the proper amounts.



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[*] posted on 21-2-2012 at 15:02


The ammonia will be oxidized even without hydrogen peroxide.

In one experiment, the electrochemical oxidation of ammonium ions to nitrate was less than 2% efficient.

One reason for the very low efficiency is probably that the oxidation is dependant on the formation of hydroxyl radicals at the anode. But at the same time, the positively charged anode would tend to repel away positively charged ammonium ions. This should not be so important if ammonium hydroxide is used, since most of the ammonia in equilibrium is actually in plain NH3 form.

Another explanation is that the density of hydroxyl radicals is very low, since the unstable radicals quickly recombine. So it is more likely that nitrogen will form from the oxidation of ammonia, since the nitrogen radicals will recombine before they have a chance to get oxidized.


[Edited on 21-2-2012 by AndersHoveland]
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[*] posted on 21-2-2012 at 19:04


Anders do you have a link to that? It rings bells but I can not find anything in my archives.


Also to the OP

Try this.
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[*] posted on 22-2-2012 at 01:27


Electrolysis of ammonia hardly works at all. Only a very small part of the ammonia is converted to ammonium ions and hydroxide ions, so the conductivity of the solution will be low, very low. You hardly get any current through the solution.

If you add hydrogen peroxide, then that hardly affects the outcome, still there is nearly no conductivity. Dilute hydrogen peroxide and ammonia do not react noticeably, so do not expect anything interesting to happen. Probably you just get (slow) decomposition of the hydrogen peroxide in the mildly alkaline solution.

The term 'basic piranha' does not apply at all. Dilute ammonia mixed with 3% hydrogen peroxide is rather tame and not dangerous at all. Piranha solution is at least 30% hydrogen peroxide, mixed with concentrated sulphuric acid. A basic equivalent might be a strong solution of NaOH in 30% hydrogen peroxide, but here we only talk of weak solutions.

If you perform electrolysis of a water-soluble ammonium salt, then at the cathode the following reaction occurs:

2NH4(+) + 2e --> 2NH3 + H2

What happens at the anode depends on the nature of the anode and the anionic part of the salt. If an inert anion (like sulfate or perchlorate) is used and a platinum anode, then oxygen is produced at the anode and a very small fraction of ammonium ions may be oxidized to a mix of nitrogen compounds.




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[*] posted on 25-2-2012 at 19:28


I recall from a prior Sciencemadness thread (see http://www.sciencemadness.org/talk/viewthread.php?action=pri... ) a claim that NH4OH + H2O2 + catalyst --Boil--> NH4NO2. Note, Wikipedia reiterates without the mention of the catalyst:

"It can also be prepared by oxidizing ammonia with ozone or hydrogen peroxide, or in a precipitation reaction of barium or lead nitrite with ammonium sulfate, or silver nitrite with ammonium chloride, or ammonium perchlorate with potassium nitrite. The precipitate is filtered off and the solution concentrated. It forms colorless crystals which are soluble in water and decompose on heating or in the presence of acid, with the formation of nitrogen.[3] 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"

Note, if NO and NO2 is formed in the presence of NH3, then a path to NH4NO2:

2 NO2 + 2 NO + 4 NH3 --> 2 NH4NO2 + 2 H2O + 2 N2 [1]

In the Sciencemadness reference, it is also stated that the nitrite yield is significant enough to demonstrate. For example, upon adding FeSO4 and H2SO4:

2 NH4NO2 + 2 H2SO4 + 2 FeSO4 --> Fe2(SO4)3 + (NH4)2SO4 + 2 H2O + 2 NO

However, this would generate an acidic environment and per the above admonition, I would test in small quantities only.

Interestingly, reputedly even dilute H2O2 could also elicit a reaction.

One of the catalysis cited included Na2CO3 (others, per my recollection, included NaOH, Pt, Zn dust,..), so what if one added Na2CO3 to the ammonia and H2O2 electrolyte? Current flow should increase and any heat generated could also move the reaction forward.

On a historical note, Watt's a long time ago did mention, per my notes, that aqueous ammonia is decomposed in the presence of finely divided Cu (note, the updated use of Zn dust, and as of 2010, Zn/TiO2 per "Comparative Study on Low Temperature Selective Catalytic Oxidation of Ammonia over Transition Metals Supported on TiO2", available at http://ieeexplore.ieee.org/xpl/freeabs_all.jsp?arnumber=5576... ) and air into "nitrite of ammonia" and other gases. On Watts, see "A dictionary of chemistry", Volume 4, page 71, at: http://books.google.com/books?q=ammonia&id=r4zPAAAAMAAJ&...
where Platinum-black is described as having the ability to quickly convert a mixture of NH3 and air into a cloud of "nitrite of ammonia". Similarly, fine copper shaken in ammoniacal air is said to produce a rapid oxidation of both the ammonia and copper (into "nitrite of copper'). This implies to me that both nitrites are formed as side reactions to the presence of NO and NO2 gases (and NH3 per equation [1] above in the formation of NH4NO2).

Caution: Any NH4NO2 formed is acutely toxic, and has some explosive issues as well (especially when dry, a high explosive with thermal and shock sensitivity). The compound has about a 2 hour half life, so don't bother to store it over night (unless in an alkali environment where the decomposition is slowed) especially in a sealed container as nitrogen gas is a decomposition product.


[Edited on 26-2-2012 by AJKOER]
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[*] posted on 7-10-2012 at 08:34


As a follow-up on the above comments on the formation of NH4NO2 (from seemingly household products like Na2CO3.H2O2 and ammonia), the reaction with NaI and Acetic acid may be able to replace the use of H2SO4/FeSO4 noted above (see "Effects of reducing reagents and temperature on conversion of nitrite and nitrate to nitric oxide and detection of NO by chemiluminescence", by Fan Yang, Eric Troncy, Martin Francœur, Bernard Vinet, Patrick Vinay, Guy Czaika and Gilbert Blaise, where the full paper is available free at http://www.clinchem.org/content/43/4/657.full ). The creation of Nitric oxide is apparently best (see Figure 2) at around 55 C.

Speculated reactions:

HOAc + NaI <--> HI + NaOAc

NH4NO2 + HI --> NH4I + HNO2

And, in anything other than very dilute, cold solutions, Nitrous acid rapidly decomposes:

2 HNO2 → NO2 + NO + H2O

and, when exposed to oxygen, NO is converted into nitrogen dioxide:

2 NO + O2 → 2 NO2

where the NO2 can be employed to form nitrates. However, the cost of any nitrate prepared by this synthesis is most likely many times more expensive than any nitrate source that can be directly purchased.


[Edited on 7-10-2012 by AJKOER]
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[*] posted on 29-3-2013 at 16:27


OK back on topic, here is an interesting account on the electrolysis of aqueous ammonia in the presence of NaOH and, separately, Cu(OH)2 from an old (1905) report (see page 242 at http://books.google.com/books?pg=PA242&lpg=PA242&dq=... from Journal Chemical Society, London, Volume 88, Part 2), to quote:

"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."

Apparently replacing NaOH with Cu(OH)2 favors the formation of nitrate over nitrites, and increasing the ammonia concentration raises the yield. Caution: product could include some copper ammonium nitrate, see discussion at http://www.pyrosociety.org.uk/forum/topic/3303-electrolysis-... . This experiment may be inherently dangerous as the author states "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" together with the observed formation of N2. From this I suspect the presence of NH4NO2 (decomposing to form nitrogen), which is inherently unstable (explosive) as the pH is lowered, which could be particular problematic in the presence also of any copper ammonium nitrate.

On the surface IMHO, this appears to be a simple, educational and safe experiment, but upon adding NaOH and/or Cu(OH)2 to the aqueous ammonia, things apparently could go very wrong, especially if one attempts to recover the dry salts.

[Edited on 30-3-2013 by AJKOER]
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[*] posted on 29-5-2013 at 17:07


Here is a more recent 1962 study on NH4NO2 formation examining various underlying theoretical models of the reaction called "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... the author cites a rate for Cu dissolution as a function of available O2 and NH3.

Some of the underlying reactions cited by the authors include:

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 this thread, an important side reaction:

2 NH3 (aq) + 3 O2 + [Cu(NH3)4](OH)2 --> [Cu(NH3)4](NO2)2 + 4 H2O

Now, I actually performed the above reaction replacing atmospheric oxygen with some dilute H2O2 to speed things up. To my surprise, Copper pennies (my Cu source) became readily covered with O2 in agreement with a cathodic reduction reaction of oxygen at the copper's surface per the author's electrochemical dissolution model. The reaction is also apparently exothermic as the solutions became warmer. Within an hour, a dark blue was apparent. In 8 hours, a different lighter shade of blue was apparent that is characteristic of the usual cupric salts. Expected products could include tetraamminediaquacopper(II) dihydroxide, [Cu(NH3)4(H2O)2](OH)2, as well a monohydroxide, tetraamminecopper(II) nitrite and also the nitrate. The important side reactions forms NH4NO2, which was somewhat apparent by more excess gas formation than I suspected (do not used a sealed vessel) with the formation of both O2 and N2 (via a nitrite decomposition reaction).

Caution: The presence of Copper Ammonium nitrite and/or Ammonium nitrite may present a potential spontaneous nitrogen gas decomposition issue, which are more likely in slightly acidic or concentrated solutions. I would also be concerned on heating an acidified form of the solution just prepared due 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... ).
--------------------------------------------------

Here is a less authoritative 2011 study ("Copper-Mediated Non-Enzymatic Formation of Nitrite from Ammonia and Hydrogen peroxide at Alkaline pH" ) that is pertinent relating to nitrite formation noted above (please see http://www.google.com/url?sa=t&rct=j&q=reaction%20of%20nh3%2Ch2o2%20and%20cu&source=web&cd=4&ved=0CDwQFjAD&url=http%3A%2F%2Fsp hinxsai.com%2Fvol3.no2%2Fchem%2Fchempdf%2FCT%3D23(646-656)AJ11.pdf&ei=iS-mUfCNN4nr0gGYw4D4BA&usg=AFQjCNFaObAi5_3NNOdt8e1DiRoiHzg9bg&bvm=bv .47008514,d.dmQ ). To quote:

"Hydrogen peroxide with lowest recorded redox
potential of - 0.68 V compared to that of Cu++ / Cu+, +
0.15 V15 acts as a strong reducing agent particularly in
presence of hydroxide ions [13], [18] to donate electrons to
copper (II) forming copper (I) oxide,

H2O2 + 2 OH- → 2 H2O + O2 + 2 e- (1)
2 Cu++ + 2 e- + H2O2 → Cu2 O + H2O (2)

Reddish-yellow cuprous oxide is rendered colorless in
presence of sufficient ammonia to form
diamminecopper (I) [15],

Cu2 O + 2 NH4OH → 2 [Cu (NH3)2] OH + H2O (3)

[ not balanced, corrected per ajkoer:
Cu2O + 4 NH3 + H2O → 2 [Cu(NH3)2]OH (3)]

Diamminecopper (I), generated from reduction of
copper (II) or added exogenously facilitates oxidation
of ammonia, a reducing agent [14], by hydrogen
peroxide,

...[Catalyst].....Cu (NH3)2]OH.........................
NH3 + 3 H2O2 -----------------> HNO2 + 5 H2O (4)

[ not balanced, corrected by ajkoer:
NH3 + 3 H2O2 -----------------> HNO2 + 4 H2O (4)]

Further studies are required to elucidate the actual role
of diamminecopper (I) in the reaction; whether it is
converted to tetramminecopper (II), or undergoes a
reversible changes during the process."

With additional ammonia, the reaction with nitrous acid proceeds as follows:

HNO2 + NH3•H2O --> NH4NO2 + H2O

Interesting observations by the author's non-electrochemical experiment includes "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 30-5-2013 by AJKOER]
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[*] posted on 29-5-2013 at 18:06


i one time put silver annode and cathode in ammonium hydroxide and the ammonia turned dark blue and had a bunch of silver crystals at the bottom.i sealed the electrodes with epoxy and air tightened the whole thing and remember that it lost its color after a couple of days.the silver dissolves very quickly and i did this back when i didnt understand where the nitrogen came from when one made silver nitrate.
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[*] posted on 29-5-2013 at 22:46


The dark blue color is due to the presence of copper. Normally, silver is alloyed with copper to make its mechanical properties better. The copper is oxidized at the anode and gives copper(II) ions, which with ammonia form the royal blue [Cu(NH3)4](2+) complex.



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