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Author: Subject: Silver / Copper Nitride
AndersHoveland
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[*] posted on 20-5-2011 at 14:49


Some pictures:

Aqueous ammonia (NH4OH) is added to silver nitrate (AgNO3). The result is a light tan precipitate (silver nitride):
http://jchemed.chem.wisc.edu/JCESoft/CCA/CCA4/STHTM/AGNO/AGN...

Aqueous ammonia is added to lead(II) nitrate. The result is a white precipitate:
http://jchemed.chem.wisc.edu/JCESoft/CCA/CCA4/STHTM/PBNO/PBN...

This does not necessarily mean that the product is lead(II) nitride because a similar reaction is also observed using zinc nitrate, which would not be expected to form a nitride.
http://jchemed.chem.wisc.edu/JCESoft/CCA/CCA4/STHTM/ZNNO/ZNN...

Lead(II) salts react with aqueous ammonia to precipitate a white basic salt, Pb2O(NO3)2, rather than the expected lead(II) hydroxide, according to the reaction:

Pb[+2] + (2)NH3(aq) + (3)H2O + 2NO3[-] <==> Pb2O(NO3)2 + H2O + (2)NH4[+]

This basic lead salt is insoluble in excess ammonia.

Mercury Nitride

Mercury(II) nitride forms a chocolate colored powder, which is slowly decomposed by water. The dry nitride tarnishes in air. Mercury(II) nitride is very explosive, and must be handled with extreme care. It detonates violently, yielding a white flame with a bluish purple border, when heated. The salt is so sensitive that it can be detonated by rubbing it with a glass stir rod. It is formed by reacting HgO with NH4OH, initially at 10°C, but thereafter heating the reaction.

Mercury(I) nitrate also reacts with ammonia to form a white precipitate, but I am unsure if this is mercury(I) nitride, or simply a less soluble ammonia complex of mercury(I) nitrate, both of which would be explosive.
http://jchemed.chem.wisc.edu/JCESoft/CCA/CCA4/STHTM/HGNO/HGN...

The reaction between mercury(I) nitrate and aqueous ammonia produces a mixture of a white basic amido salt and metallic mercury, both of which precipitate out.

(2)Hg2[+2] + (4)NH3(aq) + NO3[-] + H2O --> (2)Hg + Hg2ONH2NO3 + (3)NH4[+]

Interestingly, another site shows pictures of this reaction giving brownish black colored reaction products, in contrast to the white colored products seen in the other site.
http://www.public.asu.edu/~jpbirk/qual/qualanal/mercury.html
It is possible that different reactant ratios could give different products with correspondingly completely different observable color changes in reaction. Perhaps someone here would like to investigate this?

Another explanation for the precipitate is that the ammonia, because it is a base, causes mercury(I) to precipitate out as a double salt. Even if dissolved in plain water, mercury(I) nitrate is somewhat acidic due to its slow reaction with water:

Hg2(NO3)2 + H2O --> Hg2(NO3)(OH) + HNO3

The Hg2(NO3)(OH) separates out as a yellow precipitate. In the presence of ammonia, it may be possible that some other double salt, with an unknown structure, would form.

Higher nitrides of lead and silver?

I would think one would want to try reacting lead dioxide rather than lead(II) oxide. Despite being very coordinating, the Pb+2 ion does not seem to be much of an oxidizing ion like Ag+1 or Au+3. Lead dioxide dissolves in strong bases. Someone here might try reacting such a solution (of lead dioxide which has been dissolved in aqueous sodium hydroxide) with NH4OH, to see if an explosive lead(IV) nitride can be obtained.

Ag2O4 (silver (I,III) oxide) is also known to dissolve in strong bases. These solutions might also react with NH4OH to produce a different form of silver nitride, with a higher proportion of nitrogen. Ag3N2 ?

Notes about preparation of precursors

(1) lead dioxide may be prepared by oxidizing lead(II) nitrate with potassium ferrate, then adding dilute nitric acid

(2) silver(I,III) oxide can be easily prepared from AgNO3 and sodium persulfate. Here is a video:
http://www.youtube.com/watch?v=1_a81M9p2so

[Edited on 21-5-2011 by AndersHoveland]
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AJKOER
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[*] posted on 21-5-2011 at 13:31


Some minor corrections on some good research.

2 NH4OH + 2 AgNO3 --> Ag2O(a tan to black solid) + H2O + 2 NH4NO3

Ag2O + 4 NH4OH --> 2 [Ag(NH3)2]OH + 3 H2O

So, with excess ammonia the silver oxide dissolves forming di-amine silver hydroxide.

The latter discomposes on standing forming Ag3N (also some AgNH and AgNH2).


ON THE TOPIC OF LEAD NITRIDE Pb2N3

The compound exists but is (per Bretherick's "Handbook of reactive chemical hazards", page 1884), "Very unstable, it decomposes explosively on vacuum degassing."
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AndersHoveland
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[*] posted on 21-5-2011 at 14:46


Quote: Originally posted by AJKOER  

2 NH4OH + 2 AgNO3 --> Ag2O(a tan to black solid) + H2O + 2 NH4NO3

Ag2O + 4 NH4OH --> 2 [Ag(NH3)2]OH + 3 H2O

So, with excess ammonia...


This may potentially explain the two different results in the two sites that showed pictures of mercury(I) nitrate reacting with ammonia. With excess ammonia, the reaction would turn black, instead of the white precipitate.

When you mentioned lead nitride as "Pb2N3", did you perhaps mean Pb3N2 ?

While the Hg[+2] ion can easily form nitrides, the same is not true of the Pb[+2] ion. So it seems that the more oxidizing the ion is, the easier it is to form nitrides. Of course, if the ion is too oxidizing, it would only simply oxidize the ammonia, liberating nitrogen gas. As mercury(II) nitride readily forms, but not mercury(I) nitride, one would be inclined to think that lead(IV) nitride, if it exists, may be more chemically/thermally stable than lead(II) nitride.

"Lead dioxide... oxidizes ammonia to nitric acid, with the simultaneous formation of ammonium nitrate. It oxidizes manganese salt (free from chlorine) in the presence of nitric acid to permanganate." Encyclopaedia Britannica. Eleventh Edition. Volume XVI. edited by Hugh Chisholm. p318 (year 1911)

I do not know if lead dioxide which has been dissolved alkaline (forming the hydroxyplumbate ion, Pb(OH)6[-2] ) can oxidize ammonia, or ammonium hydroxide reacts with solid PbO2. The article is not clear about whether lead dioxide needs to be acidified to attack ammonia. While lead dioxide is an oxidizer even while alkaline, oxidizing Cr[+3] to CrO4[-2], the salt ammonium chromate does exist, so this does not imply that it would necessarily oxidize ammonia.

[Edited on 21-5-2011 by AndersHoveland]
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[*] posted on 15-7-2011 at 19:23


Then,silver or copper azide salt,more violence.
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AJKOER
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[*] posted on 29-1-2012 at 16:59






Quote: Originally posted by Formatik  
"Fulminating copper": just found this. They dissolve pure Cu in dilute HNO3 (might need some heat) and then pour in "liquid Ammonia" (likley just aqueous NH3) to get a precipitate, they dry it (with heat!) and say rubbing it causes a loud explosion. 2 grains (129.6 mg) of it over a fire for a bit is said to produce a loud explosion.


I suspect this source is referring to Copper Ammonium Nitrate Cu(NH3)4(NO3)2.

First, per Wiki: "Schweizer's reagent is the chemical complex tetraamminediaquacopper dihydroxide, [Cu(NH3)4(H2O)2](OH)2. It is prepared by precipitating copper(II) hydroxide from an aqueous solution of copper sulfate using sodium hydroxide or ammonia, then dissolving the precipitate in a solution of ammonia."

My speculation that in the last step replacing ammonia with NH4NO3 and warming forms Copper Ammonium Nitrate (or CAN), formula Cu(NH3)4(NO3)2. Speculation is based on how other Copper Ammonium salts are prepared.

http://www.orgsyn.org/orgsyn/prep.asp?prep=cv2p0142

Confusion because CAN is capable of (and has been) used as a primer in place of mercury fulminate. See:

http://books.google.com/books?id=dHbHS5GhCN4C&pg=PT70&am...

I am not recommending anyone follows this or Formatik's source synthesis, but if they so elect, please read the MSDS for CAN to avoid potential injury.


[Edited on 30-1-2012 by AJKOER]
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[*] posted on 29-1-2012 at 17:20


Back when I took chemistry lab in college, many years ago, I made a solution of (ammoniacal silver) this stuff and put it in a 50 ml ground glass reaction flask for safekeeping, along with my complete set of rented glassware for that semester, and synthesis assignment. I thought it would be safe because I didn't let it get dry. Well , ... I came back on my lab day and opened the locked drawer to find a blackened mess with many dollars worth of expensive glass blasted to bits. I realized immediately what had happened, and closed the drawer, with a few crackling noises. No sense in letting everyone know what a foolish thing I had done. Later I came back and cleaned up and totaled the losses. It could have been worse. A few hundred dollars was a big expense to college student in those days.

This stuff serves no useful purpose, so don't risk your fingers and eyes screwing with it.
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