APO
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Strontium Fulminate Synthesis
I just got some anyhydrous strontium fulminate, seeing as it is a
strong nitrating agent(Sr(NO3)2), I was wondering if mixed with pure ethanol, would it create strontium fulminate?
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AndersHoveland
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Sr(NO3)2 is not a nitrating agent, at least not by itself. Anhydrous magnesium and copper nitrates are, however (it is not possible make the
anhydrous form of either just by heating their hydrates)
Even if you added concentrated HNO3 to the strontium nitrate, fulminate would still not form. For the formation of fulminate, there usually needs to
be another metal ion that will bind with the fulminate and take it out of solution, before it further reacts.
[Edited on 30-12-2012 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.
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APO
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So if magensium is a strong nitrating agent why isn't strontium nitrate? They're in the same family. Also I don't get how adding HNO3 would not work
either, if you add strontium to fuming nitric acid and let it dissolve, then add it to ethanol would that not make strontium fulminate? That method
works for mercury and silver, so why not other metals? I really don't see why strontium nitrate and ethanol won't make the fulminate. Also I'm not
asking how to make it anyhydrous I already have anyhydrous strontium nitrate.
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AJKOER
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Here is an available online reference and more background material on strontium fulminate and related salt preparation that be be of assistance:
"Report 1833 FULMINATES: A LITERATURE SURVEY", 1965, from the U. S. Army Engineer Research and Development Laboratories.
To quote from page 4 at http://www.dtic.mil/dtic/tr/fulltext/u2/625397.pdf :
"III. FULMINATES
1. Synthesis. The preparation of mercuric fulminate, by the addition to ethyl alcohol of a solution of mercury in nitric acid, was described by Howard
in 1800. In 1802, silver fulminate was obtained by Brugnatelli by the addition of alcohol and then nitric acid to powdered silver nitrate (1). Methods
for preparing colloidal aggregates (40) and small crystals (41) of silver fulminate are described by Taylor and others. The cadmium, thallous, and
cuprous salts of fulminic acid were obtained by reaction of either mercuric or silver fulminate with the amalgam of the particular metal, in dry
methanol and a hydrogen atmosphere, the products being precipitated out with ether (42,43). Cd fulminate is stable when dry but readily decomposes in
water and is sensitive to heat or shock. The Cu(I) salt is insoluble in water and also highly explosive. Thallous fulminate is sensitive to moisture
and light and susceptible to heat and shock but does not explode violently. Langhans (44) reviewed tho preparation of fulminates (and also noted the
effect of moist mercury fulminate upon metals). The fulminates of sodium, potassium, calcium, strontium, and barium were similarly prepared by
reaction of the appropriate amalgam with mercuric fulminate, in BaO-distilled methanol (42,45). Shaking of the reaction medium was continued until Hg
was no longer detected with SnCl2 . Temperatures from -5 to -15 C were used to avoid polymerization. The solutions were filtered into cold, dry ether
in an oxygen-free atmosphere. The alkaline earth fulminates separated with 1 mole of methanol, which could not be driven off without decomposition of
the salts. Both the alkali and the alkaline earth fulminates are unstable to CO2 and moisture."
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woelen
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Quote: Originally posted by APO  | | So if magensium is a strong nitrating agent why isn't strontium nitrate? They're in the same family. Also I don't get how adding HNO3 would not work
either, if you add strontium to fuming nitric acid and let it dissolve, then add it to ethanol would that not make strontium fulminate? That method
works for mercury and silver, so why not other metals? I really don't see why strontium nitrate and ethanol won't make the fulminate.
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The reason why some metals do make fulminates and others don't is simply because not all metals have the
same properties. Silver tends to form (nearly) covalent bonds with many anionic species, while alkali metals and the heavier earth alkali metals only
form purely ionic compounds. The same is true for mercury. These covalent compounds have very different properties compared to their ionic
counterparts. Usually these covalent compounds have complicated and frequently polymeric properties which make them insoluble in most solvents,
allowing them to precipitate from such solvents.
Silver, mercury(II) and to a lesser extent copper(I) and lead(II) are known well for their capability to form insoluble covalent compounds with many
anions, such as fulminate, azide, acetylide, which easily are prepared simply by forming the corresponding ionic species or even the hydrogen compound
of the anion (e.g. silver acetylide can be made by simply bubbling C2H2 through an aqueous solution containing silver ions).
I indeed do not expect that any strontium fulminate will be formed when Sr(NO3)2 is added to alcohol. No reaction will occur at all. On addition of
fuming nitric acid I only expect a reaction between the acid and the ethanol, the Sr(NO3)2 will not react, or, when the material inflames, it will
completely oxidize the ethanol in a violent reaction.
[Edited on 2-1-13 by woelen]
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AndersHoveland
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| Quote: Originally posted by APO | | So if magensium is a strong nitrating agent why isn't strontium nitrate? They're in the same family. |
Yes, the elements in both the alkali and akaline earth families, repectively, usually have almost identical chemistry, the main difference being
increasing reactivity. However there are a small handful of instances where there are significant chemical differences.
To name a few:
Potassium forms a superoxide, whereas sodium does not. There have been instances where the surface of pieces of metallic potassium that have been
exposed to air in storage have exploded or spontaneously ignited as soon as the metal was handled or cut. My chemistry professor told us that, after
having read of this, he was reluctant to use potassium in his demonstrations any longer, as he had done in the past, and despite the beautiful purple
flame from potassium he was now using sodium instead.
Because of the smaller atomic radii, both lithium and magnesium ions are actually somewhat acidic when they are not complexed to water molecules. The
energy it takes to ionize the second electron off of magnesium is surprisingly high. The only reason Mg+2 ions even exist in anhydrous MgSO4 is
because of the lattice energy of the crystal structure. In the absence of water, beryllium has some bizarre chemistry, more in common with boron
really. Beryllium is not often encountered, being very rare, expensive, and poisonous.
One interesting bit of trivia is that, although sodium is "more reactive" than lithium, a molar equivalent of lithium actually releases more energy
reacting with water than a molar equivalent of sodium. This is because of the additional energy released from the solvation of lithium ions, which is
considerable. The reaction may only seem more violent with sodium because sodium has a lower melting point, and this greatly increases reaction rate.
But of course, in the absence of water, sodium will still displace lithium from its molten chloride salts, so sodium is indeed "more reactive".
However, lithium can burn in nitrogen, whereas molten sodium will not react with nitrogen under any conditions. The reasons for this seemingly
counterintuitive difference in chemistry is explained elsewhere in this forum.
There are also very significant differences between the different metal azides in the same family:
| Quote: Originally posted by AndersHoveland |
Lithium azide has only moderate explosive properties. Decomposition of lithium azide only forms lithium nitride and nitrogen. It can survive hammer
blows without detonation
Sodium azide is not itself an explosive. The main reason for this is rather simple- the decomposition proceeds according to the following equation:
2 NaN3 --> 2 Na + 3 N2
The reduction of sodium ions to elemental sodium is not particularly favorable, and this is one of the main reasons sodium azide is not explosive.
Sodium nitride does not form because Na3N is not very stable (sodium nitride decomposes into its elements at only 87 °C ).
Calcium azide begins to thermally decompose above 110 °C, and explodes at 158°, it is more explosive than either strontium or barium azide.
Barium azide is a sensitive explosive, with a drop height value of 10cm. It appears to be relatively insensitive to impact but highly sensitive to
friction.
H. Ficheroulle, Mem. des Poudres. 33, 7 (1956)
The temperature at which barium azide explodes is apparently highly variable, values have been reported between 152° to 221°C.
The enthalpy of formation for barium azide from its elements is actually slightly negative, -5.3 kcal/mole. The formation of barium nitride is very
favorable, the compound having an enthalpy of formation of -89.9 kcal/mole. "Nitrogen Burning of Metals", G. Petrov, Combustion, Explosion, and
Shock Waves, Volume 11, Number 3, 309-312
The decomposition products from the explosion of barium azide leave behind a significant quantity of metallic barium, whereas the explosion of calcium
azide leaves only calcium nitride.
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The general rule seems to be that the "less reactive" alkali/alkaline earth elements tend to form more stable compounds with nitrogen. This has to do
with atomic size, crystal lattice energy, and ability to form covalent bonds, since true nitride ions, N-3, are not very stable. The
heavier elements in these families prefer ionic bonds.
There are also a few obscure specific reactions in organic chemistry where there are significant differences in the reactivity of potassium hydride
compared to sodium hydride.
("Potassium hydride, highly active new hydride reagent"
Charles Allan Brown, J. Org. Chem., 1974, 39 (26), pp 3913–3918 )
So in conclusion, in most situations the elements in each of these families are completely interchangeable, but there are a few exceptions, and this
should give you a few examples.
I am always interested in the obscure differences in chemistry between elements in the same column.
[Edited on 2-1-2013 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.
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12AX7
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| Quote: Originally posted by AndersHoveland |
Potassium forms a superoxide, whereas sodium does not. There have been instances where the surface of pieces of metallic potassium that have been
exposed to air in storage have exploded or spontaneously ignited as soon as the metal was handled or cut. My chemistry professor told us that, after
having read of this, he was reluctant to use potassium in his demonstrations any longer, as he had done in the past, and despite the beautiful purple
flame from potassium he was now using sodium instead.
|
Indeed, this has been known for about as long as potassium metal itself has been known -- in one of the Faraday lectures (The Chemical History of a
Candle), Faraday was demonstrating that carbon dioxide contains oxygen, by combusting potassium in it. The potassium has to be warmed first, to
encourage it to burn:
| Quote: | | [In the preliminary process of heating, the potassium exploded.] Sometimes we get an awkward piece of potassium that explodes, or something like it,
when it burns. |
http://www.gutenberg.org/cache/epub/14474/pg14474.html Lecture VI
Very likely that "awkwardness" was due to superoxides!
Tim
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