White Yeti
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One the subject of alkali metal production via oxalate salt intermediate.
I don't know if this idea is brilliant or absurd, so bare with me.
Heating iron oxalate yields pyrophoric iron dust and carbon dioxide:
FeC2O4 + heat ----> Fe(s) + 2CO2
I was wondering if this could be done with say.... calcium oxalate, or even better, potassium oxalate (to yield alkali, and alkali earth metals
respectively).
It seems like it would work if all the water was boiled off before the reaction started. This is no problem because the reaction temperature is much
higher than that of the boiling point of water. There's one problem that remains. What is the reaction temperature? Does it exceed the boiling point
of the alkali metal in question? Surely I could install a condenser, but this is more of a safely issue, I don't want to be exposed to gaseous bases
and such.
One last question, does potassium react with carbon dioxide gas at high temperatures?
This idea seems so simple that I wonder why no one has thought of this before. There must be a catch somewhere.
Any ideas?
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bbartlog
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The pyrolysis of salts of the transition metals gives entirely different results from that of alkali salts with the same cations. For example, heating
calcium acetate yields (mostly) acetone and calcium carbonate, while doing the same with copper acetate will give you mostly acetic acid and
pyrophoric copper. This is not to say that you can assume all of the transition metals or even all the alkalis will behave exactly the same...
There are a couple of pages on the thermal decomposition of oxalates in 'The pyrolysis of carbon compounds' by Charles Dewitt Hurd, freely available
at hathitrust.org. Calcium oxalate decomposes into calcium carbonate, carbon dioxide, carbon monoxide, and carbon at >440C. I would expect other
alkali metals to also yield the carbonate with some variation in the details.
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Melgar
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Yeah, at best you'd get calcium oxide.
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AndersHoveland
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Actually, it is possible to prepare sodium metal by cautiously heating sodium azide in the absence of oxygen.
(2)NaN3 --> (2)Na + (3)N2
Similarly, the explosive decomposition of copper azide also results in the separation of the constituent elements, but this reaction happens for very
different reasons.
Cu(N3)2 --> Cu + (3)N2
But the same reaction with iron (which is dangerous) will result in iron nitride.
(3)Fe(N3)2 --> Fe3N2 + (8)N2
The iron nitride can be decomposed to elemental iron and nitrogen gas above 800°C.
Fe3N2 --> (3)Fe + N2
The decomposition of calcium azide is similar to that of iron.
Ca(N3)2 decomposes above 110degC, explosively so over 140degC. The Ca3N2 that forms only decomposes at 1600 °C, at which point the elemental calcium
simultaneously vaporizes out with the nitrogen.
Creative Way to Make Elemental Potassium?
An idea for chemical preparation of elemental potassium, which does not require electric current. It would be impractical, but very creative. Not sure
if all the reactions would work.
Ca3N2 + (6)KCl --> (3)CaCl2 + (3)K2 + N2
Distilling calcium nitride with potassium chloride in with steel-walled distillation may cause potassium to boil out. This proposed reaction would
make use of Le Chatelier's principle. Although potassium boils at 759°C, it is possible that molten potassium could be produced below this
temperature. Although lithium can burn in nitrogen, both sodium and potassium nitrides are very unstable. Sodium nitride decomposes into elemental
sodium, giving off nitrogen gas, at only 87°C.
(6)CaCl2 + Ti3N4 --> (2)Ca3N2 + (3)TiCl4
The titanium nitride (m.p. 2930°C) would be crushed into a fine powder and distilled under intense heat with calcium chloride. Titanium tetrachloride
(TiCl4) is a liquid which boils at only 137 °C.
(3)TiI4 + (16)NH3 --> Ti3N4 + (12)NH4I
I think titanium tetraiodide (b.p. 377 °C) could be reacted with anhydrous ammonia gas to form titanium nitride and ammonium iodide. I am not sure if
the NH3 could be bubbled into molten TiI4, or if the TiI4 would need to be in the vapor phase, with the intense heat required for the reaction. The
reaction would be expected to procede because TiI4 is very acidic, and because the titanium-nitrogen bonds are stronger than titanium-iodide.
Wikipedia claims that TiCl4 "with ammonia, titanium nitride is formed"; this is not surprising since TiCl4 reacts with water to form titanium dioxide
and hydrogen chloride.
Titanium tetraiodide melts at 150 °C. It can be prepared from easily obtainable materials:
(3) TiO2 + (4) AlI3 --> (3)TiI4 + (2)Al2O3
What do you think? Opinions?
[Edited on 27-9-2011 by AndersHoveland]
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White Yeti
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Good stuff Anders. I was originally planning on attempting to make some potassium graphite (awesome reducing agent). Unfortunately I don't have a
stockpile of potassium metal at my disposal and the stuff is extremely expensive because of the shipping costs. I was looking for an economically
feasible way to make small amounts of potassium metal.
The reaction is definitely creative, talk about thinking outside the box. But it seems just as suicidal as passing hundreds of amps through molten
potassium hydroxide :\ and perhaps even more expensive.
For my purposes I was wondering if I could replace the potassium with another metal that is still a good reducing agent. The first one that comes to
mind is zinc. I can easily melt it on my hot-plate and inserting some graphite rods into molten zinc shouldn't be a problem at all. Plus, I have about
half a kilo of zinc at my disposal. Aluminium is also another option if zinc doesn't work.
I didn't find any documentation on "zinc graphite" so I have no clue if this will work. I'll try it out anyway.
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bbartlog
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You can make mossy zinc via electrolysis (of a zincate solution if I remember right); it would have a large surface area. Not sure what inserting a
graphite rod into molten zinc will get you - I would expect to get graphite coated with a layer of zinc, but it has a high enough surface tension that
I hardly think you will get a large surface area, i.e. it will either fill all pores or just coat the outside, depending on the surface interaction.
Anyway, while zinc is a fine reducing agent it is not as powerful as potassium, so it depends as always on what you're actually trying to do.
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