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APO
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I have an idea for makeing the caesium and rubidium a bit easier, maybe I could add stochiometric amounts of caesium/rubidium hydroxide to magnesium
in a flask filled with mineral oil ,with a condenser,presure equalizing seperatory funnel and vaccum adapter. I connect it to a vaccum pump, then heat
it to 200c were the caesium/rubidium hydroxide and magnesium would be catalysed by addition of tertiary butanol or tertiary pentanol while in the
inert enviorment of mineral oil, the metal produced would rest on the bottom so that I could coalesce them later, and of course the hydrogen would be
sucked through the presure equalizing seperatory funnel and out of the vaccum adapter. I have a 100ml erlenmyer flask, a 300mm vigruex column, a 25ml
presure equalizing seperatory funnel, and a 90 degree vaccum adapter and a hotplate. Now I just need a vaccum pump. Does this sound like a plausible
method for making rubidium or caesium?
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elementcollector1
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According to the original patent, it's supposed to work for cesium. No idea on rubidium, but it seems that it would follow.
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APO
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Awesome.
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elementcollector1
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Incidentally, how do you intend to collect the cesium after the reaction (hopefully) suceeds? Can't take it out into the air. Could pipette it up...
maybe. I hope all your glassware is ground glass; you're going to need near-perfect vacuum seals for this metal.
I do wish some of the more experienced members in this subject were around to try it out, but...
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blogfast25
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Quote: Originally posted by APO  | | I have an idea for makeing the caesium and rubidium a bit easier, maybe I could add stochiometric amounts of caesium/rubidium hydroxide to magnesium
in a flask filled with mineral oil ,with a condenser,presure equalizing seperatory funnel and vaccum adapter. I connect it to a vaccum pump, then heat
it to 200c were the caesium/rubidium hydroxide and magnesium would be catalysed by addition of tertiary butanol or tertiary pentanol while in the
inert enviorment of mineral oil, the metal produced would rest on the bottom so that I could coalesce them later, and of course the hydrogen would be
sucked through the presure equalizing seperatory funnel and out of the vaccum adapter. I have a 100ml erlenmyer flask, a 300mm vigruex column, a 25ml
presure equalizing seperatory funnel, and a 90 degree vaccum adapter and a hotplate. Now I just need a vaccum pump. Does this sound like a plausible
method for making rubidium or caesium? |
You mean you got the 'idea' from the relevant sticky thread?
Vacuum will also reduce the BP of your solvent which will come over too. Not necessarily a problem but your catalyst will certainly come over too and
react immediately with any distilled Cs. And w/o catalyst in the reactor flask, NO REACTION!
In short: forget it!
[Edited on 6-2-2013 by blogfast25]
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elementcollector1
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I think he means to leave the cesium in there, and somehow collect it later, not distill it immediately. Although the condenser confuse me. Did you
mean a reflux condenser?
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APO
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I mean I would prepare it under oil in the apparatus I described with magnesium powder and hydroxide, and then once done I would pour the mixture into
a beaker and drop the caesium into oil for storage. Mineral oil barley refluxes at 200c so I would use a vigruex column on top of a flask as a
condenser,and on top of that I would have the rest of the apparatus.
So in short I would use a vacuum adapter on top of a pressure equalizing seperatory funnel, on top of a vigruex column, on top of a Erlenmeyer flask
being heated by a hotplate. Then then magnesium oxide, hydrogen, and the metal of choice would form under mineral oil. The magnesium oxide and metal
would rest on the bottom under the mineral oil, while the hydrogen diffuses out.
I'm confident it will work so I will post my results when finished.
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kristofvagyok
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| Quote: Originally posted by APO | So in short I would use a vacuum adapter on top of a pressure equalizing seperatory funnel, on top of a vigruex column, on top of a Erlenmeyer flask
being heated by a hotplate. Then then magnesium oxide, hydrogen, and the metal of choice would form under mineral oil. The magnesium oxide and metal
would rest on the bottom under the mineral oil, while the hydrogen diffuses out.
I'm confident it will work so I will post my results when finished. |
This writeup is not worth even for trying it. If you would like to try it than I would suggest to get a really good bed at the local hospital.
Why?
1: we do not heat Erlenmeyer flask on hotplate, because it could easily break. Rule no.1: always use round bottomed flasks.
2: if you do not have a dry inert gas (e.g.: argon) than imagine that what would happen if you would swith off the vacuum? If you have no idea than I
would tell a trick: oxygen is much better soluble in mineral oil than usually expected.
3: why is is good to make some cesium under a lot magnesium oxide/magnesium and Cs-oxide if you can't do anything with it? It will slowly turn back to
Cs-oxide even if the bottle is locked.
Get a RBF, some argon or nitrogen and make a proper setup, working as you described will only end up with a really nice accident.
I have a blog where I post my pictures from my work: http://labphoto.tumblr.com/
-Pictures from chemistry, check it out(:
"You can’t become a chemist and expect to live forever." |
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blogfast25
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Yeah, I look forward to that!
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blogfast25
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Oh but we do, we do! All the time, in very harsh conditions sometimes, not once had a problem.
In a TiO2 plant I worked at we used to dissolve CALCINED TiO2 in boiling conc. H2SO4 for at least 1/2 hour (for Ti assay), in Erlenmeyer, on top heat
(electric plate). Not once did one crack or smash. I've done this also in my home lab.
All the experimenters in the sticky K-thread used Erlenmeyers.
Decent quality boroglass suffices for this, whether Erlen or RBF.
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APO
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Hold on kristofvagyok! Heating a RBF without a special mantel is much more dangerous than using an Erlenmyer flask because the heat will be focused at
a point plus my glass is rated for 500 celcius and is heavy walled, I clearly stated before that I will use a vacuum or inert atmosphere, the only
possible point you may have is a slight reaction between the metal and the magensium oxide, I'm pretty sure they don't react under these conditions.
Caesium and Rubidium don't even react with lithium under these conditions, I supect that magnesium oxide would behave pretty inert, if anyone knows
for sure please tell me.
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phlogiston
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Very cool, Valentine!
Would this approach also work for preparing small quantities of the alkaline earth metals? I'd be interested specifically in preparing barium and
strontium metal.
Can the alkaline earth azides be decomposed in a controlled fashion?
The title of the following paper seems to suggest it is possible, but I can only access the first page of it:
Die Zersetzung der Alkali- und Erdalkali-azide im Hochvakuum zur Reindarstellung von Stickstoff
Berichte der deutschen chemischen Gesellschaft Volume 49, Issue 2, pages 1742–1745, Juli–Dezember 1916
[Edited on 8-2-2013 by phlogiston]
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"If a rocket goes up, who cares where it comes down, that's not my concern said Wernher von Braun" - Tom Lehrer |
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APO
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Well, lithium azide and sodium azide are explosive at elevated tempatures, and alkali earth metals are a little less reactive than lithium, so they
may have similar characteristics, so best not to use glass for that. But a steel crucible may be suitable.
Also for making hydrazoic acid you can mix hydrochloric acid and sodium azide and distill it to obtain your final product for making azides.
[Edited on 8-2-2013 by APO]
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APO
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I think I found a way to make azides without making the acid. Check this out: http://web.mit.edu/semenko/Public/Military%20Manuals/RogueSc...
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Zan Divine
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APO, as long as you are buying a vacuum pump, why not go the easy route?
I've posted my results for Li reduction of CsCl and Li and Ca reductions of RbCl. 660 degrees C is easy to reach. You start and a few hours later you
are ready to seal Cs in an ampoule. The synthesis and isolation/purification are all in one operation.
Warm syringes and needles (say 18-20 gauge) are ideal for handling. You can easily store the metal in a glass screw-top bottle under mineral oil and
then draw oil-free beautiful metal from its center. Handling Cs and Rb is not difficult. You will want to rent an argon cylinder and pass the gas
through a drying train first for inerting receivers.
 
[Edited on 2/14/2013 by Zan Divine]
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phlogiston
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That method will only work for poorly soluble azides. The azides of rubidium, caesium, barium and strontium are all quite soluble.
-----
"If a rocket goes up, who cares where it comes down, that's not my concern said Wernher von Braun" - Tom Lehrer |
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APO
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Aw man.
[Edited on 14-2-2013 by APO]
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Zan Divine
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| Quote: Originally posted by phlogiston | Would this approach also work for preparing small quantities of the alkaline earth metals? I'd be interested specifically in preparing barium and
strontium metal.
Can the alkaline earth azides be decomposed in a controlled fashion? |
The answer is a qualified yes. There are numerous variables involved, unfortunately.
Different crystalline forms are possible, some are more prone to violent decomposition than others. Certain impurities (including, paradoxically, Ba
metal in the case of barium azide) catalyze more rapid decomposition rates, which you don't necessarily want.
The bottom line is just as you might expect. You can, and people have, thermally decomposed group II azides, but you are rolling the dice. Any
experimental setup should be designed and run with the possibility of detonation in mind.
Much is made of the relative difficulty of preparing pure group 1 metals but Ca, Ba & Sr present much more formidable targets.
One side note...if any of you prepare HN3 you had better be VERY observant of the special rules that apply to this acid. The azide is best used as
produced and not allowed to accumulate. Ground glass may trigger an explosion even when greased. Hydrazoic acid is the only lab chemical that I have
witnessed involved in a detonation and this was in the hands of a very experienced chemist.
[Edited on 2/16/2013 by Zan Divine]
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a_bab
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I would add that decomposing a group II azide si almost like heating lead azide without getting it to detonate. It can be done, but the risk is
enormous. These group II azides are simply described as "explosive".
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Formatik
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| Quote: Originally posted by phlogiston | The title of the following paper seems to suggest it is possible, but I can only access the first page of it:
Die Zersetzung der Alkali- und Erdalkali-azide im Hochvakuum zur Reindarstellung von Stickstoff
Berichte der deutschen chemischen Gesellschaft Volume 49, Issue 2, pages 1742–1745, Juli–Dezember 1916 |
The paper is attached below. Take it with a grain of salt.
Paraffin or sand bath was used to carefully heat dry small samples (0.5g) under high vacuum. Sodium, potassium, rubidium and cesium azide were
investigated. As opposed to these lithium azide was said to be pretty explosive, and so wasn't investigated. Additionally, calcium, strontium and
barium azides were looked at. In one of the heating attempts there was an explosion with calcium azide, but the author kind of swept this aside saying
it usually goes pretty peacefully, saying the salt may have been impure or the heating jumped by too much.
But it is known barium, strontium and especially calcium azide are explosives as mentioned above. The lead block values of these are listed in
Urbanski, and are beneath or approach mercury fulminate. Calcium azide is the strongest. Sodium azide gives a result of zero in the lead block test
(normally a 10g sample fired with a No. 8 blasting cap), which factors into supporting the idea that it is not an actual explosive, and so isn't
classified as an explosive.
If you still feel so inclined to heat the (earth) alkaline azides in small amounts, use a saftey shield or put a few thick cinder blocks in front of
the apparatus trajectory and take other appropriate saftey measures like a faceshield, thick gloves, clothes, etc.
But alkali azides are not entierly free of explosion risks. Although they are not considered explosives, they all have the capacity for explosion.
Rubidium azide although has shown no reaction to hammer blows, is stated to still show clear shock sensitivity and explodes when overheated (Gmelin).
Cesium azide although it can be used to make pure cesium metal, has exploded if heated above decomposition temperatures (also in Gmelin). And also
according to references in Gmelin, even potassium azide explodes if overheated. In Federoff, it states that sodium azide explodes only when heated to
a high temperature. Though from tests done, sodium azide does not propagate detonation.
Basically, nearly all what Wiki says on the matter is misleading:
| Quote: | "Azides of heavier alkali metals (excluding lithium) or alkaline earth metals are not explosive, but decompose in a more controlled way upon heating,
releasing spectroscopically-pure N2 gas.[8]"
http://en.wikipedia.org/wiki/Hydrazoic_acid |
The alkali metal compounds above lithium, can be decomposed in a more controlled manner. The alkaline earth metal compounds also, but less so. Calcium
azide in particular is more explosive than Ba or Sr azide.
For decomposition of calcium and strontium azides, in Gmelin I've read nitride is formed so that the metal formed is not pure. Thus, not all of the N2
is necessarily driven out from these compounds. The Na or Cs metal that can be formed from azides is pure though.
Concerning hydrazoic acid:
It can be made using alkali azide and dilute H2SO4, as Curtius has stated. HCl would be too volatile to use.
Conc. H2SO4 destroys hydrazoic acid as its made from azide salts, so that this acid can't be used as was also observed by Curtius (he was talking
about making the anhydrous stuff).
In a conference of the Deutschen chemischen Gesellschaft zu Berlin 25. November 1895 (Ber. 29, 771), Curtius noted:
Wässrige Lösungen des Stickstoffwasserstoffs, auch ziemlich concentrirte, scheinen sich dagegen ohne Gefahr handhaben zu lassen, wenn man nicht
eine Flamme der Flüssigkeit nähert.
Namely, even fairly concentrated solutions of hydrazoic acid are not dangeorus to handle as long as no flame comes near the solution.
Some of this discovery was probably prompted by the accident Curtius had when he attempted to melt-close the thread of a capillary glass tube
containing 2mL of a 27% aq. solution of HN3, doing so caused a violent detonation that pulverized all the glass and nearly caused injury (he described
this in Ber. 23, 3027).
Pure hydrazoic acid is a different monster altogether. Something completley mundane like a gas bubble rising from the boiling liquid grinding against
sharp edges of glass can cause detonation. Vibrations easily set it off, but it can appear that even nothing at all also sets it off. In Gmelin and
Brethericks there is a description of its reactivity to stimulus. And the heat of explosion of liquid HN3 is 1460cal/g (HMX is 1480 cal/g). Not worth
even preparing for novelty interest.
Anyways, Curtius findings is mainly what I've held the standard of saftey to for this compound. Hydrogen azide is extremely poisonous and even making
a few milligrams outside working with a chemical respirator on it has caused me mild but noticeable effects (respiratory effects, like dyspnea-type
effects). Wafting a tiny amount of air with the hand from a flask (having a small amount of HN3) somewhat near my exposed face caused me to
immediatley quit breathing out of my nose. Hydrazoic acid is said to have pungent odor resembling PH3, it doesn't matter all that much since not all
may be able to smell it and it attacks breathing organs right away.
It is known hydrazoic acid can also be made from acidic oxidation of hydrazine sulfate in somewhat decent yield, this can also be bubbled into
alcoholic NaOH to precipitate azide salt. I've used this method a couple times and it may be good for small amounts, since hydrazoic acid is extremely
poisonous and dangerous.
Chemical Lecture Experiments by Francis Gano Benedict illustrates the oxidation of the sulfate using only colorless mediocre weak conc.
nitric acid (regular or fuming conc. HNO3 would likely destroy much, if not all HN3), the illustration can be seen below but I've added the warning,
based on what we know about HN3 flame sensitivity, so it's recommendable not to use any flame. Mild warming from a hot plate or hot bath is all that
is necessary anyways, after which the reaction becomes spontaneous. They bubble the gas into aq. AgNO3 to generate AgN3. The first few times in doing
the reaction I put a big cinder block in front of the apparatus trajectory, because of fear of explosion mainly surrounding possibility of generated
gas being able to travel back to the heat source.
The reaction between hydrazine sulfate and the acid can foam vigorously out of control, so that it is best not to do the reaction in a narrow
container. But if the receiving mixture is ethanolic NaOH (using absolute EtOH is best), when the evolved gas from the reaction mixture is bubbled
through this, NaN3 precipitates. And no condensation should come over and this should minimize any signifcant amount of nitrate being formed.
But other oxidants like aq. H2O2 with H2SO4 can be used for the oxidation where nitrate contaminant is feared, though I think the oxidation with
nitric acid is more simple and should give a relatively pure product. I've had a procedure with described crude yield for this reaction, but don't
know what happened to it. The procedure can likely be modified to use KOH, CsOH or RbOH instead of NaOH.
Also note that if you have formed an aqueous solution of NaN3, KN3, etc. you can not simply boil down the this mixture to isolate the salt, because it
causes hydrolysis of the salt and the HN3 is split off and thus released, so evaporation needs to be done at moderate temperatures (and free from
CO2).
Attachment: Ber. 49, 1742.pdf (589kB)
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Attachment: Brethericks-HN3.pdf (117kB)
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Attachment: Chemical-lecture-experiments.pdf (88kB)
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