Sciencemadness Discussion Board

Alkali metal synthesis

APO - 28-12-2012 at 20:34

Now, if I use the method found here but using rubidium hydroxide, or caesium hydroxide, will it work? Nurdrage could'nt get sodium hydroxide to work, so I asume it won't work using lithium hydroxide either.

kristofvagyok - 28-12-2012 at 20:54

It will work. The only problem when working with Cs and Rb is to use inert atmosphere, avoid self ignition and also note that CsOH is a really strong base, it will eat the glass. It would be much safer to turn CsOH and RbOH to their carbonates and make magic with that.

Here is a patent what describes production of Cs from Cs2CO3 and Mg: www.google.com/patents/US3432156

And there is also a thread somewhere about this...

Here is some Rb kept under argon:

UnintentionalChaos - 28-12-2012 at 23:06

http://www.sciencemadness.org/talk/viewthread.php?tid=6981&a...

http://www.sciencemadness.org/talk/viewthread.php?tid=20496

APO - 28-12-2012 at 23:56

Thanks! I didn't know I could use carbonates, would the Cs2CO3 and Mg method work with Rb2CO3?
Also their carbonates are way cheaper. Could this be done using thier chlorides?

12AX7 - 28-12-2012 at 23:59

Oddly enough, MgCl2 isn't favorable -- these methods depend on the stability of MgO (or Al2O3), both of which are remarkably robust.

Tim

APO - 29-12-2012 at 00:09

Would nitrates work? I have access to lots of caesium and rubidium nitrate, but I would obvioulsy prefer the elemental form.
Does tert-butanol work as a catalyst for the carbonate method? Also I will need a solvent that caesium will float in to help it coalesce into a sphere, tetrahydronaphthalene isn't dense enough.

kristofvagyok - 29-12-2012 at 05:03

Quote: Originally posted by APO  
Would nitrates work?


I don't want to be rude, but it is highly recommended to start with something easier, like... Potassium?
It could also easily ignite, it's an alkali and it's perfect to gather some experience.

And no, starting from a nitrate won't work.

Endimion17 - 29-12-2012 at 06:36

Alkali metal isolation. Elements can't be chemically synthesized.

Leaning the basics of chemistry before dealing with such dangerous substances is a must.

kavu - 29-12-2012 at 06:40

Quote: Originally posted by Endimion17  
Alkali metal isolation. Elements can't be chemically synthesized.

Leaning the basics of chemistry before dealing with such dangerous substances is a must.


Nope, synthesis is a correct term here. Isolation would indicate that the metal would already be at oxidation state (0) and just separated from something else.

Actually the most accurate term would be analysis. In the early days of chemistry braking stuff down was called analysis and combining stuff was synthesis. Nowadays synthesis is however more widely used do describe the process of making new compounds/elements from other ones.

For comparison, would the following then be the isolation of hydrogen: HCl + 2 Zn → ZnCl2 + H2?

[Edited on 29-12-2012 by kavu]

Endimion17 - 29-12-2012 at 07:46

Quote: Originally posted by kavu  
Quote: Originally posted by Endimion17  
Alkali metal isolation. Elements can't be chemically synthesized.

Leaning the basics of chemistry before dealing with such dangerous substances is a must.


Nope, synthesis is a correct term here. Isolation would indicate that the metal would already be at oxidation state (0) and just separated from something else.

Actually the most accurate term would be analysis. In the early days of chemistry braking stuff down was called analysis and combining stuff was synthesis. Nowadays synthesis is however more widely used do describe the process of making new compounds/elements from other ones.

For comparison, would the following then be the isolation of hydrogen: HCl + 2 Zn → ZnCl2 + H2?

[Edited on 29-12-2012 by kavu]


Isolation in the strictest sense would be removing a chemical element or chemical substance from a mixture, but in the broad sense can be used instead analysis. Rarely anyone uses the word "analysis" in its original meaning. Today, it's connected to quantitative and qualitative analysis and not preparative chemistry.

Chemical elements can not be chemically synthesized. They can be synthesized in nuclear reactors which is a nuclear process.
Chemical compounds are chemically synthesized, and chemical elements are chemically isolated from chemical compounds.

Yes, the process behind the equation you've wrote would be isolation of hydrogen by oxidizing it using zinc as a reducer. It's a typical school example of hydrogen isolation by a redox process.


I'll continue repeating this ad nauseam, every time I see someone making these grossly huge errors here. I can not believe people are dealing with incredibly dangerous substances under highly energetic conditions and having the lack of knowledge kids learn in elementary schools.

Before the age of YouTube, I've never ever seen any professor talking about "synthesizing sodium". I've never read a single book talking about "synthesizing chlorine". Those were nonexistent phrases. I saw "making" and "producing", but not "synthesis".

It's a disinformation, a bad meme that's spreading around because more and more uneducated people have opportunities to do amateur chemistry via Internet. I'm not saying that's a bad thing, of course. The more of them, the better, but proper information must be spread, and the bad information should be stopped and corrected.

kavu - 29-12-2012 at 09:28

Quote: Originally posted by Endimion17  

Before the age of YouTube, I've never ever seen any professor talking about "synthesizing sodium". I've never read a single book talking about "synthesizing chlorine". Those were nonexistent phrases. I saw "making" and "producing", but not "synthesis".


Synthesis has indeed become synonymous to "prepare" in some sense. It sounds odd to me as well to say "synthesis of hydrogen" or any other element for that matter. I'm not however very concerned with this synthesis business as there are terms which are much more often confused and lead to more serious problems in theoretical and experimental work. Such terms would include stereospecific, R/S/D/L/d/l/+/- and so forth.

Interestingly enough IUPAC has not given a definition for synthesis.

You are indeed very much spot on with the term isolation! I have been too busy with natural products chemistry and grown to forget the use of "isolation" when talking about elements. My bad!

[Edited on 29-12-2012 by kavu]

plante1999 - 29-12-2012 at 10:19

Synthesis mean to combine.

I would say that the previous equation was zinc chloride synthesis/metathesis with hydrogen byproduct.

Hexavalent - 29-12-2012 at 10:27

A well-known online dictionary generally defines synthesis as;

"the combining of the constituent elements of separate material or abstract entities into a single or unified entity"

in terms of chemistry as;

"the forming or building of a more complex substance or compound from elements or simpler compounds"

and analysis as;

" the separating of any material or abstract entity into its constituent elements"


Source: http://dictionary.reference.com/browse/synthesis

[Edited on 29-12-2012 by Hexavalent]

kavu - 29-12-2012 at 10:33

Quote: Originally posted by plante1999  
Synthesis mean to combine.

It pretty much depends on what the target of the synthesis is. If for example a protecting group is removed from a delicate organic structure you combine the protecting group with the removal reagent. However no one in their right mind would call that a synthesis of Me3SiFOH with the original target as a side product. This is why I think it's not all that important to get too picky with when the word synthesis is applicable and just go with the flow.

[Edited on 29-12-2012 by kavu]

Hexavalent - 29-12-2012 at 10:45

I have to agree with kavu: our poorly-controlled use of terms globally often leads to some seeming synonymous and others antonymous to one another. Indeed, for instance, is the electrolysis of water the synthesis of oxygen and hydrogen molecules - which 'weren't there' previously, or isolation as the elements cannot be chemically synthesised? Is this whole process 'analysis', as we are breaking down the water molecules?

Perhaps someone should email IUPAC, asking for exact definitions.

(Just like to announce, on this day 1 year ago, I joined sciencemadness.org, and have enjoyed every moment of activity here. I have learned so much and the knowledge gained is invaluable: thank you to you all).

DJF90 - 29-12-2012 at 11:07

Isolation is the separation of one desired component from a mixture, wherein the target material already exists. Analysis isn't really an applicable term here, because we're talking preparatively. Extraction is another term that is similar to isolation, but can also mean the production of an element from an ore.

Saying this, I can also see why synthesis is a misleading term. However, obtention of an element from of the desired element-bearing material can also be classified as the "synthesis of element X". It might help to specify somewhat. Technically its a "chemical synthesis", as opposed to a "nuclear synthesis". I dont think synthesis should be restained to the adding of mass/atoms. "Preparation" is a much less ambiguous term.

blogfast25 - 29-12-2012 at 13:11

Leaving the pissing contest on 'synthesis, isolation' aside [cough!], it would be great if someone here could actually produce Rb using the KOH + Mg + t-alcohol method. But I would suggest to have a 'dry run' with KOH first: numerous experimenters here had a hard time with KOH, never mind something really expensive as RbOH...

APO - 29-12-2012 at 13:19

Sorry for saying synthesis, you all knew what I meant. Just so all of you know I do have many safety precautions, such as eyes and respiratory protection ,also a powder based fire extinguisher. I do think making potassium would be a good start, I won't be making any alkali metals until I get the right equipment. I've only gotten lithium from a battery once, and made sodium oxide.

But I really do need a solvent that is highboiling and very dense (at least 2.00/ml) I'm pretty sure tert-butanol can be used as a catalyst for most reductions, but can we please focus on the isolation of alkali metals and not vocabulary?

APO - 29-12-2012 at 13:27

So there is magnesium reduction with a solvent and catalyst involved, this works for potassium, rubidium and caesium production. Also to make sodium you can mix magnesium powder and sodium hydroxide powder, and then light it inside a crusible. But what about lithium production?

bbartlog - 29-12-2012 at 13:48

Molten lithium will attack glass, so even if the method outlined for potassium would work, it would have to be conducted in some other sort of vessel.

APO - 29-12-2012 at 13:54

Lithium will attack glass? Even boroscilicate glass?

blogfast25 - 30-12-2012 at 08:07

APO: have you consulted the 'sticky' potassium thread above? Great information on solvents for K production!

[Edited on 30-12-2012 by blogfast25]

shannon dove - 30-12-2012 at 09:38

I suppose you could convert you nitrate to a carbonate by mixing it with charcoal and igniting it. Some of your carbonate might be contained in the smoke.

APO - 30-12-2012 at 13:06

I would rather use carbonic acid to produce carbonates. I didn't look at the sticky thread so much because I had trouble finding dense solvents, can I get an example of a solvent at least twice as dense as water?

shannon dove - 30-12-2012 at 13:18

I don't think carbonic acid is going to convert your nitrate into a carbonate.
but maybe I'm wrong.

kristofvagyok - 30-12-2012 at 13:43

Quote: Originally posted by APO  
I would rather use carbonic acid to produce carbonates. I didn't look at the sticky thread so much because I had trouble finding dense solvents, can I get an example of a solvent at least twice as dense as water?

Just a suggestion: learn some chemistry, elseway you could easily lost of a few fingers, maybe one eye and a lot blood. So be clever.

Converting Cs/Rb-nitrates to carbonates with carbonic acid is impossible. Where did you study chemistry?

Dense solvents are usually halogen containing e.g.: chloroform, carbon tetrachloride, trichloroethylene ect. And all of these WILL REACT with Cs, Rb and even with Mg. But may I ask why do you need dense solvents? If the Cs floats on the surface of something than I will give you a 100% for it's autoignition.

A little more info from Rb/Cs chemistry: mixing metallic Cs with benzene will yield a black amorphous participate what will explode on contact on almost everything, it has a chemical formula: C6H6Cs6.
Mixing metallic Cs with toluene will also end up with a reaction and yield C6H5CH2Cs. So be smart what kind of solvent would you want to use.

APO - 30-12-2012 at 14:18

I'm talking about using carbonic acid to make carbonates from hydroxides, not nitrates, sorry for being non specific. I'm looking for a solvent that is unreactive with caesium and rubidium so that they will coalesce easier, and I know that they won't auto ignite if I use an inert atmosphere. So if you think that won't work how can I make make rubidium and caesium from their carbonates efficiently and safely?

blogfast25 - 31-12-2012 at 09:48

Quote: Originally posted by APO  
I would rather use carbonic acid to produce carbonates. I didn't look at the sticky thread so much because I had trouble finding dense solvents, can I get an example of a solvent at least twice as dense as water?


The type of inert solvents you need for reducing alkali hydroxides with Mg ALL have densities close to 1. To increase density you’d need to introduce heavier exo-atoms into the molecules but then you can forget about inertness. Even fluorinated solvents react with hot alkali metals.
Quote: Originally posted by APO  
I'm talking about using carbonic acid to make carbonates from hydroxides, not nitrates, sorry for being non specific. I'm looking for a solvent that is unreactive with caesium and rubidium so that they will coalesce easier, and I know that they won't auto ignite if I use an inert atmosphere. So if you think that won't work how can I make make rubidium and caesium from their carbonates efficiently and safely?


Why do you want to use carbonates? The patent calls for hydroxides for a reason: that they can form thermodynamically favourable MgO in the presence of a catalyst like t-butanol.

Most fully saturated hydrocarbon solvents are totally unreactive towards alkali metals.


[Edited on 31-12-2012 by blogfast25]

APO - 31-12-2012 at 11:13

Maybe if I do a reduction of caesium/rubidium hydroxide with magnesium, then I could take the smaller spheres and coalesce them later. I wanted to use carbonates because they aren't very strong bases, I'm told molten caesium hydroxide
will even eat through boroscilicate glass.

blogfast25 - 1-1-2013 at 05:09

Quote: Originally posted by APO  
I wanted to use carbonates because they aren't very strong bases, I'm told molten caesium hydroxide will even eat through boroscilicate glass.


Molten CsOH maybe, but at 200 C you're still a long way off the MP. Some superficial damage to your glass may happen though: one or two experimenters reported that with KOH.

Valentine - 1-1-2013 at 10:31

If you have Cs2CO3 or CsOH, you can turn it into CsN3 by hydrazoic acid. The caesium azide will decompose on heating (under N2 or argon atmosphere) to Cs and N2. I have a good method for preparing hydrazoic acid, I have attached it (Inorganic Syntheses 1939, 1, 78-79.) If you have some ethereal solution of HN3, just pour it onto some concentrated Cs2CO3 or CsOH solution, then decant the ether (so you can recycle it) and evaporate the CsN3 solution. Of course, the aqueous method could be more efficient, but I think it's a bit more dangerous, because for CsN3 you don't have to distill or dry the ethereal HN3, but you have to distill the aqueous solution and I don't like the idea of hot HN3... I found some literature about the decomposition of the alkali azides, I have attached it, too (Zeitschrift für anorganische und allgemeine chemie 1926, 152, 52-58.)
Finally, I have found some paragraphs in the Kirk-Othmer Encyclopedia of Chemical Technology about rubidium and cesium, the latter mentions the azide way of obtaining pure Cs at 390 °C.
I have a few kgs of caesium chloride, other cesium salts, and some sodium azide too, so I will try this method soon. I think it's a way more simple and clean method than the reduction of molten CsOH or Cs2CO3 under inert atmosphere, if you have NaN3 and you don't get carried away with the amount of CsN3 to decompose at once.
It is my first post at ScienceMadness, I hope you find it useful.

Attachment: Hydrazoic acid.pdf (86kB)
This file has been downloaded 1600 times

Attachment: Alkali azide decomposition.pdf (419kB)
This file has been downloaded 661 times

APO - 1-1-2013 at 12:04

Thanks!

Valentine - 17-1-2013 at 05:55

P1060933k.jpg - 108kB
Some cesium azide is ready :)
I found a good article about the decomposition process.

Attachment: Cesium azide decomposition.pdf (575kB)
This file has been downloaded 2432 times

kristofvagyok - 17-1-2013 at 10:29

Quote: Originally posted by Valentine  
Some cesium azide is ready :)
I found a good article about the decomposition process
This recipe is also in Georg Brauer-s book: "Handbook of Preparative Inorganic Chemistry". -according to my memory but I never ever tried it because of the potential explosion.


If you have made it successfully please post the results, I also want to try it out(:

Valentine - 22-1-2013 at 08:47

I could do some preliminary experiments with cesium azide today. I used 2 g azide in a test tube with ground glass joint, a bath temperature of 400 °C, N2 atmosphere, a little piece of pure copper as catalyst, 10 mbar vacuum and I listened to classical music the whole time. Bubbling began just after the salt had melted. Slowly the colorless liquid turned to dark blue. After half an hour the bluish coloration started to fade, and the liquid became golden yellow. The cesium creeps to the walls of the tube pretty high, and it refuses to flow down. Next time I will use a lot more suitable apparatus, and I will distill the cesium directly into an ampoule.



Azide in the bath



Temperature



Bubbling



The result



The result


I forgot taking pictures when the liquid was blue, so I will take some next time.

elementcollector1 - 22-1-2013 at 10:05

Is there any condensed cesium in there, or has it all mirrored onto the tube walls?

APO - 22-1-2013 at 11:05

Can someone explain hydrazoic acid procution to me? The pdfs are sorta confusing.

Valentine - 22-1-2013 at 12:22

Quote: Originally posted by elementcollector1  
Is there any condensed cesium in there, or has it all mirrored onto the tube walls?


No, it was just a really thick mirror on the whole tube, next time I will use 15-20 g azide, and a better apparatus to get an ampoule of cesium. I still think as a preliminary experiment it was quite good.
Quote: Originally posted by APO  
Can someone explain hydrazoic acid procution to me? The pdfs are sorta confusing.


You put some concentrated aqueous solution of NaN3 into a RBF, then add some ether and you will get a two phase system. The HN3 is a weak acid, so if you add sulfuric acid to the system, a lot of HN3 will form in your lower (aqueous) layer. HN3 is quite apolar, so most of it will go into the upper (ethereal) layer. If you add enough H2SO4, you will get an almost pure aqueous Na2SO4-H2SO4 solution as a lower layer, and an ethereal HN3 solution as an upper layer. If you have a solution of a strong base (for example CsOH), just pour the upper layer from the previous flask onto it, and as a weak acid, the HN3 will neutralize the strong base (CsOH) and goes into the lower layer forming a solution of an azide salt (CsN3). It's just playing with polarity and acid-base equilibria.

APO - 22-1-2013 at 15:14

If I mix sodium bisulfate and sodium azide and heat it to the melting point of the sodium bisulfate, will I be able to distill some hydrazoic acid over?

kristofvagyok - 22-1-2013 at 23:01

Valentine, this is awesome! Congrats for the Cs!

What do you think about using a paraffin wax what would only melt under 400C? I would use it as a reaction medium and a protecting from the evaporation of the Cs.
Quote: Originally posted by APO  
If I mix sodium bisulfate and sodium azide and heat it to the melting point of the sodium bisulfate, will I be able to distill some hydrazoic acid over?


Everything you learned about chemistry in school was either a) wrong, b) a crude approximation, because hydrazoic acid will explode under those conditions.

Valentine - 23-1-2013 at 00:25

At 400°C cesium might react slowly with the forming decomposition products of wax. If the wax is stable enough, the cesium might endure the entire procedure, but you will get a messed up tube in either case. How would you dispose of the wax after the decomposition?

elementcollector1 - 23-1-2013 at 10:10

This might be the way to go for most home chemists, but... Cesium and oxygen... plus the availability of hydrazoic acid and/or ether, and cesium salts in the first place... I'm very conflicted as to trying this out. Also, it's very wet where I live, if you know what I mean. I would imagine the way to go would be to distill cesium under vacuum into a pre-formed ampoule, and seal it up while still under vacuum for best results. Or you could do what Zan Divine did and distill from a mixture of Li and CsOH under vacuum/argon in a steel distilling apparatus leading into a RBF, and then, while still under vacuum/argon, somehow get your sample into a good-sized ampoule. I don't know the specifics, and Zan's been out for a while (contacted him recently, he's not dead), so... Any further ideas?

APO - 23-1-2013 at 11:13

What's the safest way to make HN3?

Valentine - 23-1-2013 at 11:28

Quote: Originally posted by APO  
What's the safest way to make HN3?

The NaN3 - H2SO4 ehtereal method without doubt. In a previous post I described it quite well.

APO - 23-1-2013 at 11:57

Any way I can make HN3 without H2SO4?

I'm so damn scared of sulphuric acid.

Valentine - 23-1-2013 at 12:56

Quote: Originally posted by APO  
Any way I can make HN3 without H2SO4?

I'm so damn scared of sulphuric acid.

Don't be scared, sulfuric acid is your friend :)
If you still in fear, you could use 85% phosphoric acid

ScienceSquirrel - 23-1-2013 at 13:43

If you are scared of sulphuric acid you should not attempt to make hydrazoic acid.
Although it is a weak acid it is quite toxic and can explode.

http://en.wikipedia.org/wiki/Hydrazoic_acid

APO - 23-1-2013 at 14:33

I'm scared of sulphuric acid because of it's corrosiveness and that it's a liquid,so unlike most bases I can just scoop it back up, but with sulphuric acid it would probably do some damage.

However I'm experienced with toxic, explosive, and mildly corrosive compounds.

Also under what conditions will hydrazoic acid explode?


[Edited on 23-1-2013 by APO]

APO - 23-1-2013 at 14:43

Also I decided that instead of using HN3 to make caesium azide to make caesium, that I could react magnesium and caesium carbonate under a vaccum at 200 celcius and then distill the caesium at 482 fahrenheit.

However to make my rubidium I may use the azide method.

[Edited on 23-1-2013 by APO]

EdMeese - 25-1-2013 at 14:56

Anybody have any experience with Na-B''ASE solid electrolyte systems? 99.999% pure sodium from a ZnCl2-CaCl2-NaCl bath in one report, at 200 °C so lower energy Green blah blah. I've been obsessed for some time with the idea of doing this on an NaOH or NaOAc system closer to 100 °C -- still want to be able to siphon off pure liquid sodium for continuous operation -- and ready to take the plunge to build one. There's an outfit in the UK that sells BASE ceramic shapes but they start at ~100$ so I thought I'd try to make one with this Indian microwave method: Solid State Ionics 158 (2003) 199– 204. Thoughts?

"The conventional processing of sodium beta alumina (SBA) which includes synthesis and sintering in order to obtain dense bodies for use as a solid electrolyte usually requires high temperatures (f1873 K) and long holding times. A substantial amount of Na2O may be lost during such high temperature heat treatment, the consequence of which may be deleterious when SBA is used for device applications. To overcome this problem, the use of microwave as heating source for the processing of SBA was attempted in the present process. Preliminary investigations on commercial Na-h-alumina bearing the composition NaAl11O17 were carried out and it was found that Na-h-alumina was a very good absorber of microwaves. Following this, a microwave active precursor of MgO-stabilized Na-hW-alumina was synthesized through a novel solution technique, which on compaction and exposure to microwaves yielded dense (94% theoretical) and phase-pure Na-hW-alumina body. This paper presents the results of a direct single-step synthesis cum sintering of MgO-stabilized Na-hW-alumina."

APO - 5-2-2013 at 19:11

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?

elementcollector1 - 5-2-2013 at 20:33

According to the original patent, it's supposed to work for cesium. No idea on rubidium, but it seems that it would follow.

APO - 6-2-2013 at 08:15

Awesome.

elementcollector1 - 6-2-2013 at 09:15

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...

blogfast25 - 6-2-2013 at 10:08

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]

elementcollector1 - 6-2-2013 at 10:11

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?

APO - 6-2-2013 at 11:00

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.

kristofvagyok - 6-2-2013 at 11:48

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.

blogfast25 - 6-2-2013 at 14:41

Quote: Originally posted by APO  

I'm confident it will work so I will post my results when finished.


Yeah, I look forward to that!

blogfast25 - 6-2-2013 at 14:47

Quote: Originally posted by kristofvagyok  
1: we do not heat Erlenmeyer flask on hotplate, because it could easily break. Rule no.1: always use round bottomed flasks.



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.

APO - 6-2-2013 at 17:52

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.

phlogiston - 8-2-2013 at 07:43

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]

APO - 8-2-2013 at 11:11

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]

APO - 14-2-2013 at 11:29

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...

Zan Divine - 14-2-2013 at 13:36

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.

The Setup sm.jpg - 88kB all cesium cropped.jpg - 51kB




[Edited on 2/14/2013 by Zan Divine]

phlogiston - 14-2-2013 at 13:43

That method will only work for poorly soluble azides. The azides of rubidium, caesium, barium and strontium are all quite soluble.

APO - 14-2-2013 at 15:32

Aw man.



[Edited on 14-2-2013 by APO]

Zan Divine - 16-2-2013 at 09:17

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]

a_bab - 16-2-2013 at 09:33

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".

Formatik - 24-2-2013 at 17:12

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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