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Author: Subject: The short questions thread (2)
woelen
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[*] posted on 3-7-2009 at 02:20


This could indeed be done. Mix the two chemicals (crushing to a powder may be necessary, before mixing) and add a few drops of water to make a slurry of the chemicals. carefully heat the mix, but do not heat very strongly, the mix should remain wet and you should not boil off the water in the mix.

Even easier is the making of ammonia from solid crushed ammonium sulfate and solid crushed sodium hydroxide. You must add a few drops of water to the mix, such that the material is just wet. Initially the reaction sets of, but at the end you may need some heating.

Simply leading the gas into water may be somewhat risky. There is a chance of suckback into the flask in which the ammonia is made (remember: ammonia dissolves in water exceptionally well). Better is to use an inverted small funnel and put that inverted funnel just on the water surface.
Another good thing to do is start the preparation with household ammonia instead of plain water. The first 5% of ammonia you then already have in the liquid.




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[*] posted on 3-7-2009 at 03:00


Short question:
I might have the chance to buy a aspirator pump.
The seller says that it sucks 200 Mbar ( i think he means milli :P )
So how must is look at this, is it strong enough to use for normal lab apps?




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[*] posted on 3-7-2009 at 03:45


Quote: Originally posted by woelen  
Better is to use an inverted small funnel and put that inverted funnel just on the water surface.

The inverted funnel doesn't work well with NH3 as NH4OH is less dense than water.
The gas dissolves in the top few mm and is lost from the funnel when those few mm are saturated. . .
Bubbling the gas into the bottom of the container of water works well but a container or plastic bulb fitted inline to the hose to accommodate suckback is necessary.

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entropy51
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[*] posted on 3-7-2009 at 11:13


Put a small magnetic stir bar ("flea") in the beaker under the funnel and spin it just enough to mix the NH4OH collecting in the funnel with the water beneath it. NH3 sucks back so easily that the funnel is a good idea. Using a setup like this I can dissolve so much NH3 in the water that it requires external cooling; it's quite exothermic as more and more NH3 dissolves.

I agree that NaOH seems to work better than Ca(OH)2, if the NH4SO4 is finely ground.

Manimal, aren't those old Popular Science articles great? The author of that one, Kenneth Swezey, later collected the articles in a great book called Chemistry Magic. Sometimes you can find it for sale on abebooks.com.

[Edited on 3-7-2009 by entropy51]
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[*] posted on 3-7-2009 at 14:06


You should be able to get ammonia by heating almost any ammonium salt with almost any non-volatile basic compound. As a schoolboy I had my first experience of ammonia when I heated red lead with ammonium chloride. Got a huge toot of ammonia, and probably some lead volatiles too, when I gave it a good sniff.
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[*] posted on 3-7-2009 at 18:02


I will try this on a small scale with about 5g ammonium chloride and sodium hydroxide. I'm guessing the equation looks like:
NH4Cl + NaOH -> NH3 + NaCl + H2O

An issue with this is that I will not end up with particularly conc. Ammonia solution, because I will also be boiling off water...
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[*] posted on 3-7-2009 at 18:11


As Woelen said "add a few drops of water to make a slurry of the chemicals. carefully heat the mix, but do not heat very strongly, the mix should remain wet and you should not boil off the water in the mix."

That is there is little water to boil off and the NH3 comes off without strong heating.

You will not end up with a particularly concentrated Ammonia solution because 5 g of NH4Cl will yield very little NH3. If you scale it correctly, you can get quite strong ammonia solution, especially if you repeat it a few times. But keep the absorbing water cool. Otherwise it becomes quite hot and will dissolve no more NH3. Cold water absorbs much more NH3 than warm water.


[Edited on 4-7-2009 by entropy51]
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[*] posted on 3-7-2009 at 18:33


I see, so about 60 degrees as opposed to a couple of hundred?
And yes I'm not really doing it to make Ammonia solution to use, more to try it out (on a test tube scale)
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[*] posted on 3-7-2009 at 18:39


Quote: Originally posted by User  
Short question:
I might have the chance to buy a aspirator pump.
The seller says that it sucks 200 Mbar ( i think he means milli :P )
So how must is look at this, is it strong enough to use for normal lab apps?


1 bar = 750 mTorr = 750 mm Hg

So 200 mbar = 1/5 bar = 150 mm Hg

Normally aspirators connected to the sink produce 10 or 20 mm Hg depending on the water temperature, and that is often a good pressure for vacuum distillation. Many materials boil about 100C less than their atmospheric BP at 10 mm Hg. It depends on the boiling point of the material you are distilling as to whether a higher pressure is acceptable. 150 mm is on the high side for many distillations.

I have a pump that pulls about 150 mm Hg and I sometimes use it for vacuum filtrations, but it's not that good and I mostly use the sink aspirator.

A good aspirator connected to the faucet doesn't use all that much water, especially for distillations in a system free of leaks, unless you are using it 24/7.
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[*] posted on 3-7-2009 at 18:45


Quote: Originally posted by Mossydie  
I see, so about 60 degrees as opposed to a couple of hundred?
And yes I'm not really doing it to make Ammonia solution to use, more to try it out (on a test tube scale)


The test tube will contain a very strong ammonia solution and very little heat should be needed to release it, certainly less than 200 C. Normally by the time an aqueous NH3 solution reaches the boiling point of water all the NH3 has escaped as NH3 gas.
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[*] posted on 3-7-2009 at 19:10


So one could effectively concentrate dilute ammonia solution by heating it gently and bubbling the resulting gas through water - would this not be easier than the ammonium salt / basic compound method?
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[*] posted on 3-7-2009 at 19:33


Can aluminum chloride hexahydrate be dehydrated by heating somewhere below the sublimation point (178 C)?



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[*] posted on 3-7-2009 at 20:06


No, it gives off HCl leaving various chlorohydrates, and I believe finally AlO(OH). SFAIK there is no way to dehydrate the hydrate chloride that is less challenging than directly making the anhydrous chloride from the metal.

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[*] posted on 3-7-2009 at 20:11


I have seen a patent that performed the dehydration of the hydrate but N_I's words still ring true, your better off just starting anhydrous and going that route.




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[*] posted on 4-7-2009 at 07:21


Quote: Originally posted by Mossydie  
So one could effectively concentrate dilute ammonia solution by heating it gently and bubbling the resulting gas through water - would this not be easier than the ammonium salt / basic compound method?


Well you can do that, but it's a long and slow process. Dilute ammonia is maybe 6% around here but if you can find the stronger "janitorial" strength the effort would be more worthwhile. Most ammonia I see nowadays also contains detergents which may foam too much when heating. You can make tons of NH3 with a bag of (NH4)2SO4 fertilizer and NaOH or Ca(OH)2.
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[*] posted on 4-7-2009 at 09:42


I attempted to make some 2-bromopropane yesterday in the same way that smuv did in an ollllllllllld thread. This was basically H2SO4 + HBr + iPrOH. Does anyone know what specifically is that horrible acrid smell when the distillation begins?

As an aside, I'm pretty sure I got mostly diisopropyl ether and propene, followed by some actual isopropyl bromide containing (possibly, as I couldn't smell it, which is odd) dissolved bromine and a vast quantity of HBr fumes. I'm going to try a different way.

On the other, other hand, the reaction mix turned this incredible indigo-purple color halfway through, which hydrolyzed when I added some water. Is it possible that I had some polyatomic bromine cations floating around?




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[*] posted on 4-7-2009 at 11:13


UC, that might go better without the H2SO4, which I don't think you need. Could the acrid smell be SO2? HBr, some of which probably distills?

I've made 2-Br propane by distilling constant boiling HBr and iPrOH per Vogel, page 387 of the 4th edition (probably also in the older editions).
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[*] posted on 4-7-2009 at 11:26


Quote: Originally posted by entropy51  
UC, that might go better without the H2SO4, which I don't think you need. Could the acrid smell be SO2? HBr, some of which probably distills?

I've made 2-Br propane by distilling constant boiling HBr and iPrOH per Vogel, page 387 of the 4th edition (probably also in the older editions).


Well, the thread claimed 90% yield, so I gave it a shot (also because I don't have any HBr) It was more like a rancid nasty organic smell. It came over before any distillate and passed through my NaOH scrubber, so I don't know.

I think I'll just make some HBr next time and go that route.

I do seem to have squeezed a roughly 30% yield out of it, although I am not happy with this.

[Edited on 7-4-09 by UnintentionalChaos]




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[*] posted on 4-7-2009 at 11:58


UC, here's a nice HBr prep in case you don't already have a favorite method. I hope it's useful.

Attachment: HBr.pdf (1.2MB)
This file has been downloaded 2841 times
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[*] posted on 4-7-2009 at 15:06


what would one use large amounts of KBrO3 ?
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manimal
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[*] posted on 5-7-2009 at 19:35


Does anyone know the percentages of xylene isomers contained in commercial xylene solvent?

And: what the fuck happened to roguesci.org?

[Edited on 6-7-2009 by manimal]
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[*] posted on 5-7-2009 at 20:52


it varies. Look at the wiki page for xylene.
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[*] posted on 5-7-2009 at 22:00


The technical-grade stuff contains ortho-, meta-, and para-xylenes, C6H4(CH3)2, of the same composition. There may be minor amounts of ethylbenzene (which has the same composition and molecular weight as xylenes); and possibly of other substances such as ethyl-substituted homologs, toluene, trimethylbenzenes, and durene.
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[*] posted on 5-7-2009 at 23:01


Quote: Originally posted by JohnWW  
The technical-grade stuff contains ortho-, meta-, and para-xylenes, C6H4(CH3)2, of the same composition. There may be minor amounts of ethylbenzene (which has the same composition and molecular weight as xylenes); and possibly of other substances such as ethyl-substituted homologs, toluene, trimethylbenzenes, and durene.


You're just about guaranteed a few percent of ethylbenzene. This can interfere, depending on what you want the xylene for. I recrystallized sulfur from hardware store xylene a while back and when I dried it, found clumps of gummy gray material clinging to my sulfur. The reason?

ethylbenzene + sulfur -(heat)-> H2S + styrene

And some of the styrene probably polymerized on me. Come to think of it, the little globs looked an awful lot what you get when you splash acetone on expanded polystyrene foam. I've wondered if this is actually a good way to remove ethylbenzene entirely and make a more useful product. Perhaps a KMnO4 wash to cleave the styrene's double bond before distillation would be in order.

Unrelated question:

Does anyone know where to get tiny hard polystyrene prills? About the consistency of ion exchange resin beads is needed, because that is more or less the project.

[Edited on 7-6-09 by UnintentionalChaos]




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[*] posted on 6-7-2009 at 01:31


Ethylbenzene can be present in quite large amounts. According to it's MSDS, Startex brand xylene contains 20% ethylbenzene.
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