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Author: Subject: Uranium Isolation
ShadowWarrior4444
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[*] posted on 21-5-2008 at 23:07


Quote:
Originally posted by The_Davster
Your first statement is erroneous. Alpha emitters pose little harm while outside the body, however inside the body they are extremely dangerous as they emit right next to your cells, damaging them. Hence any alpha emitter becomes dangerous, especially in powder form, and alpha emitting dusts are known to cause cancer.

While the ore grindings may not be too dangerous in the low doses expected with grinding down to a fine powder, respiratory protection would be strongly advisable.
And defiantly in any later steps with purified uranium compounds, all precautions must be taken to avoid inhalation.

I would sooner work with cyanogen bromide than radioactive dusts.

EDIT: Looks like Polv' beat me to it by 5 min:P



[Edited on 21-5-2008 by The_Davster]


Apologies, wasn't paying enough attention at the time: I meant to say that Uranium has not been known to cause human cancers (http://www.atsdr.cdc.gov/toxprofiles/phs150.html,) due to its long half-life and mode of decay. Then subsequently note that the decay of Uranium can produce Radon gas which should be taken into account when storing the ore/compounds.

When working with any fine particulate it is important to wear respiratory protection. Silicosis (or a similar ailment) is not a very nice thing to have.


Quote:
One extractive method that might be practical would be to convert to crude uranyl nitrate, add a small amount of aluminium sulfate to tie up any fluoride, evaporate to dryness, and extract with acetone, isopropyl or ethyl alcohol, then evaporate off the solvent to obtain the purified uranyl nitrate.



Will extraction by nitric acid work for all ores if they are at a small enough particulate size? I seem to recall certain uranium ores were somewhat immune to extraction via nitric acid. Though, I suppose alkali extraction followed by treatment with nitric acid would yield the same results.

[Edited on 5-22-2008 by ShadowWarrior4444]




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MagicJigPipe
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[*] posted on 21-5-2008 at 23:41


I've been reading up on the separation of U isotopes. It just doesn't seem like it would be that difficult for a nation with a decent amount of U ore, modern industrial capacity and some sort of scientific knowledge base.

What am I missing here? What is it exactly that makes extraction of U235 from U ore so difficult for a decent sized nation (other than the hazards of working with fluorine if the gas centrifuge method is used)?

I know it is certainly not something that would be practical to undertake in a home laboratory (if it was someone surely would have done it by now) but if you had the right equipment, knowledge, scientists, ore and industrial capacity, what the hell would stop you (other than money)?

Could someone kindly explain this to me as I'm sure it's not that simple...

Also, I wonder if there is actually some law that says you can't "purify" uranium ore. Or perhaps one that says you cannot extract U235 from mixed isotopes. It wouldn't surprise me.




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[*] posted on 22-5-2008 at 01:28


I may be wrong, but I seem to remember that Most of laws re Isotopes are related to Transuranics, even taking the Am241 chip out of a smoke detector is breaking the law.

some laws about shipment, you can send ore and specimens overseas but need special shipment/licensing to send it in a state Not found in nature.




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[*] posted on 22-5-2008 at 07:48


It would be quite easy, too, to suspect that both detrimental qualities, once they're within living tissue, interact synergistically.

P
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[*] posted on 22-5-2008 at 09:35


It seems that the nitric acid extraction would be the easiest. I will still try these four methods-

-Acidify the peroxide solution to 2.5.
-Add peroxide to the sulfuric acid mix, once neutalized to 2.5.
Both will hopefully precipitate Uranium peroxide.

-Add ammonia to the hydroxide solution, if it extracts anything, to precipitate the diuranate.
-Simple extration with nitric acid.

I also have an idea for producing pure uranium metal using the iodide.



The uranium iodide(dark red) sublimes onto the tungsten wire(red), decomposes to uranium metal, flows down to the other end, drips off, and cools. Not only is temperature control not needed, but neither is voltage control, as the uranium metal never builds up on the wire.

It would be performed in a clay or metal vessel, under helium gas. (what I always use for inert atmospheres) If one had the equipment, would performing this under a vacuum help?

[Edited on 22-5-2008 by StevenRS]
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[*] posted on 22-5-2008 at 11:27


http://carlwillis.wordpress.com/2008/02/20/uranium-chemistry...

I think this will solve most (if not all) of your problems regarding the home chemistry of uranium :) Hope it helps :)
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[*] posted on 22-5-2008 at 13:51


It seems that the best way to extract pure uranium from a leachate is the very selective precipitation of uranium peroxide. Using the carbonate leaching method, all one would have to do is add hydrogen peroxide to the leachate to precipitate uranium. Easy enough. (Back to the method I wanted to avoid) Someone correct me if this would not work.

Now it gets even better. Maybe the uranium peroxide could be used to oxidize a iodide to iodine, which would react with uranium in situ to create uranium iodide for further purification without need for another oxidizer.



[Edited on 22-5-2008 by StevenRS]
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[*] posted on 22-5-2008 at 15:11


I don't think uranYL peroxide is in fact an oxidizer, or, if it is, it's a very weak one. since it can be made with H2O2 (a weak oxidizer).
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[*] posted on 22-5-2008 at 15:41


Yea, I think so. But even if it is very weak, iodide is very easily oxidized. I will just have to test it.

By the way, great link Ragnarok. It led me to this, another good link.

http://www.geocities.com/norm_alara/
A simple sulfuric acid extraction.
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[*] posted on 22-5-2008 at 15:54


Quite nice first post, don't ya think ?
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[*] posted on 23-5-2008 at 05:48


Quote:
Originally posted by StevenRS
...

I also have an idea for producing pure uranium metal using the iodide.



The uranium iodide(dark red) sublimes onto the tungsten wire(red), decomposes to uranium metal, flows down to the other end, drips off, and cools. Not only is temperature control not needed, but neither is voltage control, as the uranium metal never builds up on the wire.
...

[Edited on 22-5-2008 by StevenRS]


I think you'd have to do that under reduced pressure, as UI4 starts to beak down into low volatility lower iodides before it boils.

You'll also have to chill the uranium, as when warm it will react with free I2.

I suspect the molten salt electrolysis is a better route to the metal on a small scale.

BTW - you may need more steps in your isolation process, to get of other metals that follow uranium through one stage or another. You're starting with ore, there's a lot of crap in there to separate out.
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[*] posted on 23-5-2008 at 08:15


Hmm... Maybe one could use the other lower volatility iodides to you advantage? Would they decompose on contact with the tungsten filament?
Using the other iodides would allow for lower temperatures, possibly.

Cooling one end would not be a problem, just submerse it in sand or give it an aluminum heat sink. If even greater cooling is needed, use water.



I do not think that many other metal peroxides are entirely insoluble, most just decompose anyway.

I wonder how selective the peroxide precipitation is?

(does anyone know how to make the picture bigger?)
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Ragnarok
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[*] posted on 23-5-2008 at 09:38


You could use electrolysis in a mixed media. What I mean by that is having an uranium halide in an aqueous layer with a lead sacrificial anode and sulfuric acid and then having the cathode in an imiscible organic layer (dibutyl ether or C6H5-Cl) loaded with LiPF6 or LiClO4 for conductivity. The uranium cations would pass from one layer to the other and they would be deposited from the organic layer as uranium metal, not low valence oxides.
This is just theory, as i haven't tried it. I am going to try it on lithium for my bachelors' degree disertation, and I could give you the results in a few months. A complexation equilibrium in the organic phase would help a lot with cation transfer.
My 2 cents, hope it helps.
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ShadowWarrior4444
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[*] posted on 23-5-2008 at 11:20


Quote:
Originally posted by Ragnarok
You could use electrolysis in a mixed media. What I mean by that is having an uranium halide in an aqueous layer with a lead sacrificial anode and sulfuric acid and then having the cathode in an imiscible organic layer (dibutyl ether or C6H5-Cl) loaded with LiPF6 or LiClO4 for conductivity. The uranium cations would pass from one layer to the other and they would be deposited from the organic layer as uranium metal, not low valence oxides.
This is just theory, as i haven't tried it. I am going to try it on lithium for my bachelors' degree disertation, and I could give you the results in a few months. A complexation equilibrium in the organic phase would help a lot with cation transfer.
My 2 cents, hope it helps.



Could a membrane cell be constructed using an organic solvent in the cathode chamber and water in the anode chamber? Naturally this might only be suitable for producing alkali metals; uranium will not pass the membrane. Though, using a membrane cell would avoid any difficulty extracting the alkali metal.

The above article mentioned forming UO2 coatings via electrolysis using 12 volts and patience; this suggests that any solvent electrolysis would likely have poor yields. The molten salt electrolysis might show promise though! Molten salts of uranium are already part of a newer power-plant design: http://en.wikipedia.org/wiki/Molten_salt_reactor. Electrolysis of a fluoride is not to be recommended, especially of the molten variety; Uranium Tetrachloride is produced from carbon tetrachloride and UO2 industrially, though there may be a more useful way involving chlorine gas. The gas could be recovered as part of the electrolytic process.

[Edited on 5-23-2008 by ShadowWarrior4444]




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Ragnarok
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[*] posted on 23-5-2008 at 12:13


If you have two immiscible solvents, the interface acts like a membrane. Hidrophobic ions like tertaphenyborate will stick in the organic phase, while hidrophylic ions like OH(-) or NH4(+) will tend to stick to the water phase. Anyway, is you start from UF4 you can form complexes of U(IV) in the organic phase that don't have water as a ligand and are less likely to deposit as UO2. The reduction of H(+) can not happen, since the proton complex [(H2O)3.H2O](+) can not pass into the organic layer. If you use ethers, you can go as low as to be able to deposit metallic lithium; uranium would not be such a problem.
The acidic solution redox potentials that may be involved are:
UO2(2+) -> UO2(+) +0.17 V (6+)->(5+)
UO2(+) -> U(4+) +0.38 V (5+)->(4+)
UO2(2+) -> U(4+) +0.27 V (6+)->(4+)
U(4+) -> U(3+) -0.52 V
U(3+) -> U(2+) -4.70 V
U(2+) -> U(0) -0.10 V
U(4+) -> U(0) -1.38 V
(data from Shriver, Atkins, Langford - Inorganic Chemistry and referenced to the standard hidrogen electrode)

Acording to the redox potentials, the only things that are possible in water are reductions from 6+ to 4+ and maybe the reduction from 2+ to uranium metal.

In an organic phase, the -1.38 potential of the3 last redox couple is no big deal. So either use a salt that is slightly soluble in an organic phase or a two phase system.
Making a complex of uranium with a highly organic-soluble ligand in the organic phase would be a huge help to the ion transfer between phases. Maybe a LIX ligand that is used in two-phase extreaction from sulfuric leachate?
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[*] posted on 23-5-2008 at 17:19


I wonder if simple decomposition of a uranium halide would work? Just have some uranium iodide in a metal tube, get it hot enough to decompose the iodide, and then blow an inert gas though the tube to drive off the halogen, driving the equilibrium to the right?
UI<sub>4</sub> <--> U + 2I<sub>2</sub>

You could then condense the iodine vapor for reuse, or just bubble it through uranyl tricarbonate to form more iodide (and iodate).



(How do I made these bigger!) The inert gas goes in the little tube, down to the heated iodide, and then out the big tube.
Easy, simple, and produces liquid uranium metal!

The only problem I see is that the uranium iodide might just evaporate off before decomposing, maybe a different salt that decomposes to U metal more easily without evaporating/sublimating first could be used?

[Edited on 23-5-2008 by StevenRS]
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[*] posted on 23-5-2008 at 17:30


Uranium halides will evaporate before they thermally decompose.
However, an arkel-de-boer process could be used on the uranium halide vapour.




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[*] posted on 23-5-2008 at 17:52


Maybe it would be able to catalytically decompose the iodides at a lower temperature?
But if not, the arkel-de-boer process does not seem so hard to do.
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[*] posted on 23-5-2008 at 23:28


Quote:
Originally posted by StevenRS
I wonder if simple decomposition of a uranium halide would work? Just have some uranium iodide in a metal tube, get it hot enough to decompose the iodide, and then blow an inert gas though the tube to drive off the halogen, driving the equilibrium to the right?
UI<sub>4</sub> <--> U + 2I<sub>2</sub>

You could then condense the iodine vapor for reuse, or just bubble it through uranyl tricarbonate to form more iodide (and iodate).



Um, no. UI4 sublimes to some extent, and at the same time decomposes into lower iodides.

Whatever container you hope to do this in will have to be inert to hot I2, UI4, and uranium.

Running it at a temperature much below the melting point of uranium would very likely result in the formation of finely powdered uranium, ready to ignite on contact with O2.

uranyl tricarbonate = [(UO2)(CO3)3] 4- iodine is not to give UI4 from that, way too much oxygen in it. Nor can you form the iodide around much water.
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[*] posted on 23-5-2008 at 23:32


this is just a thought, but after reading uranium halides in here and wanting to reduce these to the metal, I was thinking about Uranotypes and photography developer, I wonder if it would be possible to use such things as Phenidone, Metol, Hydroquinone or the likes to reduce the halide to the metal?

Woelen would be the best one to ask about this method.




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[*] posted on 24-5-2008 at 00:32


Uranium is sort of between manganese and magnesium in the electromotive series, similar to some of the lanthanum group, in the range where the metal reacts with water to create bubbles. The relatively low potentials of photographic developers just isn't enough to do the job.
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[*] posted on 24-5-2008 at 01:04


I did think about the water aspect, and considered an organic solvent instead, and from some research I did with regard to photography, these developers although often used in a Basic soln to make the react faster (within seconds in the darkroom), will also work as they are without making it basic, it just takes a lot longer.

but again, the problem would be that if it did work, you would still only end up with very fine pyrophoric powder :(




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[*] posted on 24-5-2008 at 05:55


Maybe sintering the powder in an inert atmosphere? Although I think it would react with N2. Maybe a H2 or Ar atmosphere? It would look and feel like bulk metal, only it's density and mechanical resistances would be smaller by a few %.
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[*] posted on 24-5-2008 at 06:57


The electrodeposition
of uranium has been performed at
aluminum cylinder cathodes from uranyl
nitrate/isopropyl alcohol solutions.
Uranium metal can be prepared from
a combination electrodeposition/thermal
decomposition process by first forming
a mercury amalgam and subsequently
heating the amalgam to produce pure
uranium metal. Hasegawa et al. utilized
this methodology to prepare uranium
metal with purity higher than
a commercial grade of approximately
99.95%. The electrochemical cell consisted
of anode and cathode compartments
separated by a proton-specific cationexchange membrane. The platinum anode
compartment was filled with 1 M sulfuric
acid, while the cathode compartment
contained 0.5–1.0 M HCl solution. The
initial step in the process was reduction
of U(VI) to U(IV) at −0.6 V versus
SCE, followed by pH adjustment with an
acetic acid/sodium acetate solution, and
finally amalgamation from −2.0 to −2.3 V
versus SCE. Most of the mercury was removed
from the amalgam at 250 ◦C in
a vacuum (<1 × 10−6 torr) before heating
to 1200–1300 ◦C for 1h. Martinot
and coauthors have reported the electrodeposition
of uranium metal from an
organic solvent medium [72]. The report
mostly focuses on La metal electrodeposition,
but the conclusion with regard
to uranium is that macroscopic quantities
of metal can be deposited from γ -
butyrolactone/tetrahydrofuran (60/40 vol
%) solutions. The current density at the
tungsten working electrode surface must
be set between 20 and 40 mA cm−2 for
plating to occur. Since reduction of the
solvent is a competing process, setting the
current density too high results in the inhibition
of the plating of uranium. Results
from Inductively Coupled Plasma (ICP)
analyses of the dissolved metal were used
to calculate the faradaic yield (about 39%
at 20 mA cm−2).


Theese are straight from an encyclopedia of electrochemistry.
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[*] posted on 24-5-2008 at 07:05


I suspect any handling of U metal will be accompanied by many of the difficulties of Ti, i.e., B, C, N, O, P, S and so on will all cause trouble.

Uranium strikes me as a softer, lower melting metal, so maybe it behaves better. Still, I believe it forms extensive intermetallics with most metals, similar to the rare earths. I don't have many phase diagrams with it, unfortunately.

Tim




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