Uranium Isolation

Y'think soaking it in molten sodium chlorate would do it?

...Hey, I have a lot of sodium chlorate on hand, okay?...

Tim

which means it can be major hassles to filter, and not likely to be formed from carbonate complexes.

Uranyl iodate has a low solubility, as does the sulfide and oxalate; the oxalate forms soluble complexes while the sulfide dissolves in dilute acids and is readily oxidised by air.

The extraction, ion exchange, and volatility methods tend to be industrial type ones, unless you're set up to work with UF6. I think you'll have better luck looking in old chemistry texts from before WW-II into the Victorian era; they use less exotic reagents and conditions.

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



[Edited on 21-5-2008 by The_Davster]

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 [Edited on 21-5-2008 by The_Davster]

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]

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

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]

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?)

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?
UI₄ <--> U + 2I₂

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]

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

If someone here has the heat of formation (@ 298 K) of UCl₄, then it would be possible to estimate the end-temperature of the reaction in adiabatic conditions. With a boiling point of 1,412 C (for MgCl₂ the reduction reaction may well be energetic enough to cause the MgCl2 to sublimate off, leaving the unprotected, very hot metal exposed to air. This could explain what happened in your case. In many reductions with Al or Mg, unless the slag sufficiently protects the newly formed metal from the air, inert atmosphere will be necessary to obtain the clean, unoxidised product.

Great!

From this paper the standard molar enthalpy of reaction at 298 K for U (s) + 2 Cl₂ (g) ---> UCl₄ (s) to be ΔH = - 1,018 kJ (per mol of UCl₄ (see page 5, Table 3) can be gleaned.

For UCl₄ (s) + 2 Mg (s) ---> U (s) + 2 MgCl₂ (s) we then obtain ΔH = - 264 kJ (per mol of UCl₄. That's only barely exothermic and almost certainly insufficient to melt both the U and the MgCl₂. To be confirmed/infirmed tomorrow...

Very interesting description of the UF₄/Mg reduction process indeed! But fluorination of U is probably outside the capability envelope of most backyard scientists and the UCl₄-route looks more promising from that perspective, so it's a case of finding the most suitable reductant. Ideally (for backyard science), the RT UCl₄/reductant mixture would be lit with an ignition pill and the temperature developed during reduction would be high enough to melt both the metal and the slag, leading to gravitational separation of the (in this case very heavy) metal from the molten chloride/U metal mix.

Al itself can be excluded as UCl₄ + 4/3 Al ---> U + 4/3 AlCl₃ is not exothermic (ΔH = + 77 kJ at 298 K).

The main reductants to be considered are Mg and K, IMHO (at least at first glance).

The enthalpy needed to heat 1 mol of U from 298 K to its MP (1,405 K) is approximately 39 kJ (including heat of fusion).

The enthalpy needed to heat 2 moles of MgCl₂ from 298 K to 1,405 K (including heat of fusion at MP = 987 K) is approximately 227 kJ.and for UCl₄ + 2 Mg ---> U + 2 MgCl₂ the heat of reaction (at 298 K) is about - 266 kJ (per mol of U). So basically (by sheer coincidence) 39 kJ + 227 kJ = 266 kJ. The Mg reduction, started from 298 K and carried out in adiabatic conditions would probably just about reach the MP of uranium. In reality we're cutting it fine and heat losses will probably lead to slightly lower end-temperature and thus poor metal/slag separation.

The enthalpy needed to heat 4 moles of KCl from 298 K to 1,405 K (including heat of fusion at MP = 1,049 K) is approximately 364 kJ.and for UCl₄ + 4 K ---> U + 4 KCl the heat of reaction (at 298 K) is - 730 kJ. Since as 730 kJ > 39 kJ + 364 kJ (730 kJ > 403 kJ), the reduction with K has kilojoules to spare and should heat to well past the MP of U. Perhaps this partly explains why in1841, Eugène-Melchior Péligot, who was Professor of Analytical Chemistry at the Conservatoire National des Arts et Métiers (Central School of Arts and Manufactures) in Paris, isolated the first sample of uranium metal by heating uranium tetrachloride with potassium.

Interestingly, the results with Li are very similar to those with K: heat of reaction (at 298 K) is about - 614 kJ, the enthalpy to heat 4 moles of LiCl from 298 to 1,405 K (including heat of fusion at MP = 881 K) is about 334 kJ, so 614 kJ > 39 kJ + 334 kJ (614 kJ > 373 kJ). This too should run to well past the MP of U, at least in adiabatic conditions.

Li being easier to handle than K, it would be my own choice of reductant, if I possessed some UCl₄.

Li is of course more difficult to obtain/produce in fine form than Mg and a mixture (layered perhaps?) of UCl₄ and coarse Li may have to be furnace heated in order to trigger the reduction reaction... But both powdered and pelletised lithium are commercially available (but maybe not to backyarders...)

[Edited on 19-6-2008 by blogfast25]

[Edited on 19-6-2008 by blogfast25]

[Edited on 20-6-2008 by blogfast25]