Sciencemadness Discussion Board

Uranium Isolation

StevenRS - 18-5-2008 at 16:27

I have recently acquired a Geiger counter, and have found some "hot rocks", some reading over 19,000 CPM. I think they are carnotite, and some have green crystals with black flecks, possibly uraninite.

It could be and interesting process to purify this ore to actual metal, or even a salt. I was wondering what methods could be used for maximum uranium extraction, and then a method to extract the uranium salts from the iron and other metal salts sure to be present.

I have heard of the alkaline extraction method, but this produces low yields. The industry uses it because it is cheap.

Also, most uranium salts are soluble, including the carbonate, so maybe converting the mixed salts to carbonate, filtering out the other insoluble salt out would be possible as a method of purification.

Converting this carbonate (or any other salt) to an oxide would be easy, but converting U<sub>3</sub>O<sub>8</sub> to U could be difficult.

Any thoughts?

crazyboy - 18-5-2008 at 16:34

http://www.unitednuclear.com/extract.htm


The second page is missing maybe someone could fill it in for him?

[Edited on 18-5-2008 by crazyboy]

StevenRS - 18-5-2008 at 16:48

I have read that page, and it is informative, but the alkaline method produces low yields and with my small supply of ore this is not ideal.

Would a stronger base, KOH maybe, work better?

ShadowWarrior4444 - 18-5-2008 at 20:00

Quote:
Originally posted by crazyboy
http://www.unitednuclear.com/extract.htm


The second page is missing maybe someone could fill it in for him?

[Edited on 18-5-2008 by crazyboy]


I found it amusing that on the solvent extraction page for uranium they claim that they haven’t been able to update it because the process requires sulfuric acid--which they apparently 'haven’t gotten in yet.' *raises an eyebrow*

I personally don’t like united nuclear, not just for their prices.

That said, Uranyl Hydroxide does exist, though it is commonly precipitated from oxidized uranium at neutral pH. Certain papers describe the volatilization of uranium under strongly oxidizing conditions, which may allow its distillation:
http://dx.doi.org/10.1016/j.jnucmat.2005.07.013

Uranyl Chloride may be an option for extraction as well, it forms from uranium dioxide at red heat under chlorine gas, or dissolving uranium oxide in HCl. It is unstable, and decomposed by light, which may facilitate recovery:

The company Indian Rare Earths Limited (IREL) has developed a process to extract uranium from the Western and Eastern coastal dune sands of India. After pre-processing with high intensity magnetic separators and fine grinding, the mineral sands (known as monazite), are digested with caustic soda at about 120C and water. The hydroxide concentrate is further digested with concentrated hydrochloric acid to solubilise all hydroxides to form a feed solution composed of chlorides of uranium and other rare earth elements including thorium. The solution is subjected to solvent extraction with dual solvent systems to produce uranyl chloride and thorium oxalate. The crude uranyl chloride solution is subsequently refined to nuclear grade ammonium diuranate by a purification process involving precipitation and solvent extraction in a nitrate media.

This may be the best solution, as it will also recover other interesting metals and is likely the extraction of choice when you are not positive it is uranium causing the radioactivity. I should note that Uranyl Chloride is reported as being toxic to the point of humor if ingested.

Another interesting note:
It is precipitated during production by adding aqueous ammonium hydroxide after uranium extraction by tertiary amines in an organic kerosene solvent. This precipitate is then thickened and centrifuged before being calcined to uranium oxide. Canadian practice favours the production of uranium oxide from ammonium diuranate, rather than from uranyl nitrate as is the case elsewhere.

There is also, of course, digestion with nitric acid, as the united nuclear site mentions.

Ancillary:
http://en.wikipedia.org/wiki/In-situ_leaching
http://en.wikipedia.org/wiki/Uranium_mining

12AX7 - 18-5-2008 at 20:06

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

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

Tim

pantone159 - 18-5-2008 at 20:37

Lee, Concise Inorg. Chem., 5th ed, has this to say about industrial U extraction... paraphrasing...

Ore is crushed and concentrated via flotation. Then it is roasted in air, and leached with H2SO4 in the presence of MnO2 to ensure all U goes as U(VI). This is precipitated as Na diuranate ('yellowcake'). This dissolves in HNO3 as uranyl nitrate, and can be extracted into tributyl phosphate (c. 20% accd another source) in kerosene.

I wonder how available tributyl phosphate is.

Quote:
Originally posted by ShadowWarrior4444
I personally don’t like united nuclear, not just for their prices.


I wonder why? It is true, the guy behind UN is definitely a kook, he is the one responsible for the alien dissection movie myth and a lot of the Area 51 nonsense. But, they do have some interesting stuff, and not all of their prices are bad, I guess it depends on what. Given the limited number of hobbyist suppliers, I am willing to tolerate eccentricity.

MagicJigPipe - 18-5-2008 at 21:16

It seems like tributyl phosphate would be an ester of butanol and phosphoric acid.

Seems like it wouldn't be exceedingly difficult to make.

Wikipedia has this to say (just remember to take it with a grain of salt):
Quote:

Tributyl phosphate is manufactured by esterification of orthophosphoric acid with butyl alcohol. A laboratory synthesis proceeds with phosphorus oxychloride: [1]

POCl3 + 3 C4H9OH → PO(OC4H9)3 + 3 HCl

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


OrgSyn citation on Wikipedia page:

^ G. R. Dutton and C. R. Noller (1943). "n-Butyl phosphate". Org. Synth.; Coll. Vol. 2: 109.
http://www.orgsyn.org/orgsyn/prep.asp?prep=cv2p0109

That sucks... It would probably be easier to find and buy tributyl phosphate (TBP) than to find or make phosphorus oxychloride. I'm sure there are other ways, however...

[Edited on 5-18-2008 by MagicJigPipe]

YT2095 - 18-5-2008 at 23:02

Quote:
Originally posted by pantone159
Lee, Concise Inorg. Chem., 5th ed, has this to say about industrial U extraction... paraphrasing...

Ore is crushed and concentrated via flotation. Then it is roasted in air, and leached with H2SO4 in the presence of MnO2 to ensure all U goes as U(VI). This is precipitated as Na diuranate ('yellowcake'). This dissolves in HNO3 as uranyl nitrate,


once you get to this stage, make a ppt with oxalic acid, as uranyl oxalate is not very soluble at all, then heat the oxalate to decomp into uranium oxide, then use as thermit to form the metal.

not_important - 19-5-2008 at 01:08

Some older processes, more focussed on radium, would crush and finely grind the ore, air roast it to oxidise sulfides and remove arsenic, then fuse it with potassium or sodium bisulfate (believe you can use equal molar mix of K/Na sulfate and ammonium sulfate). Crush the cooled melt, extract with warm dilute H2SO4, filter the extract and add an excess of Na2CO3 to precipitate the iron group, aluminium, any remaining calcium, and a few other metals. Concentrate, but not so far as to cause crystals or ppt to form, add strong NH3 + NH4Cl to remove vanadium as NH4VO4. Filter, boil the solution to get a ppt of ammonium diuranate (NH4)2U2O7; igniting this in air at ~500 C gives UO3.

Additional purification would involve reactions as suggested above. Another useful purification step is the precipitation of uranyl peroxide at a pH of 3-4 using 35 to 50 percent H2O2.

roamingnome - 19-5-2008 at 17:16

Also look into TRU resin it catches 3+ Actinides

of course thorium removal takes HF acid... and they are slow

What would you like to do with your nuclear material?
A nuclear battery wold be just dandy


SO... i really encourage you to remember Paul Brown!!!!!!!

http://users.erols.com/iri/Pauleulogy.htm

http://www.rhfweb.com/paulresnuc.htm



ive read about his invention and i think the concept is quite valid

think nuclear powered tesla coil.


actually
http://www.dow.com/liquidseps/prod/pt_u.htm

these resins are more for Uranium

[Edited on 20-5-2008 by roamingnome]

Fleaker - 19-5-2008 at 17:31

Anyone have a proposal for uranyl acetate to uranium metal without ridiculous reducing agents (Li) and dangerous gases?

The_Davster - 19-5-2008 at 18:15

Perhaps http://www.sciencemadness.org/lanl1_a/lib-www/la-pubs/003180... could be modified for the acetate?
Of course, one would have to take precautions to avoid uranium vapours from aerosols and non-zero vapour pressures of uranium compounds.

not_important - 19-5-2008 at 22:49

Uranium is difficult to reduce. The lower oxides are rather stable, the element combines with many non- and semi- metals so it tends to form carbides UC2 mp 2400 C), nitrides (U3N4 decomposes in vac ~1400 C), &ct.

Uranyl acetate decomposes below 300 C, the carbon and oxygen content also make it an unlikely direct precursor.

The methods I've seen listed are basically reduction of U3O8 with Mg in an arc furnace in H2, the action of K, Na, or Na-Mg allow on UCl4 or UCl4.2NaCl at 700 to 800 C; electrolysis of UCl4 or UCl3 in fused alkali or alkali+alkaline earth chlorides; electrolysis of aqueous solution of UCl4 with a mercury cathode under an atmosphere of hydrogen followed by distillation of the mercury; and the Ames process - effectively a thermit reaction using calcium or magnesium. All but the fused salts electrolysis and Ames process give a pyrophoric powdered metal.

StevenRS - 20-5-2008 at 11:48

I have heard that uranium oxide(s) can be reduced using aluminum and magnesium at high temperature. Sound fun. I wonder how high of a temp?

Maybe someone could compile a list of common, easily formed, insoluble uranium salts? I am pretty sure it will be a short list...

I think Ammonium uranyl carbonate is insoluble. Could one leach the uranium form an ore with a carbonate/bicarbonate, and then convert the resulting Uranyl Tricarbonate to the ammonium salt with ammonia gas?

One could convert the tricarbonate to the sulfate, and extract that, as I believe it is also insoluble.

Is the hydroxide soluble in basic solutions? If so, that opens up a possible route of extraction from the ore.

Edit: Found this on wikipedia--
"In the classical procedure for extracting uranium, pitchblende is broken up and mixed with sulfuric and nitric acids. The uranium dissolves to form uranyl sulfate, and sodium hydroxide is added to make the uranium precipitate as sodium diuranate, This older method of extracting uranium from its uraninite ores has been replaced in current practice by such procedures as solvent extraction, ion exchange, and volatility methods."

Seems easy enough. Would the sulfuric acid method work on other ores than urananite?



[Edited on 20-5-2008 by StevenRS]

not_important - 20-5-2008 at 22:02

The reduction of uranium oxides does not appear to be a very desirable route. Historically such efforts tended to give the lower oxides or mixtures of uranium with oxides, nitrides, or carbides. Such methods also tend to produce the uranium as a pyrophoric finely powdered metal.

They have been done. One early old process was to mix "uranium oxide" with an excess of powdered magnesium, compact that into rods, then place the rods in a special arc furnace in an atmosphere of hydrogen. After cooling the hydrogen was flushed out with CO2, the finely powdered uranium was then washed with dilute acetic acid to remove excess magnesium.

If you have carnotite, you have vanadium which tends to follow uranium around to some extent. Removing the vanadium as ammonium metavanadate, or uranium as peroxide, are two methods of getting good separation, I believe. Thorium and some cerium will go with uranium with the peroxide, but they are easy to separate at other stages.

Extraction of ores with alkalies tends to drag along silica, which must be separated at a later stage; repeated fuming down with H2SO4, then dissolving the metal ions out with acid, is a common way of doing that.

Changing the oxidation state of the uranium and precipitating it or impurities, followed by another change in oxidation state to switch solubility, seemed to be common. A solution of a uranyl salt in water with alcohol or glucose, exposed to sunlight will slowly form an insoluble violet to black hydrate oxide of mixed oxidation states; air oxidises it to UO3.

Quote:
All uranates, even those of the alkalies and ammonium, are insoluble. Ammonia will precipitate from a solution of a uranyl salt the di-uranate (NH)2[U2O7], a yellow powder practically insoluble in water, but easily soluble in ammonium carbonate solution; this salt can be used to estimate uranium, but is liable to go over into the colloidal state

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.

Rideal_furnace_for_uranium_metal.png - 38kB

pantone159 - 20-5-2008 at 22:15

btw - Lee also added re U metal prepn:

Final steps are conversion to UF4 followed by reduction with Ca or Mg to yield metal U.

Even if U metal is too hard, just getting any purified U compound from ore would be cool.

StevenRS - 21-5-2008 at 03:46

This might be an easy method of getting pure metal, and it only requires a lightbulb! (kinda)
Place uranium iodide in close contact with a white hot tungsten filament, and the iodide sublimates, contacts the wire, and is decomposed, yielding very pure metal, minus the tungsten. As the amount of deposited metal increases, the resistance goes down, so more current must be applied.

not_important - 21-5-2008 at 05:34

I knew this was familiar

http://www.sciencemadness.org/talk/viewthread.php?tid=6714&a...

And it appears the van Arkel and de Boer method is a little tricky when used with uranium, the decomposition temperature of the iodides is high enough that it must be done near the melting point of the metal, so careful temperature control must be done.

http://jphyscol.journaldephysique.org/index.php?option=artic...

van Arkel - de Boer generally uses a small amount if the iodide, and a lot of the free metal in powdered form, in an evacuated envelope along with the heated filament. The powdered metal is kept hot enough to react with free iodine, which is formed when the iodide decomposes on the filament. If just the iodide is used, the iodine formed increases the pressure, both conducting heat away from the filament and reversing the decomposition. With some elements a mixture of a volatile halide and hydrogen is passed over the hot filament in a flow-through system.


And this could be useful in checking solubilities :

http://www.archive.org/details/dictionaryofchem014380mbp

Pulverulescent - 21-5-2008 at 07:20

Quote:
Originally posted by StevenRS
It could be and interesting process to purify this ore to actual metal, or even a salt.

I was wondering what methods could be used for maximum uranium extraction.

Any thoughts?


Here's one!

The most interesting thing about uranium besides its extreme toxicity is its radioactivity.

However, if you can get your hands on enough of the 235 isotope things'll get really really interesting.

And some guy, Osama bin. . .something might just be interested, too.

You could call him. . .

P

StevenRS - 21-5-2008 at 08:43

I finally got my ore ground to powder, I mean dust. The stuff flows like a liquid, three days in a ball mill with 3/4 chrome steel media will do that. I liked 3 extractiom methods, the peroxide extraction, sodium hydroxide, and sulfuric acid. The peroxide has already reacted a bit, and the solution is a bit yellow. L

ShadowWarrior4444 - 21-5-2008 at 11:06

Quote:
Originally posted by Pulverulescent
Here's one!

The most interesting thing about uranium besides its extreme toxicity is its radioactivity.

However, if you can get your hands on enough of the 235 isotope things'll get really really interesting.

And some guy, Osama bin. . .something might just be interested, too.

You could call him. . .

P


Not all radioactive substances are equally detrimental to life, this is a common misconception. Uranium emits Alpha radiation, which cannot pass through solid matter nor a few meters of air. Alpha radiation has not ever been known to cause a human cancer, as well. This is likely because it is in the form of a helium nucleus, whereas the other forms of radiation are subatomic, and therefore more able to penetrate shielding and damage cells. Uranium can be effectively shielded by a piece of paper.

And the term "extreme toxicity" I reserve for things like Cyanogen Bromide, Dimethyl Mercury, and on a bad day Nickel Carbonyl.

Osama would not know what to do with raw 235; now Iran, the ayatollahs would pay quite a sum for the services of an enrichment expert. (As they have for Russian ex-patriots with *any* knowledge of nuclear physics.)

Quote:
I finally got my ore ground to powder, I mean dust. The stuff flows like a liquid, three days in a ball mill with 3/4 chrome steel media will do that. I liked 3 extraction methods, the peroxide extraction, sodium hydroxide, and sulfuric acid. The peroxide has already reacted a bit, and the solution is a bit yellow. L



I'm quite interested in seeing how the peroxide extraction turns out, as of the three you've chosen, it is the only one that is not used appreciably on an industrial level. Can you photograph each stage of the process? [This might make a good addition to the prepublication section--perhaps a treatise on OTC ore extraction processes.]

[Edited on 5-21-2008 by ShadowWarrior4444]

StevenRS - 21-5-2008 at 18:43

I had to go out of town, so my test samples are 500 miles away, butt I will definatly take pictures when I return.

What else would the peroxide method extract, other than cerium or thorium?

I wonder what could be used to precipitate the uranium out of this solution?
Maybe conversion to the sulfate? Or the carbonate, followed by neutralization?

I just read about a very efficent extraction process using sulfuric acid and an oxidiser, sodium chlorate. Maybe H2O2 could be substituted?

"To produce this form of yellowcake (magnesium diuranate), crushed ore is mixed with hot water to a 58% solids slurry. The solids slurry is then processed through a series of tanks, where sulfuric acid, sodium chlorate, and steam are used to extract the uranium from the solids slurry. The average leaching efficiency for this process is 98.5%. The uranium-bearing solution is then decanted and directed to a solvent extraction (SX) process for further purification. In this extraction step, the dissolved uranium is transferred from the feed solution into the organic solvent. Next a stripping step recovers the uranium into a sodium chloride aqueous phase after which the barren solvent is recycled. The average efficiency of the SX circuit is 99.9%. The high-grade “pregnant” strip solution from SX goes to the next stage where magnesia slurry is added to precipitate magnesium diuranate."

It does sound a bit compicated, though. I would prefer something that did not use an organic solvent extraction phase, and is less steps.



[Edited on 21-5-2008 by StevenRS]

Polverone - 21-5-2008 at 19:10

Quote:
Originally posted by ShadowWarrior4444
Not all radioactive substances are equally detrimental to life, this is a common misconception. Uranium emits Alpha radiation, which cannot pass through solid matter nor a few meters of air. Alpha radiation has not ever been known to cause a human cancer, as well. This is likely because it is in the form of a helium nucleus, whereas the other forms of radiation are subatomic, and therefore more able to penetrate shielding and damage cells.

Uranium is relatively harmless because it has such a long half-life. Alpha emitters are the least dangerous materials externally but among the most damaging inside the body. The relatively high mass and charge of an alpha particle leads to high linear energy transfer which in turn leads to concentrated damage on a cellular level. The famous alpha emitter radium contributed to many human cancers before its properties were well-understood. The Russian dissident Alexander Litvinenko was killed in a high profile incident only 2 years ago with the alpha emitter Polonium-210. Alpha emitters as a group are certainly not a safer category of radioactive materials.

The_Davster - 21-5-2008 at 19:15

Quote:
Originally posted by ShadowWarrior4444
Alpha radiation has not ever been known to cause a human cancer, as well. This is likely because it is in the form of a helium nucleus, whereas the other forms of radiation are subatomic, and therefore more able to penetrate shielding and damage cells. Uranium can be effectively shielded by a piece of paper.

And the term "extreme toxicity" I reserve for things like Cyanogen Bromide, Dimethyl Mercury, and on a bad day Nickel Carbonyl.



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]

not_important - 21-5-2008 at 20:37

Quote:
Originally posted by StevenRS

I wonder what could be used to precipitate the uranium out of this solution?
Maybe conversion to the sulfate? Or the carbonate, followed by neutralization?

I just read about a very efficent extraction process using sulfuric acid and an oxidiser, sodium chlorate. Maybe H2O2 could be substituted?
...


We've been listing those sorts of things. The sulfate, unlike simple sulfates such as those of lead or barium, is a basic sulfate UOSO4 . 2H2O formed under proper conditions of pH and concentration. The simple sulfates, such as U(SO4)2 . n H2O or UO2SO4 . n H2O, are soluble in water.

The Radiochemistry of Uranium
http://library.lanl.gov/cgi-bin/getfile?rc000030.pdf
http://www.sciencemadness.org/lanl1_a/lib-www/books/rc000068...
has some useful information:

Quote:
Peruranates are formed when uranyl solutions containing hydrogen peroxide are made alkaline. The composition of the peruranates depends upon the concentration of the alkali and peroxide. The followlng groupO have been identified: M2U2O10 . XH20, M2U06 . XH20, M6U2013 . XH20, and M4U08 . XH20. The peruranates are generally soluble In water. The least soluble are those of the M2U2010 . XH20 group. The peruranates are soluble In dilute mineral acid.


and
Quote:
Hydrogen peroxide precipitates uranium
peroxide, U04 . XH20, from sllghtly acldlc solutions. The
reaction occurs In the pH range 0.5-3.5. The optimum range
IS 2.0-2.5. Hydrogen Ions released with the formation of
uranlvm peroxide are neutralized with ammonia or ammonium
acetate. Complete preclpltiatlon requires an excess of
hydrogen peroxide. Quantltatilve separation may be effected
by freezing the aolutlon, allowing It to &tend, and filtering
at 2°C. The separation from most elements Is good since It
Is done from an acidic solution. plutonium, thorium,
hafnium, zlrconlum, and vanadium aleo precipitate. Iron
interferes by catalytically decomposing hydrogen peroxide.
Small quantities of Iron may be complexed with acetic, lactlc,
or malonlc acid. Low yields may result from the use of
malonic. acid. Ammonlum, potaaslum, and alkaline earths retard
the rate of precipitation. Complexlng Ions such as
oxalate, tartrate, sulfate, and fluoride In large quant%tiea,
also Interfere. Fluoride Ion may be complexed with alumi-
-.-


Th e extraction processes generally use counterflow extraction, doing them in a home lab means doing a series of extractions and back extractions, often not practical except for those with OCD.

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.

ShadowWarrior4444 - 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]

MagicJigPipe - 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.

YT2095 - 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.

Pulverulescent - 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

StevenRS - 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]

Ragnarok - 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 :)

StevenRS - 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]

Ragnarok - 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).

StevenRS - 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.

Ragnarok - 22-5-2008 at 15:54

Quite nice first post, don't ya think ?

not_important - 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.

StevenRS - 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?)

Ragnarok - 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.

ShadowWarrior4444 - 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]

Ragnarok - 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?

StevenRS - 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]

The_Davster - 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.

StevenRS - 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.

not_important - 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.

YT2095 - 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.

not_important - 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.

YT2095 - 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 :(

Ragnarok - 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 %.

Ragnarok - 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.

12AX7 - 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

StevenRS - 24-5-2008 at 08:55

Uranium actually melts (1132.2 °C) fairly easily. It could be cast under an inert atmosphere without difficulty using a special method were the mold and reservoir are one piece, and both are heated together, never exposing the uranium to the atmosphere.

Cast uranium bullets? That would be very cool.

Simply decomposing the uranium iodide in a tube was to good to be true, just a half-serious idea.

Decomposing the uranium on a tungsten filament hot enough to melt the uranium metal formed seems the way to go. Anyone have the resources to test it with a different metal with similar properties?

This would avoid dealing with pyrophoric powders or fluorine and its compounds.

The electrolysis method sounds pretty good too, it avoids high temperatures, but electrolysis is a slow process, and creating bulk uranium would be difficult.

One more thing, fired clay is resistant to hot halogens, correct?

[Edited on 24-5-2008 by StevenRS]

StevenRS - 29-5-2008 at 08:40

I am finally back in town, and this is my method for pure uranium metal.

4Al + 3UCl4 --> 3U + 4AlCl3

This would be performed in a crucible, partially sealed from the air. When completed, the crucible would be heated enough to drive off the volatile AlCl3, producing pure uranium.

Possibly, an excess of aluminum could be added to form a flux on top of the uranium if they do not alloy.

12AX7 - 29-5-2008 at 10:16

I would think the Kroll process would be quite suitable, giving a bit more reasonable end product, MgCl2, instead of that weird AlCl3. MgCl2 is still volatile enough to drive off the metal, and would be an excellent flux. An MgO crucible and inert gas cover would probably be sufficient. Though UCl4 probably has to be led in slowly, as in the Kroll process proper, to control reaction temperature. Maybe something less reactive, but not too much less reactive, would be suitable. Um, let's see Zn, Al, Mg -- dangit, there aren't many common metals within that narrow range of reduction potentials!

Al and U probably form numerous intermetallics, a shame I don't have the phase diagram. For example, the Al-La system contains an intermetallic with melting point substantially higher than either metal -- maybe not too big a stretch to call that an aluminide, as if. If uranium is more white-metal-like, the Al-Pb system is immiscible with only minor solubility, but I feel that's unlikely.

Tim

not_important - 29-5-2008 at 16:01

Quote:

This would be performed in a crucible, partially sealed from the air. When completed, the crucible would be heated enough to drive off the volatile AlCl3, producing pure uranium.

Possibly, an excess of aluminum could be added to form a flux on top of the uranium if they do not alloy.


I suspect you'll find that anything less than a good seal will reduce if not thwart the effectiveness of this approach. Uranium loves oxygen

Quote:
One more thing, fired clay is resistant to hot halogens, correct?

Not entirely, depends on the clay. More important would be the porosity of the ceramic; much of the non-glazed commercial ceramics are porous enough that they'll bubble when immersed in water and absorb water as well.


Segregation in uranium-aluminum alloys
http://www.ornl.gov/info/reports/1958/3445603505721.pdf


THE ALLOY SYSTEMS URANIUM-ALUMINUM AND URANIUM-IRON
http://www.osti.gov/bridge/servlets/purl/4449905-4O7Sdg/4449...

blogfast25 - 31-5-2008 at 05:12

This patent relies on the reduction of UCl4 with Mg:

Quote:
c. Reduction of Uranium Tetrachloride to Metallic Uranium

The uranium tetrachloride powder formed and collected in the chlorination operation is now fed to a reduction operation where the uranium tetrachloride will be reduced to form metallic uranium. The uranium tetrachloride is reduced by contacting it with a metal which is a greater reducing agent than uranium in the electromotive-force series, whereby the uranium tetrachloride will reduce to metallic uranium while the other metal is oxidized to form the corresponding chloride of the metal.

A preferred group of reducing metals which may be useful in carrying out this step of the process is either lithium or the alkaline earth metals such as calcium, magnesium, barium, and strontium. Preferably the reducing metal is either calcium or magnesium metal, and most preferably the reducing metal is magnesium. The following equation., using magnesium as the reducing metal by way of illustration, and not of limitation, shows the uranium reduction reaction.


No mention of Al though.

I can't find the heat of formation of UCl<sub>4</sub>. For AlCl<sub>3</sub> it's - 706 kJ/mol, for MgCl<sub>2</sub> it's - 641 kJ/mol, for KCl - 437 kJ/mol.

If the heat of formation of UCl<sub>4</sub> is ΔH, then the reaction enthalpies are:

UCl<sub>4</sub> + 4/3 Al ---> U + 4/3 AlCl<sub>3</sub>

Reaction enthalpy = - ΔH - 941 (kJ/mol)


UCl<sub>4</sub> + 2 Mg ---> U + 2 MgCl<sub>2</sub>

Reaction enthalpy = - ΔH - 1282 (kJ/mol)


UCl<sub>4</sub>. + 4 K ---> U + 4 KCl

Reaction enthalpy = - ΔH - 1748 (kJ/mol)

So, Al would be the least exothermic.

It doesn't sound implausible that the heat of formation of UCl4 would be somewhere in the 900 kJ/mol range, in which case higher temperature and removal of the AlCl3 would be needed to make the Al reduction proceed. For TiCl4, the heat of formation is - 815 kJ/mol.

[Edited on 31-5-2008 by blogfast25]

[Edited on 31-5-2008 by blogfast25]

StevenRS - 14-6-2008 at 12:42

I attempted the reduction of UCl4 with magnesium, and it was a success and a failure at the same time. It reacted quite quickly, but when the reduction was over, the uranium caught fire. (I think.) I expected this a little. All I need to do is stop the uranium from caching fire, and then heat it to drive of most of the magnesium chloride. I have a propane furnace, so heat is not a problem for me at all.

blogfast25 - 17-6-2008 at 09:32

Quote:
Originally posted by StevenRS
It reacted quite quickly, but when the reduction was over, the uranium caught fire. (I think.) I expected this a little.


If someone here has the heat of formation (@ 298 K) of UCl<sub>4</sub>, 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 MgCl<sub>2</sub>;) 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.

len1 - 18-6-2008 at 02:44

Quote:
Originally posted by MagicJigPipe
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)?


This is the one exception were I think its OK for one to just talk science without the slightest intention or expertise in actually doing it.

Isotopes can not be separated chemically, so the only method must be based on the mass difference. Gas diffusion separation uses the 0.8% mass differece [U(238)F6]/[U(235)F6]. Diffusion of a gas is related to the speed of its particles. I think many people may remember the first year high school experiment of mixing HCl with NH3, the latter move faster on average and so the NH4Cl fog forms predominantly in the HCl container. Thus in a 'single-stage' diffusion 0.4% more U235 will diffuse than its proportion in the original mixture. Repeat this n times using (1.004)^n and you can get the U235 fraction close to 1. Imperfections in the centrifuges will mean that the apparatus loses separation power as you get close to that 1. For a chain reaction you need enrichment to only about 5%.

There is no reason why any, even small, nation cant do this. To whit Israel with 3mln people. Most just dont want to.

As far as ore goes, you need 17kg U235 for one atom bomb thats a 10cm cube, or about 2.3ton natural uranium - thats a 45cm cube, which equates to 2.7ton of uranium ore concentrate U3O8. The average U concentration is 0.3% in Australian ore - so one atomic bomb = 690ton uranium ore - I need to order about 35 truck loads. Actually the centrifuge process is only about 70% efficient, while U smelting extracts max 85% U from Australian ore - so make that 1200ton, or 60 truck loads - not that much really - if only I had settled in Olympic Dam before 1975.

It is obvious from the above that an atom bomb is a huge waste of uranium - 99.5% is wasted (explains why depleted U costs 5$/kg). If you enrich your uranium to only 5%, and run it in a beeder, youll extract the energy from the U235, while the neutrons from the latter - 2.4 on average - will convert 2.4 its amount to Pu - another useful element. This way only 98% of natural U is wasted - but still too much. I think an H-bomb (actually its a lithium deuteride an isotope of LiH bomb) is a much 'greener' solution. It wastes only 45% natural uranium - U238 also fissions (I presume mostly thru Pu), it just doesnt release enough neutrons for a chain reaction, but in the neutron rich environment of an H-bomb, it reacts yielding about 50% of total energy.

[Edited on 18-6-2008 by len1]

blogfast25 - 18-6-2008 at 10:04

The following patent seems to suggest that the heat of the reduction reaction of UCl<sub>4</sub> with Mg is not enough to obtain the Uranium metal in liquid form and goes on to discuss the possibility of obtaining a lower meting alloy with Iron:

Process for continuous production of metallic uranium and uranium alloys:

Quote:
The uranium tetrachloride powder formed and collected in the chlorination
operation is now fed to a reduction operation where the uranium
tetrachloride will be reduced to form metallic uranium. The uranium
tetrachloride is reduced by contacting it with a metal which is a greater
reducing agent than uranium in the electromotive-force series, whereby the
uranium tetrachloride will reduce to metallic uranium while the other
metal is oxidized to form the corresponding chloride of the metal.


A preferred group of reducing metals which may be useful in carrying out
this step of the process is either lithium or the alkaline earth metals
such as calcium, magnesium, barium, and strontium. Preferably the reducing
metal is either calcium or magnesium metal, and most preferably the
reducing metal is magnesium. The following equation., using magnesium as
the reducing metal by way of illustration, and not of limitation, shows
the uranium reduction reaction.


UCl4 +2Mg ---> U+2MgCl2 (4)


While the above reaction is capable of forming metallic uranium, without
further additives, the resulting metallic uranium has a melting point of
about 1132° C. This necessitates carrying out the reaction at this
temperature or higher in order to maintain the uranium in liquid form in
the reactor to facilitate its removal when the reaction is carried out on
a continuous basis.


It would, therefore, be preferable to add to the reaction another metal
which is capable of alloying with the uranium to form an alloy or alloys
with lower melting temperatures than pure uranium. Typically, such metals
are those which form eutectic systems with uranium. These metals should
not interfere with the reduction reaction being carried out. A preferred
metal additive for this purpose is iron which, for example, will alloy
with uranium at a mole ratio of about 33 mole % iron, 67 mole % uranium
to form a low melting eutectic alloy having a melting point of about
725° C.


Other metals which could be used instead of iron, i.e., metals which can
form an alloy with a melting point lower than that of pure uranium
without, however, interfering with the uranium reduction reaction,
include: (a) one or more metals which form eutectic alloy systems with
uranium in which uranium is the major alloying constituent (i.e., in the
order of 60 mole % or higher), such as chromium, manganese, cobalt,
nickel, and the platinum metals ruthenium, rhodium, palladium, osmium,
iridium, and platinum; and (b) one or more metals which form eutectic
alloy systems with uranium in which uranium is the minor alloying
constituent (i.e., in the order of at least about 1 mole %, but less than
about 15 mole %) such as aluminum, gold, silver, copper, germanium, and
zinc.


And this here purchasable paper probably contains the value of UCl<sub>4</sub>'s heat of formation...

Nicodem - 18-6-2008 at 10:30

Quote:
Originally posted by blogfast25
And this here purchasable paper probably contains the value of UCl<sub>4</sub>'s heat of formation...

Though articles are otherwise requested in References, here it is for the those interested:

The standard molar enthalpy of formation of UNCl
M. Akabori, F. Kobayashi, H. Hayashi, T. Ogawa, M. E. Huntelaar, A. S. Booij and P. van Vlaanderen
The Journal of Chemical Thermodynamics, 34 (2002) 1461-1466.

Attachment: The standard molar enthalpy of formation of UNCl.pdf (90kB)
This file has been downloaded 1565 times


blogfast25 - 18-6-2008 at 12:05

Great!

From this paper the standard molar enthalpy of reaction at 298 K for U (s) + 2 Cl<sub>2</sub> (g) ---> UCl<sub>4</sub> (s) to be ΔH = - 1,018 kJ (per mol of UCl<sub>4</sub>;) (see page 5, Table 3) can be gleaned.

For UCl<sub>4</sub> (s) + 2 Mg (s) ---> U (s) + 2 MgCl<sub>2</sub> (s) we then obtain ΔH = - 264 kJ (per mol of UCl<sub>4</sub>;). That's only barely exothermic and almost certainly insufficient to melt both the U and the MgCl<sub>2</sub>. To be confirmed/infirmed tomorrow...

not_important - 18-6-2008 at 17:36

While checking on the Fray-Farthing-Chen-Cambridge process for another thread, I noticed that UO2 can be reduced by FFC means. This results in lightly sintered powdered uranium in frozen CaCl2 electrolyte, so the workup would need to be done in a inert atmosphere followed by melting and casting.


There's a reasonable description of the standard Mg-UF4 process here
http://www.barc.gov.in/webpages/letter/2000/200009-04.pdf

and some early 20th century experience here

http://books.google.com/books?id=lC4OAAAAYAAJ&pg=RA3-PA1...

which includes the comment
Quote:
An interesting feature developed in these experiments was the ease
with which metallic uranium oxidizes. In working with a charge which
produced uranium, or metal and carbide, if the button of sponge metal
was removed from the furnace while still red hot, and exposed to the air,
it immediately oxidized to black uranium oxide, a 25-lb. button being
converted completely to oxide in less than 5 minutes.

blogfast25 - 19-6-2008 at 09:03

Very interesting description of the UF<sub>4</sub>/Mg reduction process indeed! But fluorination of U is probably outside the capability envelope of most backyard scientists and the UCl<sub>4</sub>-route looks more promising from that perspective, so it's a case of finding the most suitable reductant. Ideally (for backyard science), the RT UCl<sub>4</sub>/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 UCl<sub>4</sub> + 4/3 Al ---> U + 4/3 AlCl<sub>3</sub> 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 MgCl<sub>2</sub> from 298 K to 1,405 K (including heat of fusion at MP = 987 K) is approximately 227 kJ.and for UCl<sub>4</sub> + 2 Mg ---> U + 2 MgCl<sub>2</sub> 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 UCl<sub>4</sub> + 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 UCl<sub>4</sub>. :(

Li is of course more difficult to obtain/produce in fine form than Mg and a mixture (layered perhaps?) of UCl<sub>4</sub> 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]

Fleaker - 28-12-2008 at 13:24

Just remembered this thread and the use for lithium.

I wonder if lithium for deOxing copper alloys would be sufficient. I suppose one could easily cut one of the copper capsules in half with a band saw and melt out the Li under argon.


This company sells it.
http://www.milward.com/index.htm

a description of what it's used for in the casting industry:
http://backyardmetalcasting.com/forums/viewtopic.php?t=927&a...

[Edited on 28-12-2008 by Fleaker]