Plutonium hexafluoride

From Wikipedia, the free encyclopedia
[1]
Stereo structural formula of plutonium hexafluoride
Plutonium-hexafluoride-3D-vdW.png
Neptunium(VI)-fluoride-3D-balls.png
Names
IUPAC name
plutonium(VI) fluoride
Identifiers
3D model (JSmol)
ChemSpider
  • InChI=1S/6FH.Pu/h6*1H;/q;;;;;;+6/p-6 ☒N
    Key: OJSBUHMRXCPOJV-UHFFFAOYSA-H checkY
  • F[Pu](F)(F)(F)(F)F
Properties
PuF
6
Appearance Dark red, opaque crystals
Density 5.08 g·cm−3
Melting point 52 °C (126 °F; 325 K)
Boiling point 62 °C (144 °F; 335 K)
Structure
Orthorhombic, oP28
Space group
 Pnma, No. 62
octahedral (Oh)
Dipole moment
 0 D
Related compounds
Related fluoroplutoniums
 Plutonium trifluoride

Plutonium tetrafluoride

Hazards
GHS labelling:
GHS03: OxidizingGHS05: CorrosiveGHS06: ToxicGHS09: Environmental hazard
Signal word
 Danger
NFPA 704 (fire diamond)
4
0
4
Special hazard RA: Radioactive. E.g. plutonium
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
☒N  (what is checkY☒N ?)
Infobox references

Plutonium hexafluoride is the highest fluoride of plutonium, and is of interest for laser enrichment of plutonium, in particular for the production of pure plutonium-239 from irradiated uranium. This pure plutonium is needed to avoid premature ignition of low-mass nuclear weapon designs by neutrons produced by spontaneous fission of plutonium-240.

Preparation[]

It is prepared by fluorination of plutonium tetrafluoride (PuF4) by powerful fluorinating agents such as elemental fluorine.[2][3][4][5]

PuF
4 + F
2 → PuF
6

This reaction is endothermic. The product forms relatively quickly at temperatures of 750 °C, and high yields may be obtained by quickly condensing the product and removing it from equilibrium.[5]

It can also be obtained by fluorination of plutonium(III) fluoride or plutonium(IV) oxide.[4]

2 PuF
3 + 3 F
2 → 2 PuF
6
PuO
2 + 3 F
2 → PuF
6 + O
2

In 1984, the synthesis of plutonium hexafluoride was achieved at unprecedented low temperatures through the use of dioxygen difluoride. Previous techniques needed temperatures so high that the plutonium hexafluoride produced would decompose rapidly.[6] Hydrogen fluoride is not sufficient;[7] even though it is a powerful fluorinating agent. Room temperature syntheses are also possible by using krypton difluoride[8] or irradiation with UV light.[9]

Properties[]

Physical properties[]

Plutonium hexafluoride is a red-brown volatile crystalline solid;[1] the heat of sublimation is 12.1 kcal/mol[2] and the heat of vaporization 7.4 kcal/mol. It crystallizes in the orthorhombic crystal system. As a gas, the molecule has octahedral symmetry (point group Oh)

Chemical properties[]

Plutonium hexafluoride is relatively hard to handle, being very corrosive and prone to auto-radiolysis.[10][11]

Reactions with other compounds[]

PuF6 is stable in dry air, but reacts vigorously with water, including atmospheric moisture, to form plutonium(VI) oxyfluoride and hydrofluoric acid.[3][12]

PuF
6 + 2 H
2O → PuO
2F
2 + 4 HF

It can be stored for a long time in a quartz or pyrex ampoule, provided there are no traces of moisture, the glass has been thoroughly outgassed, and any traces of hydrogen fluoride have been removed from the compound.[13]

A significant reaction of PuF6 is the reduction to plutonium dioxide. Carbon monoxide generated from an oxygen-methane flame is an example of a good reducing agent for producing actinide oxides directly from the hexafluorides.

Decomposition reactions[]

Plutonium hexafluoride decomposes to plutonium tetrafluoride and fluorine gas.

  • It can undergo thermal decomposition, which does not occur at room temperature but proceeds very quickly at 280 °C.[5]
  • Another possibility is auto-radiolysis, that is decomposing due to its own radioactivity. Emitted alpha particles moving through the crystal lattice cause bonds to be broken, leading to decomposition to lower fluorides and fluorine gas. The decomposition rate through alpha radiation is 1.5% per day on average in the solid phase, but is significantly smaller in the gas phase.[5] It also decomposes from gamma radiation.[14]
  • Under laser irradiation at a wavelength of less than 520 nm, it decomposes to and fluorine;[15] after more irradiation it decomposes further to plutonium tetrafluoride.[16]

Uses[]

Plutonium hexafluoride plays a role in the enrichment of plutonium, in particular for the isolation of the fissile isotope 239Pu from irradiated uranium. For use in nuclear weaponry, the 241Pu present must be removed for two reasons:

  • It generates enough neutrons by spontaneous fission to cause an uncontrollable reaction.
  • It undergoes beta decay to form 241Am, leading to the accumulation of americium over long periods of storage which must be removed.

The separation of plutonium and the americium contained proceeds through a reaction with dioxygen difluoride. PuF4 that has been stored for a long time is fluorinated at room temperature to gaseous PuF6, which is separated and reduced back to PuF4, whereas any AmF4 present does not undergo the same conversion. The product thus contains very little amounts of americium, which becomes concentrated in the unreacted solid.[17]

The separation of the hexafluorides of uranium and plutonium is important in the reprocessing of nuclear waste.[18][19] From a molten salt mixture containing both elements, uranium can largely be removed by fluorination to UF6, which is stable at higher temperatures, with only small amounts of plutonium escaping as PuF6.[20]

References[]

  1. ^ a b Lide, David R. (2009). Handbook of Chemistry and Physics (90 ed.). Boca Raton, Florida: CRC Press. pp. 4–81. ISBN 978-1-4200-9084-0. (webelements.com)
  2. ^ a b Florin, Alan E.; Tannenbaum, Irving R.; Lemons, Joe F. (1956). "Preparation and properties of plutonium hexafluoride and identification of plutonium(VI) oxyfluoride". Journal of Inorganic and Nuclear Chemistry. 2 (5–6): 368–379. doi:10.1016/0022-1902(56)80091-2.
  3. ^ a b A. E. Florin (9 November 1950). "Plutonium Hexafluoride: Second Report On The Preparation and Properties (LA-1168)" (PDF). Los Alamos Scientific Laboratory.
  4. ^ a b Mandleberg, C.J.; Rae, H.K.; Hurst, R.; Long, G.; Davies, D.; Francis, K.E. (1956). "Plutonium hexafluoride". Journal of Inorganic and Nuclear Chemistry. 2 (5–6): 358–367. doi:10.1016/0022-1902(56)80090-0.
  5. ^ a b c d Weinstock, Bernard; Malm, John G. (July 1956). "The properties of plutonium hexafluoride". Journal of Inorganic and Nuclear Chemistry. 2 (5–6): 380–394. doi:10.1016/0022-1902(56)80092-4.
  6. ^ Malm, J. G.; Eller, P. G.; Asprey, L. B. (1984). "Low temperature synthesis of plutonium hexafluoride using dioxygen difluoride". Journal of the American Chemical Society. 106 (9): 2726–2727. doi:10.1021/ja00321a056.
  7. ^ Evaluation of the U.S. Department of Energy's Alternatives for the Removal and Disposition of Molten Salt Reactor Experiment Fluoride Salts. 1997. p. 42. doi:10.17226/5538. ISBN 978-0-309-05684-7.
  8. ^ Asprey, L. B.; Eller, P. G.; Kinkead, Scott A. (1986). "Formation of actinide hexafluorides at ambient temperatures with krypton difluoride". Inorganic Chemistry. 25 (5): 670–672. doi:10.1021/ic00225a016. ISSN 0020-1669.
  9. ^ Trevorrow, L.E.; Gerding, T.J.; Steindler, M.J. (1969). "Ultraviolet-activated synthesis of plutonium hexafluoride at room temperature". Inorganic and Nuclear Chemistry Letters. 5 (10): 837–839. doi:10.1016/0020-1650(69)80068-1.
  10. ^ Bibler, Ned E. (23 August 1979). "α and β Radiolysis of Plutonium Hexafluoride Vapor". J. Phys. Chem. 83 (17): 2179–2186. doi:10.1021/j100480a001.
  11. ^ Steindler, M.J.; Steidl, D.V.; Fischer, J. (November 1964). "The decomposition of plutonium hexafluoride by gamma radiation". Journal of Inorganic and Nuclear Chemistry. 26 (11): 1869–1878. doi:10.1016/0022-1902(64)80011-7.
  12. ^ Kessie, R. W. (1967). "Plutonium and Uranium Hexafluoride Hydrolysis Kinetics". Industrial & Engineering Chemistry Process Design and Development. 6 (1): 105–111. doi:10.1021/i260021a018. ISSN 0196-4305.
  13. ^ Malm, John G.; Weinstock, Bernard; Weaver, E. Eugene (1958). "The Preparation and Properties of NpF 5 ; a Comparison with PuF 5". The Journal of Physical Chemistry. 62 (12): 1506–1508. doi:10.1021/j150570a009. ISSN 0022-3654.
  14. ^ Steindler, M.J.; Steidl, D.V.; Fischer, J. (1964). "The decomposition of plutonium hexafluoride by gamma radiation". Journal of Inorganic and Nuclear Chemistry. 26 (11): 1869–1878. doi:10.1016/0022-1902(64)80011-7.
  15. ^ US 4670239, Sherman W. Rabideau & George M. Campbell, "Photochemical Preparation of Plutonium Pentafluoride", published June 2, 1987, assigned to The United States of America 
  16. ^ Lobikov, E. A.; Prusakov, V. N.; Serik, V. F. (August–September 1992). "Plutonium Hexafluoride Decomposition under the Action of Laser Radiation". . 58 (2–3): 277. doi:10.1016/S0022-1139(00)80734-4.
  17. ^ Mills, T.R.; Reese, L.W. (1994). "Separation of plutonium and americium by low-temperature fluorination". Journal of Alloys and Compounds. 213–214: 360–362. doi:10.1016/0925-8388(94)90931-8.
  18. ^ Moser, W.Scott; Navratil, James D. (1984). "Review of major plutonium pyrochemical technology". Journal of the Less Common Metals. 100: 171–187. doi:10.1016/0022-5088(84)90062-6.
  19. ^ Drobyshevskii, Yu. V.; Ezhov, V. K.; Lobikov, E. A.; Prusakov, V. N.; Serik, V. F.; Sokolov, V. B. (2002). "Application of Physical Methods for Reducing Plutonium Hexafluoride". Atomic Energy. 93 (1): 578–588. doi:10.1023/A:1020840716387. S2CID 100100314.
  20. ^ Read "Evaluation of the U.S. Department of Energy's Alternatives for the Removal and Disposition of Molten Salt Reactor Experiment Fluoride Salts" at NAP.edu. 1997. doi:10.17226/5538. ISBN 978-0-309-05684-7.

Retrieved from ""