Plutonium hexafluoride

Last updated
Plutonium hexafluoride [1]
Plutonium hexafluoride.svg
Plutonium-hexafluoride-3D-vdW.png
Neptunium(VI)-fluoride-3D-balls.png
Names
IUPAC name
plutonium(VI) fluoride
Identifiers
3D model (JSmol)
ChemSpider
PubChem CID
  • InChI=1S/6FH.Pu/h6*1H;/q;;;;;;+6/p-6 X mark.svgN
    Key: OJSBUHMRXCPOJV-UHFFFAOYSA-H Yes check.svgY
  • F[Pu](F)(F)(F)(F)F
Properties
PuF
6
AppearanceDark 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
Pnma, No. 62
octahedral (Oh)
0 D
Related compounds
Related fluoroplutoniums
 Plutonium trifluoride

Plutonium tetrafluoride

Hazards
GHS labelling:
GHS-pictogram-rondflam.svg GHS-pictogram-acid.svg GHS-pictogram-skull.svg GHS-pictogram-pollu.svg
Danger
NFPA 704 (fire diamond)
NFPA 704.svgHealth 4: Very short exposure could cause death or major residual injury. E.g. VX gasFlammability 0: Will not burn. E.g. waterInstability 4: Readily capable of detonation or explosive decomposition at normal temperatures and pressures. E.g. nitroglycerinSpecial hazard RA: Radioactive. E.g. plutonium
4
0
4
Radiation warning symbol 3.svg
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
X mark.svgN  verify  (what is  Yes check.svgYX mark.svgN ?)

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 isotope of plutonium is needed to avoid premature ignition of low-mass nuclear weapon designs by neutrons produced by spontaneous fission of plutonium-240.

Contents

Preparation

Plutonium hexafluoride 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, plutonium(IV) oxide, or plutonium(IV) oxalate at approximately 700 °C: [4] [6]

2 PuF
3 + 3 F
2 → 2 PuF
6
PuO
2 + 3 F
2 → PuF
6 + O
2
Pu(C2O4)2 + 3 F
2 → PuF
6 + 4 CO
2

Alternatively, plutonium(IV) fluoride oxidizes in an 800-°C oxygen atmosphere to plutonium hexafluoride and plutonium(IV) oxide: [7]

3 PuF
4 + O
2 → 2 PuF
6 + PuO
2

In 1984, the synthesis of plutonium hexafluoride at near–room-temperatures was achieved through the use of dioxygen difluoride. [8] [9] Hydrogen fluoride is not sufficient [10] : 42  even though it is a powerful fluorinating agent. Room temperature syntheses are also possible by using krypton difluoride [11] or irradiation with UV light. [12]

Properties

Physical properties

Phase diagram Plutonium hexafluoride phase diagram.svg
Phase diagram

Plutonium hexafluoride is a red-brown volatile solid, [1] [4] crystallizing in the orthorhombic crystal system with space group Pnma and lattice parameters a = 995 pm,b = 902 pm, and c = 526 pm. [13] It sublimes around 60 °C with heat 12.1 kcal/mol to a gas of octahedral molecules [2] with plutonium-fluorine bond lengths of 197.1 pm. [14] At high pressure, the gas condenses, with a triple point at 51.58 °C and 710 hPa (530 Torr); the heat of vaporization is 7.4 kcal/mol. [13] At temperatures below -180 °C, plutonium hexafluoride is colorless. [4]

Plutonium hexafluoride is paramagnetic, with molar magnetic susceptibility 0.173 mm3/mol. [15]

Spectroscopic properties

Plutonium hexafluoride admits six different oscillation modes: stretching modes v1, v2, and v3 and rotational modes v4, v5, and v6. [16] [17] The PuF
6 Raman spectrum cannot be observed, because irradiation at 564.1 nm induces photochemical decomposition. [18] Irradation at 532 nm induces fluorescence at 1900 nm and 4800 nm; irradiation at 1064 nm induces fluorescence about 2300 nm. [19] [20]

Absorption modes for PuF
6 [21]
Oscillationν1ν2ν3ν4ν5ν6
SymbolA1gEgF1uF1uF2gF2u
Wavelength (cm−1)628523615203211171
IR active?−−++−−
Raman active?++−−+−

Chemical properties

Plutonium hexafluoride is relatively hard to handle, being very corrosive, poisonous, and prone to auto-radiolysis. [22] [23] [24]

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] [25]

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. [26]

An important reaction involving PuF6 is the reduction to plutonium dioxide. Carbon monoxide generated from an oxygen-methane flame can perform the reduction. [27]

Decomposition reactions

Plutonium hexafluoride typically decomposes to plutonium tetrafluoride and fluorine gas. Thermal decomposition does not occur at room temperature, [28] [29] but proceeds very quickly at 280 °C. [5] [26] In the absence of any external cause for decomposition, the alpha-particle current from plutonium decay will generate auto-radiolysis, at a rate of 1.5%/day (half-time 1.5 months) in solid phase. [5] [23] [30] Storage in gas phase at pressures 50–100 torr (70–130 mbar) appears to minimize auto-radiolysis, and long-term recombination with freed fluorine does occur. [31] [ unreliable source? ]

Likewise, the compound is photosensitive, decomposing (possibly to plutonium pentafluoride and fluorine) under laser irradiation at a wavelength of less than 520 nm. [32]

Exposure to laser radiation at 564.1 nm or gamma rays will also induce rapid dissolution. [18] [24]

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:

The separation between plutonium and the americium contained proceeds through reaction with dioxygen difluoride. Aged PuF4 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. [33]

Separation of the hexafluorides of uranium and plutonium is also important in the reprocessing of nuclear waste. [34] [35] [36] 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. [10]

History

Shortly after plutonium's discovery and isolation in 1940, chemists began to postulate the existence of plutonium hexafluoride. Early experiments, which sought to mimic methods for the construction of uranium hexafluoride, had conflicting results; and definitive proof only appeared in 1942. [37] The Second World War then interrupted the publication of further research. [22]

Initial experiments, undertaken with extremely small quantities of plutonium, showed that a volatile plutonium compound would develop in a stream of fluorine gas only at temperatures exceeding 700 °C. Subsequent experiments showed that plutonium on a copper plate volatilized in a 500-°C fluorine stream, and that the reaction rate decreased with atomic number in the series uranium > neptunium > plutonium. [38] Brown and Hill, using milligram-scale samples of plutonium, completed in 1942 a distillation experiment with uranium hexafluoride, suggesting that higher fluorides of plutonium ought be unstable, and decompose to plutonium tetrafluoride at room temperature. Nevertheless, the vapor pressure of the compound appeared to correspond to that of uranium hexafluoride. [39] Davidson, Katz, and Orlemann showed in 1943 that plutonium in a nickel vessel volatilized under a fluorine atmosphere, and that the reaction product precipitated on a platinum surface. [40]

Fisher, Vaslow, and Tevebaugh conjectured that the higher fluorides exhibited a positive enthalpy of formation, that their formation would be endothermic, and consequently only stabilized at high temperatures. [41]

In 1944, Alan E. Florin  [ de ] prepared a volatile compound of plutonium believed to be the elusive plutonium hexafluoride, but the product decomposed prior to identification. The fluid substance would collect onto cooled glass and liquify, but then the fluoride atoms would react with the glass. [42]

By comparison between uranium and plutonium compounds, Brewer, Bromley, Gilles, and Lofgren computed the thermodynamic characteristics of plutonium hexafluoride. [43]

In 1950, Florin's efforts finally yielded the synthesis, [3] [44] and improved thermodynamic data and a new apparatus for its production soon followed. [2] Around the same time, British workers also developed a method for the production of PuF6. [4] [7]

Related Research Articles

<span class="mw-page-title-main">Neptunium</span> Chemical element with atomic number 93 (Np)

Neptunium is a chemical element; it has symbol Np and atomic number 93. A radioactive actinide metal, neptunium is the first transuranic element. It is named after Neptune, the planet beyond Uranus in the Solar System, which uranium is named after. A neptunium atom has 93 protons and 93 electrons, of which seven are valence electrons. Neptunium metal is silvery and tarnishes when exposed to air. The element occurs in three allotropic forms and it normally exhibits five oxidation states, ranging from +3 to +7. Like all actinides, it is radioactive, poisonous, pyrophoric, and capable of accumulating in bones, which makes the handling of neptunium dangerous.

<span class="mw-page-title-main">Nuclear reprocessing</span> Chemical operations that separate fissile material from spent fuel to be recycled as new fuel

Nuclear reprocessing is the chemical separation of fission products and actinides from spent nuclear fuel. Originally, reprocessing was used solely to extract plutonium for producing nuclear weapons. With commercialization of nuclear power, the reprocessed plutonium was recycled back into MOX nuclear fuel for thermal reactors. The reprocessed uranium, also known as the spent fuel material, can in principle also be re-used as fuel, but that is only economical when uranium supply is low and prices are high. Nuclear reprocessing may extend beyond fuel and include the reprocessing of other nuclear reactor material, such as Zircaloy cladding.

<span class="mw-page-title-main">Uranium hexafluoride</span> Chemical compound

Uranium hexafluoride, sometimes called hex, is an inorganic compound with the formula UF6. Uranium hexafluoride is a volatile, toxic white solid that is used in the process of enriching uranium, which produces fuel for nuclear reactors and nuclear weapons.

<span class="mw-page-title-main">Tungsten hexafluoride</span> Chemical compound

Tungsten(VI) fluoride, also known as tungsten hexafluoride, is an inorganic compound with the formula WF6. It is a toxic, corrosive, colorless gas, with a density of about 13 kg/m3 (22 lb/cu yd). It is the only known gaseous transition metal compound and the densest known gas under standard ambient temperature and pressure. WF6 is commonly used by the semiconductor industry to form tungsten films, through the process of chemical vapor deposition. This layer is used in a low-resistivity metallic "interconnect". It is one of seventeen known binary hexafluorides.

Dioxygen difluoride is a compound of fluorine and oxygen with the molecular formula O2F2. It can exist as an orange-colored solid which melts into a red liquid at −163 °C (110 K). It is an extremely strong oxidant and decomposes into oxygen and fluorine even at −160 °C (113 K) at a rate of 4% per day — its lifetime at room temperature is thus extremely short. Dioxygen difluoride reacts vigorously with nearly every chemical it encounters (including ordinary ice) leading to its onomatopoeic nickname FOOF (a play on its chemical structure and its explosive tendencies).

<span class="mw-page-title-main">Xenon tetrafluoride</span> Chemical compound

Xenon tetrafluoride is a chemical compound with chemical formula XeF
4. It was the first discovered binary compound of a noble gas. It is produced by the chemical reaction of xenon with fluorine:

<span class="mw-page-title-main">Xenon hexafluoride</span> Chemical compound

Xenon hexafluoride is a noble gas compound with the formula XeF6. It is one of the three binary fluorides of xenon that have been studied experimentally, the other two being XeF2 and XeF4. All known are exergonic and stable at normal temperatures. XeF6 is the strongest fluorinating agent of the series. It is a colorless solid that readily sublimes into intensely yellow vapors.

Fluoride volatility is the tendency of highly fluorinated molecules to vaporize at comparatively low temperatures. Heptafluorides, hexafluorides and pentafluorides have much lower boiling points than the lower-valence fluorides. Most difluorides and trifluorides have high boiling points, while most tetrafluorides and monofluorides fall in between. The term "fluoride volatility" is jargon used particularly in the context of separation of radionuclides.

<span class="mw-page-title-main">Plutonium tetrafluoride</span> Chemical compound

Plutonium(IV) fluoride is a chemical compound with the formula PuF4. This salt is generally a brown solid but can appear a variety of colors depending on the grain size, purity, moisture content, lighting, and presence of contaminants. Its primary use in the United States has been as an intermediary product in the production of plutonium metal for nuclear weapons usage.

<span class="mw-page-title-main">Technetium hexafluoride</span> Chemical compound

Technetium hexafluoride or technetium(VI) fluoride (TcF6) is a yellow inorganic compound with a low melting point. It was first identified in 1961. In this compound, technetium has an oxidation state of +6, the highest oxidation state found in the technetium halides. In this respect, technetium differs from rhenium, which forms a heptafluoride, ReF7. Technetium hexafluoride occurs as an impurity in uranium hexafluoride, as technetium is a fission product of uranium (spontaneous fission in natural uranium, possible contamination from induced fission inside the reactor in reprocessed uranium). The fact that the boiling point of the hexafluorides of uranium and technetium are very close to each other presents a problem in using fluoride volatility in nuclear reprocessing.

<span class="mw-page-title-main">Krypton difluoride</span> Chemical compound

Krypton difluoride, KrF2 is a chemical compound of krypton and fluorine. It was the first compound of krypton discovered. It is a volatile, colourless solid at room temperature. The structure of the KrF2 molecule is linear, with Kr−F distances of 188.9 pm. It reacts with strong Lewis acids to form salts of the KrF+ and Kr
2F+
3 cations.

<span class="mw-page-title-main">Manganese(IV) fluoride</span> Chemical compound

Manganese tetrafluoride, MnF4, is the highest fluoride of manganese. It is a powerful oxidizing agent and is used as a means of purifying elemental fluorine.

<span class="mw-page-title-main">Plutonium(III) fluoride</span> Chemical compound

Plutonium(III) fluoride or plutonium trifluoride is the chemical compound composed of plutonium and fluorine with the formula PuF3. This salt forms violet crystals. Plutonium(III) fluoride has the LaF3 structure where the coordination around the plutonium atoms is complex and usually described as tri-capped trigonal prismatic.

A hexafluoride is a chemical compound with the general formula QXnF6, QXnF6m−, or QXnF6m+. Many molecules fit this formula. An important hexafluoride is hexafluorosilicic acid (H2SiF6), which is a byproduct of the mining of phosphate rock. In the nuclear industry, uranium hexafluoride (UF6) is an important intermediate in the purification of this element.

<span class="mw-page-title-main">Molybdenum hexafluoride</span> Chemical compound

Molybdenum hexafluoride, also molybdenum(VI) fluoride, is the inorganic compound with the formula MoF6. It is the highest fluoride of molybdenum. It is a colourless solid and melts just below room temperature and boils in 34 °C. It is one of the seventeen known binary hexafluorides.

<span class="mw-page-title-main">Rhenium hexafluoride</span> Chemical compound

Rhenium hexafluoride, also rhenium(VI) fluoride, (ReF6) is a compound of rhenium and fluorine and one of the seventeen known binary hexafluorides.

Chromium pentafluoride is the inorganic compound with the chemical formula CrF5. It is a red volatile solid that melts at 34 °C. It is the highest known chromium fluoride, since the hypothetical chromium hexafluoride has not yet been synthesized.

<span class="mw-page-title-main">Neptunium(VI) fluoride</span> Chemical compound

Neptunium(VI) fluoride (NpF6) is the highest fluoride of neptunium, it is also one of seventeen known binary hexafluorides. It is a volatile orange crystalline solid. It is relatively hard to handle, being very corrosive, volatile and radioactive. Neptunium hexafluoride is stable in dry air but reacts vigorously with water.

Fluorine forms a great variety of chemical compounds, within which it always adopts an oxidation state of −1. With other atoms, fluorine forms either polar covalent bonds or ionic bonds. Most frequently, covalent bonds involving fluorine atoms are single bonds, although at least two examples of a higher order bond exist. Fluoride may act as a bridging ligand between two metals in some complex molecules. Molecules containing fluorine may also exhibit hydrogen bonding. Fluorine's chemistry includes inorganic compounds formed with hydrogen, metals, nonmetals, and even noble gases; as well as a diverse set of organic compounds. For many elements the highest known oxidation state can be achieved in a fluoride. For some elements this is achieved exclusively in a fluoride, for others exclusively in an oxide; and for still others the highest oxidation states of oxides and fluorides are always equal.

Neptunium compounds are compounds containing the element neptunium (Np). Neptunium has five ionic oxidation states ranging from +3 to +7 when forming chemical compounds, which can be simultaneously observed in solutions. It is the heaviest actinide that can lose all its valence electrons in a stable compound. The most stable state in solution is +5, but the valence +4 is preferred in solid neptunium compounds. Neptunium metal is very reactive. Ions of neptunium are prone to hydrolysis and formation of coordination compounds.

References

  1. 1 2 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. 1 2 3 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. Originally published as
  3. 1 2 3 Florin, Alan E. (9 November 1950). Plutonium Hexafluoride: Second Report on the Preparation and Properties (PDF) (Technical report). Los Alamos Scientific Laboratory. LAMS-1168.
  4. 1 2 3 4 5 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. Originally published as
    • Mandleberg, C. J.; Rae, H. K.; Hurst, R.; Long, G.; Davis, D.; Francis, K. E. (April 1953). Plutonium Hexafluoride: Preparation and Some Physical Properties (Technical report). Vol. I. Atomic Energy Research Establishment. C/R-1172.
    • Hurst, R.; Mandleberg, C. J.; Rae, H. K.; Davis, D.; Francis, K. E. (January 1953). Plutonium Hexafluoride: Preparation and Some Physical Properties (Technical report). Vol. II. Atomic Energy Research Establishment. C/R-1312.
  5. 1 2 3 4 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. ↑ Dawson, J. K.; Truswell, A. E. (22 February 1951). The Preparation of Plutonium Trifluoride and Tetrafluoride by the Use of Hydrogen Fluoride (Technical report). Atomic Energy Research Establishment. C/R-662.
  7. 1 2 Mandleberg, C. J.; et al. (1952). (Technical report). Atomic Energy Research Establishment. C/R-157.{{cite tech report}}: Missing or empty |title= (help)
  8. ↑ 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.
  9. ↑ Erilov, P. E.; Titov, V. V.; Serik, V. F.; Sokolov, V. B. (2002). "Low-Temperature Synthesis of Plutonium Hexafluoride". Atomic Energy. 92 (1): 57–63. doi:10.1023/A:1015106730457. S2CID   96612181.
  10. 1 2 Evaluation of the U.S. Department of Energy's Alternatives for the Removal and Disposition of Molten Salt Reactor Experiment Fluoride Salts. Washington, DC: National Academies Press. 1997. doi:10.17226/5538. ISBN   978-0-309-05684-7 – via NAP.edu.
  11. ↑ 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.
  12. ↑ 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.
  13. 1 2 Gmelins Handbuch der anorganischen Chemie[Gmelin's Handbook of Inorganic Chemistry]. 71 (Transurane [Transuranics]) (in German). Vol. C. pp. 108–114.
  14. ↑ Kimura, Masao; Schomaker, Verner; Smith, Darwin W.; Weinstock, Bernard (May 1968). "Electron-Diffraction Investigation of the Hexafluorides of Tungsten, Osmium, Iridium, Uranium, Neptunium, and Plutonium". The Journal of Chemical Physics. 48 (9): 4001–4012. Bibcode:1968JChPh..48.4001K. doi:10.1063/1.1669727. ISSN   0021-9606.
  15. ↑ Gruen, D. M.; Malm, J. G.; Weinstock, B. (April 1956). "Magnetic Susceptibility of Plutonium Hexafluoride". The Journal of Chemical Physics. 24 (4): 905–906. Bibcode:1956JChPh..24..905G. doi:10.1063/1.1742635. ISSN   0021-9606.
  16. ↑ Steindler, Martin J.; Gunther, William H. (August 1964). "The absorption spectrum of plutonium hexafluoride". Spectrochimica Acta. 20 (8): 1319–1322. Bibcode:1964AcSpe..20.1319S. doi:10.1016/0371-1951(64)80159-4.
  17. ↑ Walters, R.T.; Briesmeister, R.A. (January 1984). "Absorption spectrum of plutonium hexafluoride in the 3000–9000 Å spectral region". Spectrochimica Acta Part A: Molecular Spectroscopy. 40 (7): 587–589. Bibcode:1984AcSpA..40..587W. doi:10.1016/0584-8539(84)80108-7.
  18. 1 2 Hawkins, N. J.; Mattraw, H. C.; Sabol, W. W. (24 May 1954). Infrared Spectrum and Thermodynamic Properties of PuF6 (Technical report). Knolls Atomic Power Laboratory. KAPL-1007.
  19. ↑ Beitz, James V.; Williams, Clayton W.; Carnall, W. T. (March 1982). "Fluorescence studies of neptunium and plutonium hexafluoride vapors". The Journal of Chemical Physics. 76 (5): 2756–2757. Bibcode:1982JChPh..76.2756B. doi:10.1063/1.443223. ISSN   0021-9606.
  20. ↑ Beitz, James V.; Williams, Clayton W.; Carnall, W. T. (19 May 1983). "11. Plutonium Hexafluoride Gas Photophysics and Photochemistry". In Carnall, William T.; Choppin, Gregory R. (eds.). Plutonium Chemistry. ACS Symposium Series. Vol. 216. Washington, D.C.: American Chemical Society. pp. 155–172. doi:10.1021/bk-1983-0216.ch011. ISBN   978-0-8412-0772-1.
  21. ↑
  22. 1 2 Steindler, Martin J. (1 August 1963). Laboratory Investigations in Support of Fluid-bed Fluoride Volatility Processes (Technical report). Vol. II: The Properties of Plutonium Hexafluoride. Argonne National Laboratory. doi:10.2172/4170539. ANL-6753.
  23. 1 2 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.
  24. 1 2 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.
  25. ↑ 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.
  26. 1 2 Malm, John G.; Weinstock, Bernard; Weaver, E. Eugene (1958). "The Preparation and Properties of NpF5; a Comparison with PuF5". The Journal of Physical Chemistry. 62 (12): 1506–1508. doi:10.1021/j150570a009. ISSN   0022-3654.
  27. ↑ Pokidyshev, A. M.; Tsarenko, I. A.; Serik, V. F.; Sokolov, V. B. (October 2003). "Reduction of Plutonium Hexafluoride Using Gaseous Reagents". Atomic Energy. 95 (4): 701–708. doi:10.1023/B:ATEN.0000010988.94533.24. ISSN   1063-4258. S2CID   93145477.
  28. ↑ Trevorrow, L. E.; Shinn, W. A.; Steunenberg, R. K. (March 1961). "The Thermal Decomposition of Plutonium Hexafluoride". The Journal of Physical Chemistry. 65 (3): 398–403. doi:10.1021/j100821a003. ISSN   0022-3654.
  29. ↑ Fischer, J.; Trevorrow, L.; Shinn, W. (October 1961). "The Kinetics and Mechanism of the Thermal Decomposition of Plutonium Hexafluoride". The Journal of Physical Chemistry. 65 (10): 1843–1846. doi:10.1021/j100827a036. ISSN   0022-3654.
  30. ↑
    • Steindler 1963
    • Wagner, R. P.; Shinn, W. A.; Fischer, J.; Steindler, Martin J. (1 May 1963). Laboratory Investigations in Support of Fluid-bed Fluoride Volatility Processes (Technical report). Vol. VII: The Decomposition of Gaseous Plutonium Hexafluoride by Alpha Radiation. Argonne National Laboratory. doi:10.2172/4628896. ANL-7013.
  31. ↑ Morse, L. R. (2005), "PuF6 gas pressure in aged cylinders" (personal communication to D. L. Clark), Los Alamos, NM.
  32. ↑ US 4670239,Sherman W. Rabideau&George M. Campbell,"Photochemical Preparation of Plutonium Pentafluoride",published June 2, 1987, assigned to The United States of America, but see also Lobikov, E. A.; Prusakov, V. N.; Serik, V. F. (August–September 1992). "Plutonium Hexafluoride Decomposition under the Action of Laser Radiation". Journal of Fluorine Chemistry . 58 (2–3): 277. doi:10.1016/S0022-1139(00)80734-4, in which the decay product is identified as tetrafluoride instead.
  33. ↑ 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.
  34. ↑
    • US 3708568A,Gilliher, W.; Harris, R.& Ledoux, R.,"Removal of Plutonium from Plutonium Hexafluoride-Uranium Hexafluoride Mixtures",published 1973-01-02, assigned to Atomic Energy Commission
    • US 4172114A,Mitsuhiro Nishimura et al,"Method for purifying plutonium hexafluoride",published 1979-10-23, assigned to Japan Atomic Energy Research Institute
  35. ↑ 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. OSTI   6168468.
  36. ↑ 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.
  37. ↑ Seaborg, G. T. (1942). (Technical report). University of Chicago Metallurgical Laboratory. CN-125.{{cite tech report}}: Missing or empty |title= (help)
  38. ↑ Brown, H. S.; Hill, O. F.; Jaffay, A. H. (1942). (Technical report). University of Chicago Metallurgical Laboratory. CN-343.{{cite tech report}}: Missing or empty |title= (help)
  39. ↑ Brown, H. S.; Hill, O. F. (12 November 1942). (Technical report). University of Chicago Metallurgical Laboratory. CN-363.{{cite tech report}}: Missing or empty |title= (help)
  40. ↑ Davidson, N. R.; Katz, J. J.; Orlemann, O. F. (11 October 1943). (Technical report). University of Chicago Metallurgical Laboratory. CN-987.{{cite tech report}}: Missing or empty |title= (help)
  41. ↑ Fisher, R. W.; Vaslow, F.; Tevebaugh, A. D. (10 August 1944). (Technical report). Iowa State College. CN-1783.{{cite tech report}}: Missing or empty |title= (help)
  42. ↑ Florin, Alan E. (1 October 1944). (Technical report). University of Chicago Metallurgical Laboratory. CN-2159.{{cite tech report}}: Missing or empty |title= (help)
  43. ↑
  44. ↑ Florin, Alan E. (16 October 1950). Plutonium Hexafluoride, Plutonium (VI) Oxyfluoride: Preparation, Identification, and Some Properties (PDF) (Technical report). Los Alamos Scientific Laboratory. LAMS-1118.

This page is based on this Wikipedia article
Text is available under the CC BY-SA 4.0 license; additional terms may apply.
Images, videos and audio are available under their respective licenses.