Plutonium hexafluoride

Plutonium hexafluoride is prepared by fluorination of plutonium tetrafluoride (PuF₄) by powerful fluorinating agents such as elemental fluorine.{{cite journal |last1=Florin |first1=Alan E. |last2=Tannenbaum |first2=Irving R. |last3=Lemons |first3=Joe F. |year=1956 |title=Preparation and properties of plutonium hexafluoride and identification of plutonium(VI) oxyfluoride |journal=Journal of Inorganic and Nuclear Chemistry |volume=2 |issue=5–6 |pages=368–379 |doi=10.1016/0022-1902(56)80091-2}} Originally published as

• {{cite tech report|first=Alan E.|last=Florin|title=Thermodynamic Properties of Plutonium Hexafluoride: a Preliminary Report|institution=Los Alamos Scientific Laboratory|number=LAMS-1587|date=15 May 1953|url=http://www.fas.org/sgp/othergov/doe/lanl/lib-www/la-pubs/00318358.pdf}}
• {{cite tech report|first1=I. R.|last1=Tannenbaum|first2=Alan E.|last2=Florin|title=An Improved Apparatus for the Production of Plutonium Hexafluoride|institution=Los Alamos Scientific Laboratory|number=LA-1580|date=15 May 1953|url=http://www.fas.org/sgp/othergov/doe/lanl/lib-www/la-pubs/00371935.pdf}}

{{cite tech report|first=Alan E.|last=Florin|title=Plutonium Hexafluoride: Second Report on the Preparation and Properties|institution=Los Alamos Scientific Laboratory|number=LAMS-1168|date=9 November 1950|url=http://www.fas.org/sgp/othergov/doe/lanl/lib-www/la-pubs/00419717.pdf}}

{{cite journal|last1=Mandleberg|first1=C.J.|last2=Rae|first2=H.K.|last3=Hurst|first3=R.|last4=Long|first4=G.|last5=Davies|first5=D.|last6=Francis|first6=K.E.|year=1956|title=Plutonium hexafluoride|journal=Journal of Inorganic and Nuclear Chemistry|volume=2|issue=5–6|pages=358–367|doi=10.1016/0022-1902(56)80090-0}} Originally published as

• {{cite tech report|first1=C. J.|last1=Mandleberg|first2=H. K.|last2=Rae|first3=R.|last3=Hurst|first4=G.|last4=Long|first5=D.|last5=Davis|first6=K. E.|last6=Francis|title=Plutonium Hexafluoride: Preparation and Some Physical Properties|volume=I|institution=Atomic Energy Research Establishment|number=C/R-1172|date=April 1953}}
• {{cite tech report|first2=C. J.|last2=Mandleberg|first3=H. K.|last3=Rae|first1=R.|last1=Hurst|first4=D.|last4=Davis|first5=K. E.|last5=Francis|title=Plutonium Hexafluoride: Preparation and Some Physical Properties|volume=II|institution=Atomic Energy Research Establishment|number=C/R-1312|date=January 1953}}{{cite journal|last1=Weinstock|first1=Bernard|last2=Malm|first2=John G.|date=July 1956|title=The properties of plutonium hexafluoride|journal=Journal of Inorganic and Nuclear Chemistry|volume=2|issue=5–6|pages=380–394|doi=10.1016/0022-1902(56)80092-4}}

:{{chem|Pu|F|4}} + {{chem|F|2}} → {{chem|Pu|F|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.

It can also be obtained by fluorination of plutonium(III) fluoride, plutonium(IV) oxide, or plutonium(IV) oxalate at approximately 700 °C:{{Cite tech report|first1=J. K.|last1=Dawson|first2=A. E.|last2=Truswell|title=The Preparation of Plutonium Trifluoride and Tetrafluoride by the Use of Hydrogen Fluoride|institution=Atomic Energy Research Establishment|number=C/R-662|date=22 February 1951}}

:2 {{chem|Pu|F|3}} + 3 {{chem|F|2}} → 2 {{chem|Pu|F|6}}

:{{chem|Pu|O|2}} + 3 {{chem|F|2}} → {{chem|Pu|F|6}} + {{chem|O|2}}

:{{Chem2|Pu(C2O4)2}} + 3 {{Chem|F|2}} → {{Chem|Pu|F|6}} + 4 {{Chem|C|O|2}}

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

:3 {{Chem|Pu|F|4}} + {{Chem|O|2}} → 2 {{Chem|Pu|F|6}} + {{Chem|Pu|O|2}}

In 1984, the synthesis of plutonium hexafluoride at near–room-temperatures was achieved through the use of dioxygen difluoride.{{cite journal|last1=Malm|first1=J. G.|last2=Eller|first2=P. G.|last3=Asprey|first3=L. B.|year=1984|title=Low temperature synthesis of plutonium hexafluoride using dioxygen difluoride|journal=Journal of the American Chemical Society|volume=106|issue=9|pages=2726–2727|doi=10.1021/ja00321a056}}{{Cite journal |last1=Erilov |first1=P. E. |last2=Titov |first2=V. V. |last3=Serik |first3=V. F. |last4=Sokolov |first4=V. B. |date=2002 |title=Low-Temperature Synthesis of Plutonium Hexafluoride |journal=Atomic Energy |volume=92 |issue=1 |pages=57–63 |doi=10.1023/A:1015106730457|s2cid=96612181 }} Hydrogen fluoride is not sufficient{{rp|42}} even though it is a powerful fluorinating agent. Room temperature syntheses are also possible by using krypton difluoride{{Cite journal|last1=Asprey|first1=L. B.|last2=Eller|first2=P. G.|last3=Kinkead|first3=Scott A.|date=1986|title=Formation of actinide hexafluorides at ambient temperatures with krypton difluoride|url=https://pubs.acs.org/doi/abs/10.1021/ic00225a016|journal=Inorganic Chemistry|language=en|volume=25|issue=5|pages=670–672|doi=10.1021/ic00225a016|issn=0020-1669}} or irradiation with UV light.{{Cite journal|last1=Trevorrow|first1=L.E.|last2=Gerding|first2=T.J.|last3=Steindler|first3=M.J.|date=1969|title=Ultraviolet-activated synthesis of plutonium hexafluoride at room temperature|url=https://linkinghub.elsevier.com/retrieve/pii/0020165069800681|journal=Inorganic and Nuclear Chemistry Letters|language=en|volume=5|issue=10|pages=837–839|doi=10.1016/0020-1650(69)80068-1}}

File:Plutonium hexafluoride phase diagram.svg

Plutonium hexafluoride is a red-brown volatile solid,{{cite book |last=Lide |first=David R. |url=https://archive.org/details/handbookchemistr00lide |title=Handbook of Chemistry and Physics |publisher=CRC Press |year=2009 |isbn=978-1-4200-9084-0 |edition=90 |location=Boca Raton, Florida |pages=[https://archive.org/details/handbookchemistr00lide/page/n1301 4]–81 |url-access=limited}} ([https://www.webelements.com/compounds/plutonium/plutonium_hexafluoride.html webelements.com]) crystallizing in the orthorhombic crystal system with space group Pnma and lattice parameters {{Math|1=a = 995 pm}}, {{Math|1=b = 902 pm}}, and {{Math|1=c = 526 pm}}.{{Cite book |title=Gmelins Handbuch der anorganischen Chemie |series=71 ({{lang|de|Transurane}} [Transuranics]) |volume=C |pages=108–114 |language=de |trans-title=Gmelin's Handbook of Inorganic Chemistry}} It sublimes around 60 °C with heat 12.1 kcal/mol to a gas of octahedral molecules with plutonium-fluorine bond lengths of 197.1 pm.{{Cite journal |last1=Kimura |first1=Masao |last2=Schomaker |first2=Verner |last3=Smith |first3=Darwin W. |last4=Weinstock |first4=Bernard |date=May 1968 |title=Electron-Diffraction Investigation of the Hexafluorides of Tungsten, Osmium, Iridium, Uranium, Neptunium, and Plutonium |url=https://pubs.aip.org/aip/jcp/article/48/9/4001-4012/770361 |journal=The Journal of Chemical Physics |language=en |volume=48 |issue=9 |pages=4001–4012 |doi=10.1063/1.1669727 |bibcode=1968JChPh..48.4001K |issn=0021-9606}} At high pressure, the gas condenses, with a triple point at 51.58 °C and {{Convert|710|hPa|Torr|abbr=on}}; the heat of vaporization is 7.4 kcal/mol. At temperatures below -180 °C, plutonium hexafluoride is colorless.

Plutonium hexafluoride is paramagnetic, with molar magnetic susceptibility 0.173 mm³/mol.{{Cite journal |last1=Gruen |first1=D. M. |last2=Malm |first2=J. G. |last3=Weinstock |first3=B. |date=April 1956 |title=Magnetic Susceptibility of Plutonium Hexafluoride |url=https://pubs.aip.org/aip/jcp/article/24/4/905-906/74261 |journal=The Journal of Chemical Physics |language=en |volume=24 |issue=4 |pages=905–906 |doi=10.1063/1.1742635 |bibcode=1956JChPh..24..905G |issn=0021-9606}}

Plutonium hexafluoride admits six different oscillation modes: stretching modes {{Math|v₁}}, {{Math|v₂}}, and {{Math|v₃}} and rotational modes {{Math|v₄}}, {{Math|v₅}}, and {{Math|v₆}}.{{Cite journal |last1=Steindler |first1=Martin J. |last2=Gunther |first2=William H. |date=August 1964 |title=The absorption spectrum of plutonium hexafluoride |url=https://linkinghub.elsevier.com/retrieve/pii/0371195164801594 |journal=Spectrochimica Acta |language=en |volume=20 |issue=8 |pages=1319–1322 |doi=10.1016/0371-1951(64)80159-4|bibcode=1964AcSpe..20.1319S }}{{Cite journal |last1=Walters |first1=R.T. |last2=Briesmeister |first2=R.A. |date=January 1984 |title=Absorption spectrum of plutonium hexafluoride in the 3000–9000 Å spectral region |url=https://linkinghub.elsevier.com/retrieve/pii/0584853984801087 |journal=Spectrochimica Acta Part A: Molecular Spectroscopy |language=en |volume=40 |issue=7 |pages=587–589 |doi=10.1016/0584-8539(84)80108-7|bibcode=1984AcSpA..40..587W }} The {{Chem|Pu|F|6}} Raman spectrum cannot be observed, because irradiation at 564.1 nm induces photochemical decomposition.{{cite tech report|first1=N. J.|last1=Hawkins|first2=H. C.|last2=Mattraw|first3=W. W.|last3=Sabol|title=Infrared Spectrum and Thermodynamic Properties of PuF₆|institution=Knolls Atomic Power Laboratory|number=KAPL-1007|date=24 May 1954}} Irradation at 532 nm induces fluorescence at 1900 nm and 4800 nm; irradiation at 1064 nm induces fluorescence about 2300 nm.{{Cite journal |last1=Beitz |first1=James V. |last2=Williams |first2=Clayton W. |last3=Carnall |first3=W. T. |date=March 1982 |title=Fluorescence studies of neptunium and plutonium hexafluoride vapors |url=https://pubs.aip.org/aip/jcp/article/76/5/2756-2757/451674 |journal=The Journal of Chemical Physics |language=en |volume=76 |issue=5 |pages=2756–2757 |doi=10.1063/1.443223 |bibcode=1982JChPh..76.2756B |issn=0021-9606}}{{Cite book |last1=Beitz |first1=James V. |url=https://pubs.acs.org/doi/book/10.1021/bk-1983-0216 |title=Plutonium Chemistry |last2=Williams |first2=Clayton W. |last3=Carnall |first3=W. T. |date=1983-05-19 |publisher=American Chemical Society |isbn=978-0-8412-0772-1 |editor-last=Carnall |editor-first=William T. |series=ACS Symposium Series |volume=216 |location=Washington, D.C. |pages=155–172 |language=en |chapter=11. Plutonium Hexafluoride Gas Photophysics and Photochemistry |doi=10.1021/bk-1983-0216.ch011 |editor-last2=Choppin |editor-first2=Gregory R.}}

Plutonium hexafluoride is relatively hard to handle, being very corrosive, poisonous, and prone to auto-radiolysis.{{cite journal|last=Bibler|first=Ned E.|date=August 23, 1979|title=α and β Radiolysis of Plutonium Hexafluoride Vapor|journal=J. Phys. Chem.|volume=83|pages=2179–2186|doi=10.1021/j100480a001|number=17}}{{cite journal|last1=Steindler|first1=M.J.|last2=Steidl|first2=D.V.|last3=Fischer|first3=J.|date=November 1964|title=The decomposition of plutonium hexafluoride by gamma radiation|journal=Journal of Inorganic and Nuclear Chemistry|volume=26|issue=11|pages=1869–1878|doi=10.1016/0022-1902(64)80011-7}}

PuF₆ is stable in dry air, but reacts vigorously with water, including atmospheric moisture, to form plutonium(VI) oxyfluoride and hydrofluoric acid.{{Cite journal|last=Kessie|first=R. W.|date=1967|title=Plutonium and Uranium Hexafluoride Hydrolysis Kinetics|url=https://pubs.acs.org/doi/abs/10.1021/i260021a018|journal=Industrial & Engineering Chemistry Process Design and Development|language=en|volume=6|issue=1|pages=105–111|doi=10.1021/i260021a018|issn=0196-4305}}

:{{chem|Pu|F|6}} + 2 {{chem|H|2|O}} → {{chem|Pu|O|2|F|2}} + 4 {{chem|H|F}}

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.{{Cite journal|last1=Malm|first1=John G.|last2=Weinstock|first2=Bernard|last3=Weaver|first3=E. Eugene|date=1958|title=The Preparation and Properties of NpF₅; a Comparison with PuF₅|url=https://pubs.acs.org/doi/10.1021/j150570a009|journal=The Journal of Physical Chemistry|language=en|volume=62|issue=12|pages=1506–1508|doi=10.1021/j150570a009|issn=0022-3654}}

An important reaction involving PuF₆ is the reduction to plutonium dioxide. Carbon monoxide generated from an oxygen-methane flame can perform the reduction.{{Cite journal |last1=Pokidyshev |first1=A. M. |last2=Tsarenko |first2=I. A. |last3=Serik |first3=V. F. |last4=Sokolov |first4=V. B. |date=October 2003 |title=Reduction of Plutonium Hexafluoride Using Gaseous Reagents |url=http://link.springer.com/10.1023/B:ATEN.0000010988.94533.24 |journal=Atomic Energy |language=en |volume=95 |issue=4 |pages=701–708 |doi=10.1023/B:ATEN.0000010988.94533.24 |s2cid=93145477 |issn=1063-4258}}

Plutonium hexafluoride typically decomposes to plutonium tetrafluoride and fluorine gas. Thermal decomposition does not occur at room temperature,{{Cite journal |last1=Trevorrow |first1=L. E. |last2=Shinn |first2=W. A. |last3=Steunenberg |first3=R. K. |date=March 1961 |title=The Thermal Decomposition of Plutonium Hexafluoride |url=https://pubs.acs.org/doi/abs/10.1021/j100821a003 |journal=The Journal of Physical Chemistry |language=en |volume=65 |issue=3 |pages=398–403 |doi=10.1021/j100821a003 |issn=0022-3654}}{{Cite journal |last1=Fischer |first1=J. |last2=Trevorrow |first2=L. |last3=Shinn |first3=W. |date=October 1961 |title=The Kinetics and Mechanism of the Thermal Decomposition of Plutonium Hexafluoride |url=https://pubs.acs.org/doi/abs/10.1021/j100827a036 |journal=The Journal of Physical Chemistry |language=en |volume=65 |issue=10 |pages=1843–1846 |doi=10.1021/j100827a036 |issn=0022-3654}} but proceeds very quickly at 280 °C. 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.

• {{harvnb|Steindler|1963}}
• {{Cite tech report|last1=Wagner|first1=R. P.|first2=W. A.|last2=Shinn|first3=J.|last3=Fischer|first4=Martin J.|last4=Steindler|title=Laboratory Investigations in Support of Fluid-bed Fluoride Volatility Processes|volume=VII: The Decomposition of Gaseous Plutonium Hexafluoride by Alpha Radiation|institution=Argonne National Laboratory|number=ANL-7013|date=1 May 1963|doi=10.2172/4628896}} 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.Morse, L. R. (2005), "PuF₆ gas pressure in aged cylinders" (personal communication to D. L. Clark), Los Alamos, NM.{{rs?|date=April 2023}}

Likewise, the compound is photosensitive, decomposing (possibly to plutonium pentafluoride and fluorine) under laser irradiation at a wavelength of less than 520 nm.{{cite patent|country=US|number=4670239|title=Photochemical Preparation of Plutonium Pentafluoride|status=|pubdate=June 2, 1987|fdate=December 20, 1977|invent1=Sherman W. Rabideau|invent2=George M. Campbell|assign1=The United States of America|url=http://www.freepatentsonline.com/4670239.html|postscript=,}} but see also {{cite journal |journal=Journal of Fluorine Chemistry |doi=10.1016/S0022-1139(00)80734-4 |first1=E. A. |last1=Lobikov |first2=V. N. |last2=Prusakov |first3=V. F. |last3=Serik |title=Plutonium Hexafluoride Decomposition under the Action of Laser Radiation |date=August–September 1992 |volume=58 |issue=2–3 |page=277|postscript=,}} in which the decay product is identified as tetrafluoride instead.

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

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 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.{{Cite journal|last1=Mills|first1=T.R.|last2=Reese|first2=L.W.|date=1994|title=Separation of plutonium and americium by low-temperature fluorination|url=https://linkinghub.elsevier.com/retrieve/pii/0925838894909318|journal=Journal of Alloys and Compounds|language=en|volume=213-214|pages=360–362|doi=10.1016/0925-8388(94)90931-8}}

Separation of the hexafluorides of uranium and plutonium is also important in the reprocessing of nuclear waste.

• {{cite patent|country=US|number=3708568A|title=Removal of Plutonium from Plutonium Hexafluoride-Uranium Hexafluoride Mixtures|fdate=1970-10-20|pubdate=1973-01-02|assignee=Atomic Energy Commission|invent1=Gilliher, W.|invent2=Harris, R.|invent3=Ledoux, R.}}
• {{cite patent|country=US|number=4172114A|title=Method for purifying plutonium hexafluoride|fdate=1977-08-24|pubdate=1979-10-23|assignee=Japan Atomic Energy Research Institute|inventor=Mitsuhiro Nishimura et al}}

{{Cite journal|last1=Moser|first1=W.Scott|last2=Navratil|first2=James D.|date=1984|title=Review of major plutonium pyrochemical technology|url=https://linkinghub.elsevier.com/retrieve/pii/0022508884900626|journal=Journal of the Less Common Metals|language=en|volume=100|pages=171–187|doi=10.1016/0022-5088(84)90062-6|osti=6168468 }}{{Cite journal|last1=Drobyshevskii|first1=Yu. V.|last2=Ezhov|first2=V. K.|last3=Lobikov|first3=E. A.|last4=Prusakov|first4=V. N.|last5=Serik|first5=V. F.|last6=Sokolov|first6=V. B.|date=2002|title=Application of Physical Methods for Reducing Plutonium Hexafluoride|url=http://link.springer.com/10.1023/A:1020840716387|journal=Atomic Energy|volume=93|issue=1|pages=578–588|doi=10.1023/A:1020840716387|s2cid=100100314}} 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.{{Cite book|url=https://www.nap.edu/read/5538/chapter/5|title=Evaluation of the U.S. Department of Energy's Alternatives for the Removal and Disposition of Molten Salt Reactor Experiment Fluoride Salts|publisher=National Academies Press|location=Washington, DC|via=NAP.edu|year=1997|doi=10.17226/5538|isbn=978-0-309-05684-7|language=en}}

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.{{Cite tech report|last=Seaborg|first=G. T.|author-link=Glenn Seaborg|date=1942|institution=University of Chicago Metallurgical Laboratory|number=CN-125}} The Second World War then interrupted the publication of further research.{{Cite tech report|last=Steindler|first=Martin J.|title=Laboratory Investigations in Support of Fluid-bed Fluoride Volatility Processes|volume=II: The Properties of Plutonium Hexafluoride|institution=Argonne National Laboratory|number=ANL-6753|date=1 August 1963|doi=10.2172/4170539}}

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.{{Cite tech report|first1=H. S.|last1=Brown|first2=O. F.|last2=Hill|first3=A. H.|last3=Jaffay|institution=University of Chicago Metallurgical Laboratory|number=CN-343|date=1942}} 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.{{Cite tech report|last1=Brown|first1=H. S.|first2=O. F.|last2=Hill|institution=University of Chicago Metallurgical Laboratory|number=CN-363|date=12 November 1942}} 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.{{Cite tech report|first1=N. R.|last1=Davidson|first2=J. J.|last2=Katz|first3=O. F.|last3=Orlemann|institution=University of Chicago Metallurgical Laboratory|number=CN-987|date=11 October 1943}}

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.{{Cite tech report|first1=R. W.|last1=Fisher|first2=F.|last2=Vaslow|first3=A. D.|last3=Tevebaugh|institution=Iowa State College|number=CN-1783|date=10 August 1944}}

In 1944, {{Ill|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.{{Cite tech report|first=Alan E.|last=Florin|institution=University of Chicago Metallurgical Laboratory|number=CN-2159|date=1 October 1944}}

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

• {{cite tech report|first1=L.|last1=Brewer|first2=L.|last2=Bromley|first3=P. W.|last3=Gilles|first4=N. L.|last4=Lofgren|institution=University of California Radiation Laboratory|number=CN-3300|date=10 October 1945}}
• {{cite tech report|first1=L.|last1=Brewer|first2=L.|last2=Bromley|first3=P. W.|last3=Gilles|first4=N. L.|last4=Lofgren|institution=University of California Radiation Laboratory|number=CN-3378|date=1 December 1945}}
• {{cite tech report|first1=L.|last1=Brewer|first2=L.|last2=Bromley|first3=P. W.|last3=Gilles|first4=N. L.|last4=Lofgren|title=The Higher Fluorides of Plutonium|institution=University of California Radiation Laboratory|number=UCRL-633|date=20 March 1950}}

In 1950, Florin's efforts finally yielded the synthesis,{{cite tech report|first=Alan E.|last=Florin|title=Plutonium Hexafluoride, Plutonium (VI) Oxyfluoride: Preparation, Identification, and Some Properties|institution=Los Alamos Scientific Laboratory|number=LAMS-1118|date=16 October 1950|url=http://www.fas.org/sgp/othergov/doe/lanl/lib-www/la-pubs/00424522.pdf}} and improved thermodynamic data and a new apparatus for its production soon followed. Around the same time, British workers also developed a method for the production of PuF₆.{{cite tech report|author=Mandleberg, C. J. |display-authors=etal |institution=Atomic Energy Research Establishment|number=C/R-157|date=1952}}

{{Reflist}}

{{Hexafluorides}}

{{Plutonium compounds}}

{{fluorine compounds}}

{{Actinide halides}}

{{Use dmy dates|date=March 2018}}

Category:Plutonium compounds

Category:Hexafluorides

Category:Octahedral compounds

Category:Actinide halides

Category:Nuclear materials