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Reactivity of transition metal fluorides—XII
J. inorg, nucl Chem. VoI. 43. pp. 123t-12~8. 1981 Printed in Great Britain All rights reserved
0022-1902181/061231~98502.00/0 Copyright © 1981 Pergamon Pres!, ltd
REACTIVITY OF TRANSITION METAL FLUOR1DES--XII PLUTONIUM HEXAFLUORIDEt R. C. BURNS and T. A. O'DONNELL* Department of InorganicChemistry,Universityof Melbourne. Parkville,Victoria 3052, Australia and C. H. RANDALL AustralianAtomic Energy CommissionResearch Establishment,Lucas Heights, N.S.W. 2232, Australia
(Received 12 September 1980; receivedfor publication 16 October I980) Abstract--In a continuing study of the relative chemical reactivities of the higher fluorides of the d- and f-transition elements, oxidation-reductionreactions of plutonium hexafluoridewith a selected series of non-metal fluorides in low oxidationstates have been examined.This investigationhas also includedreactions with hydrogen and with various covalent chloridesand bromides.The results of this study have been correlated with those from previous investigations,and the reactivity of plutonium hexafluorideis considered in the light of plutonium as an f-transition element. Specialtechniques were developedfor the study of compoundsof this radioactiveelement,and a new procedurefor the preparationof plutoniumhexafluoridehas been devisedwhich requiresapparatus very much simpler than previouslyreported. INTRODUCTION Many studies of the relative chemical reactivities of the higher fluorides of the d-transition metals have established that the oxidant strength of a particular fluoride increases with increasing atomic number across any row of the Periodic Classification, but decreases with increasing atomic number in any group [1-3]. In addition, it has also been shown that significant differences develop progressively between successive pairs of second and third row elements moving from zirconium and hafnium across the transition metals series. The above trends, as established by chemical studies, are in agreement with trends in thermodynamic and other physical properties of the compounds [1-4]. Historically, the early members of the actinides were thought to form the beginnings of a new d-transition metal row because of the similarities of the formulae of the known compounds and many of their chemical properties to those of their respective d-transition metal congeners. Thus, for example, uranium was originally regarded as a member of sub-group VI. However, after the preparation of the post-uranium elements and publication of Seaborg's [5, 6] actinide hypothesis this view was not widely held. On the basis of the available physical and chemical evidence, the latter being very limited at the time, Seaborg argued that for actinium fZ = 89) and thorium (Z = 90) the binding energies of the 6d electrons were lower than for the 5f electrons but that at about protactinium ( Z = 91) a crossover occurs, so that for uranium (Z = 92) and the post-uranium elements the order is reversed leading to the progressive development of f-transition metal character in the actinide series. However, there are major differences in chemical behaviour between the earlier actinides and the formally corresponding lanthanides. In the lanthanides themselves the corresponding crossover (Sd vs 4f orbital energies) occurs at the beginning of the lanthanide series, *Author to whom correspondenceshould be addressed. +Part XI: J. lnorg. Nucl. Chem. 42. 1613 (1980). 1231
and more abruptly than is postulated for the actinides. The 4f orbitals do not extend very far from the nuclei by comparison with the s, p and d orbitals of higher principal quantum number. They are very effectiw~ly screened by these outer orbitals and take little or no part in the bonding of compounds of these elements. It is postulated that, for the actinides, relative spatial exte,nsion of the 5f orbitals is much greater than for the 4f and that 5f orbitals are available for compound formation. In earlier studies in this series selected oxidationreduction and halogen-exchange reactions of UF6 and PaF5 were studied[7,8], and comparisons with the formally analogous higher fluorides of the elements of subgroups VI (Cr, Me and W) and V (V, Nb and Ta), respectively, were examined[9,10]. For UF6 and the higher fluorides of sub-group VI the order of oxidant strength was found to be CrF6 ~ UF6 -> MoF6 > WF6 and it was argued that if uranium were a typical d-transition element the order should be CrF6 ~>MoF6 > WF~ > [-JF6. It was proposed that the experimental order of reactivity showed differences in bonding between UF6 and the other hexafluorides, providing chemical evidence for participation of 5f orbitals in the covalent bonding in UF6 and for the classification of uranium as an f-transition element. Similarly, for PaF5 and the higher fluorides of sub-group V, the experimental evidence indicates that while protactinium behaves somewhat like a d-transition element, the observed differences indicate that this element, like uranium, forms part of a series with properties different from both the d-transition elements and the lanthanides. In order to study further the effects of a d- to fcrossover and the emergence of f-transition metal character in the post-uranium elements, within the con-. text of studies of the type previously undertaken, it was decided to compare the chemical reactivity of PuF6 with that of the actinide fluorides UF6 and PaF5 and with OsF,, and RuF6, the formally analogous d-transition metal congeners of PuF6, which were also recently examined [2, 3]. In most of the earlier work on the study of tlhe relative chemical reactivities of the higher metal
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R. C. BURNS et al.
fluorides conducted at Melbourne, reactions were studied using fairly large quantities of fluorides (0.5-1g), employing apparatus and procedures generally following those described by Canterford and O'Donnell[ll]. However, in this study all work was carried out in glove-boxes because of the radioactive nature of the compounds and, as a safety requirement, reactions were studied using only about 50mg of hexafluoride. Thus special apparatus and procedures were developed for use in this investigation and, moreover, a simple method for the preparation of PuF6 was devised. The reliability of the procedures and apparatus was checked prior to the current study by investigating the reactivity of RuF6 which, at the time, was expected to be fairly similar to PuF6 in its chemical reactions with those selected reagents used in these investigations. When the technological importance of PuF6 in nuclear fuel processing and related contexts is considered, it is surprising that so little of the chemistry of this hexafluoride has been reported. Although its decomposition reaction and its interaction with alkali and alkaline earth metal fluorides are well-documented, few other reactions have been studied in any systematic way. A summary of reported results can be found in reviews by Steindler [12] and Vanka [13], and also as discussed by Cleveland [14]. EXPERIMENTAL
Apparatus and materials. All work was performed at the laboratories of the Australian Atomic Energy Research Establishment, Lucas Heights, N.S.W., on a vacuum line constructed from monel, stainless steel, and KeI-F. In general, the apparatus was similar to that described by Canterford and O'Donnell[ll], but had been adapted for use under glove-box conditions, as indicated above. Exentsive use was made of Whitey SS-1KS4, SS-IVS4, SS-14DKM4-S4 valves and valves constructed from Kel-F[15], while Whitey Sample Cylinders were used for the storage and handling of fluorine, and in the preparation of PuF6. Reactions were carried out in small, specially designed KeI-F reaction tubes (of capacity 3 cm3),which could, be fitted with an expansion bulb (ca. 25 cm3) if a high pressure was anticipated as a result of reaction, or for handling gases such as hydrogen as reactants. Preparation of plutonium hexafluoride. The apparatus used for preparation of this thermally unstable hexafluoride has been described previously[3] where it was used successfully for the preparation of RuF6. In principle, the method of preparation involves immediate cooling of the hexafluoride upon formation as in earlier procedures which relied on very complex, costly and sophisticated equipment but, unlike those earlier methods, requires only simple apparatus based on a stainless steel cylinder, the base of which can be heated and which is provided with an external cooling trough which can be maintained at -78°C fitted to the upper part of the cylinder (for details see Ref. [3]). In this work PuO2 (ca. 0.8g, supplied by the AAEC; >96% 239pu) was first treated in the reaction cylinder with excess hydrogen fluoride at 450-550°C for 2 hr to produce a mixture of lower fluorides of plutonium (predominantly PuF3), water and oxygen. The excess hydrogen fluoride and other products were then removed by evacuation. This procedure was repeated once. Fluorine, slightly in excess of the amount required to form 50 mg of PuF6, was admitted to the cylinder and PuF6 was prepared by heating the lower fluoride(s) and fluorine at 500-550°C for 1 hr. Upon formation, the PuF6 condensed in the cooled (-78°C) region of the reactor and, after cooling of the cylinder and removal of the excess fluorine, the coolant was removed from the external trough and the hexafluoride used directly from the vessel. Waste hydrogen fluoride and fluorine were passed through a granulated alumina tower, to give water and oxygen, prior to disposal. Other reagents. Anhydrous hydrogen fluoride (AHF) Matheson, 99.8% min) was distilled on a column similar to that
described by Shamir and Netzer[16], while C1F3 (Matheson, 98% rain) was doubly distilled in a vacuum before use. The perfluoroalkanes, n-C6Fi4 and n-CTFI6 (P.C.R.), were stored over P4Oto and were distilled directly before use. All other reagents were prepared and/or purified as described previously[9], or were of reagent grade purity. Reaction procedures. All volatile reagents were transferred by distillation in a vacuum, while solid materials were handled in a glove bag filled with dry argon. Reactions were carried out by condensing excess reductant, if volatile, onto the hexafluoride at -196°C, whereas for solid reductants the hexafluoride was condensed directly onto excess reagent at -196°C. In either case the reactants were then allowed to warm slowly to room temperature during which time any reaction was noted. Previous studies involving OsF6, IrF6 and RuF6 have shown that some reactions of these hexafluorides are best studied under controlled conditions at low temperatures, using an inert solvent as a "thermal sink';, because of their extremely exothermic nature[2, 3]. Now PuF6 is thermally unstable at room temperature and so would be expected to be extremely reactive[l]. Therefore the reaction of PuF6 with PF3 was conducted in liquid AHF and the majority of the reactions of PuF6 with CCI4, BCI3 and BBr3 were conducted in an n-perfluoroalkane (either n-C6FI4 or n-CTFI6) solvent. The procedure employed for this type of reaction has been described previously [21. Identification o/products. Upon completion of a reaction the volatile species were removed and, after fractionation if necessary, were identified by their physical properties and IR spectra. Gas-phase IR spectra were recorded on either a Perkin-Elmer 225 grating IR spectrophotometer or a Perkin-Elmer Model 21 double beam IR spectrophotometer, the samples being contained in a monel-bodied cell fitted with detachable AgC1 windows. Solid products of reaction were identified by X-ray powder photography and by their physical properties. Samples for X-ray analysis were loaded into ca. 0.4 mm o.d. Pyrex capillaries under anhydrous conditions and photographs were taken using a Philips Debye-Scherrer camera of diameter 11.46cm, on a Philips Industries Ltd. X-ray generator, Model PW 1008/30. Nickel-filtered Cu-K~ radiation (Ks = 1.5418,~) was employed for all samples and exposures required about 12 hr. RESULTS AND DISCUSSION
A summary of results of the reactions of PuF6 that were studied in this investigation is given in Table 1. As observed in previous studies of other hexafluorides, the reactions of PuF6 with the lower fluorides of the nonmetals, when they occurred, involved simple oxidationreduction, with oxidation of the lower fluoride to the corresponding higher fluoride and reduction of the hexafluoride to oxidation state IV. Also, in some reactions evidence was found for adduct formation between the products of reaction and this is discussed in detail below. Other reactions of PuF6, for example those with hydrogen and the covalent chlorides and bromides, also involved oxidation-reduction and, where possible, comparisons are made with the related uranium and neptunium systems.
Oxidation-reduction reactions o[ plutonium hexafluoride In anhydrous HF, PuF6 was found to oxidize PF3 to PF5 well below room temperature, with reduction of the hexafluoride to PuF4 as shown by X-ray powder photography. Arsenic trifluoride also reacted with PuF6 below room temperature, but in this case no free AsF5 could be detected as a gaseous product. Furthermore, the solid product of reaction was pale orange in colour, whereas PuF4 is light brown to pale pink depending on the method of preparation. X-ray examination of the solid product did show, however, a rather weak and diffuse pattern of PuF4. The solid product exhibited no great solubility in anhydrous HF and showed no evidence of
Reactivityof transition metal fluorides--XII Fable I. Oxidation-reductionreactions of plutonium hexafluoride Reactant
Products
I>F3:AHF (a)
P u F . , PF~
AsI:,
PuFf. PuF..AsF~ adduct
Si~F
PuF4, SbF5
BiF,
No reaction
C!F~
No reaction
R~F.
No reaction
!t2/AEF
Pu(IV),
Ci~I,,,'PFA
PuF., CF3CI, CF2C12, C12
BI?!~/PFA
PuF.,
BBr~/PFA
PuFf, BF3 (c) , Br2
gBr,
PuFf, BF3, Br2
Pu(V)
(b)
BF3 (c) , C12
(a)
Key:
AHF = anhydrous hydrogen fluoride;
(b)
For a dlscussJon of this reaction,
~c>
No mixed halides of boron were detected.
PFA = n-perfluoroalkane
(n-C6Fx~ or n-CTF~6). see text.
liberating AsFs on heating at 90°C for l hr. A similar result was also obtained when the reaction was repeated using approximately stoichiometric quantities of reactants. Interestingly, it should also be noted that the reaction between UF6 and AsF3, although it does not proceed readily, gives UF4 as one product, and, like the above reaction, no evidence of free AsF5 as a gaseous product [7, 17]. Recently, it has been shown that UF4 may be solubilized to some extent in anhydrous HF by reaction with AsF5 over many days[18]. After decantation and removal of the solvent a solid product of empirical composition approximating to UF4'2AsF5 remains and it is suggested that the residue might be UF2(AsF,)2[18]. It is assumed here that an involatile adduct is formed between PuF4 and AsFs, but its composition is unknown. The reaction of PuF6 and SbF3 occurred just below room temperature to give PuF4 and SbFs. The formation of both products was inferred by identification of the salts /3:K2PuF6 (isomorphous with /3,-K2UF6 and ~,K2NpF6) and KSbF6 by X-ray powder photography, after conversion of the reaction products to the above salts by treatment with potassium fluoride in anhydrous H F In view of the probable complexity of the actual reaction mixtures, no attempt was made to investigate if there was any adduct formation between the products of reaction and excess SbF3, and/or between the products themselves, although the latter seems highly probable, following the AsF3 reaction. Ihe two remaining lower fluorides, BiF3 and C1F3, showed no reaction with PuF6 on standing for several hours. In the latter case this is not surprising as CIF3 may be used to prepare PuF6 by fluorination of PuF4 at high temperatures[12, 14]. Also, no reaction was observed to occur between PuF~ and the highest thermally stable fluoride of ruthenium, RuFs. In previous investigations in this series the reactions of a particular fluoride with the reagents SF4, SeF4 and CS2 have also been used as a guide to relative chemical
1233
reactivity. However, it should be noted that prior to t]his study PuF6 has been shown to oxidize SF4 to SF6, with reduction of the hexafluoride to PuF4 [19-21], while CS~, and PuF6 react to give PuF4, with CF4, SF4 and SF6 as gaseous products[22]. As SeF4 has been found to be comparable to SF4, but stronger than SbF3, in reducing ability towards related hexafluorides, reaction with this potential reductant was not investigated in this study. The reaction between PuF6 and hydrogen has previously been studied by Hawkins[23], who observed that reaction proceeded slowly at room temperature to give HF and a solid which was presumed to be PuF4, perhaps with the intermediate formation of the uncharacterized pentafluoride. In two related studies, Levy and Wilson[24] and Asprey and Paine[25], found that UF6 and hydrogen react in anhydrous HF to give /3-UF5 but, in the presence of platinum foil or UVradiation, UF4 is formed. In the present study the reaction between PuF6 and hydrogen was investigated in anhydrous HF. Reaction between PuF6 and an excess of hydrogen occurred slowly at room temperature to give an insoluble white to pale pink solid, the reaction being allowed to continue until all of the PuF6 had been consumed. The solid reaction product was somewhat sensitive to moisture, and gave some Pu(VI) (as PuO:~~+) on addition to weakly acidic or neutral solution, in addition to insoluble PuF4. This suggests that some Pu(V) formed on reduction which, like UFs, disproportionates into the dioxo-cation (VI) and the insoluble tetrafluoride in aqueous solution in the presence of F ion, a strong complexing agent. Unfortunately, the solid was amorphous to X-rays. The reaction was repeated, with identical results, and on addition of a solution of excess rubidium fluoride in anhydrous HF to a sample of the reaction product a light-pink solid was isolated. This material turned brown on exposure to moisture, with evidence for some Pu(VI) on dissolution in aqueous solution, which is uncharacteristic of any known rubidium (or other alkali metal) complex fluoride of Pu(IV). In another reaction, using approximately stoichiometric quantities of reagents, the reaction was terminated with a little PuF~ still present, and a small amount of an involatile crimson-red material was also observed, which exhibited properties similar 1:o those described above. The formation of a brick-red material has previously been observed during fluorination of PuF4 and was tentatively identified as Pu4F~v by chemical analysis and X-ray powder photography[26]. Although no conclusive evidence for PuF~ was obtained in this study it would appear that some plutonium is in oxidation state intermediate between VI and IV on reduction of PuF6 by hydrogen, although Pu(IV) appears to be the primary product, particularly in the presence of excess reductant. Moreover, it is suggested that this reaction might be strongly controlled by kinetic factors such that any Pu(V) that is formed may be reduced at a comparaNe rate to that of PuF6. Thus the reaction could proceed through the intermediate formation of mixed oxidation state products, analogous to U2F9and U , F : . ultimately leading to PuF4. The formation of Pu(V), either as PuF~ or as a complex fluoride, was also attempted by fluorine oxidation of PuF4 in anhydrous HF or in an RbF. 2HF melt at 90°C over several hours. However, neither reaction proved successful. It should be noted that under similar conditions UF4 in anhydrous HF is oxidized to /3-UF~ at room temperature, with slow formation of UF~ on
1234
R. C. BURNS et al.
standing, and forms salts oI composition MUF6 (M = Na, Rb, Cs) in the respective melts [27, 28]. In contrast, NpF4 is not oxidized in anhydrous HF at 25°C but does give CsNpF6 in a CsF.2HF melt at 70°C[27]. These results reflect the trend toward progressively decreasing stability, and hence difficulty of attaining, high oxidation states across the actinide row from uranium onwards. Thus complex fluorides of Pu(V), containing rubidium and cesium, have only been prepared by direct fluorination of the respective Pu(IV) salts at 300--400°C for several hours [29].
Halogen-exchange reactions of plutonium hexafluoride In their extensive survey of the reactions of UF6 with various covalent chlorides (and bromides), O'Donnell et al.[7] found that the formation of different products could be rationalized on the basis of a competition between two reactions, depending on their kinetic involvement, viz. simple halogen exchange of UF6 to give UC16 and the appropriate exchange products, and the secondary reaction of UF6 and UCI6 to give UF4 and chlorine (for which AG25oc= -29 kcal/mol of UF6). For excess of the halogen-exchange reagent the final product obtained in each reaction was found to depend on the comparative rates of the two reactions, such that when halogen exchange was fast, e.g. with BCI3 and AIC13, only UC16 was produced but, when halogen exchange was slow, e.g. with CC14 and SiCI4, the UF6 was effectively in excess and reacted with the UC16 as it formed to give UF4. For TiCI4 both products were observed so that the two reactions have comparable rates in this case. For excess of UF6, only UF4 was observed as a reduction product, in accord with the above arguments. On reaction of UF6 with the covalent bromides BBr3 and PBr5 only UF4 was observed, and it was concluded that in these reactions UF6 was reduced directly to UF4, reflecting the thermodynamic instability of the higher uranium bromides [30]. In this analogous study the reactions of PuF6 with the covalent chlorides, BCI3 and CCI4, reagents which had shown fast and slow reactions with UF6, respectively, and with the covalent bromide, BBr3 were investigated. The reaction of PuF6 and CC14 in the gas phase has previously been studied by Sabol and Hawkins[31,32], who observed that reaction proceeded rapidly at room temperature to form CFCI3 and CF2Clz, while PuF6 and CFCI3 react to give CF2C12 and CF3C1. For neither reaction was the reduction product reported, but was presumed to be PuF4 in each case. In this study, reaction between PuF6 and CC14 was observed just below room temperature in n-perfluoroalkane to give CF2C12, CF3CI and chlorine as gaseous products, while the solid product was identified as PuF4 by X-ray powder photography. Reaction of PuF6 with either BCI3 or BBr3 in nperfluoroalkane occurred well below room temperature to give PuF4 and BF3 in each reaction, and chlorine or bromine, respectively. Also, reaction of PuF6 and BBr3 in the absence of n-perfluoroalkane gave a violet-blue solid which, based on its colour, was presumably PuF3. It would appear, therefore, that all of the reactions of PuF6 with the covalent chlorides and bromides under conditions where the heat of reaction can be dissipated effectively involve simple oxidation-reduction, such as, 3PuF6 + 2BX3 P F A > 3PuF4+2BF3+3X2; (X = CI, Br).
Under less controlled conditions, a typical reaction is PuF6 + BBr3 ~
3 PuF3 + BF3 + ~ Br2.
The highest thermodynamically stable solid chlorides and bromides of plutonium are the trihalides, although PuC14 has been observed in the gas phase at high temperatures[14], while both tetrahalides are found at room temperature in solid complexes with various amides, dimethyl sulphoxide or triphenylphosphine oxide [14]. Hence, the formation of lower fluorides in the above reactions, probably by direct reduction of the hexafluoride, is consistent with the thermodynamic instability of the "higher" chlorides and bromides of plutonium. It should be noted that Peacock and Edelstein[33, 34] have reported that NpF6 and BC13react to give NpF4, BF3 and chlorine, so that an analogous argument is applicable to the neptunium system as well. In the reactions of PuF6 with the reagents listed above, and others[12], it should be noted that the hexafluoride is reduced almost exclusively to the tetrafluoride. The formation of a solid product, like PuF4, in such a wide variety of reactions may be attributed to either or both of the following factors; a very stable oxidation state or, for a compound which is predominantly ionic in character, a very high lattice energy. For plutonium, it would appear that a combination of both factors allows the product PuF4 to dominate the reduction reactions of the hexafluoride. In general, Pu(IV) is the most stable oxidation state (see also below) in aqueous solution, although because of the various equilibria and kinetics of conversion from one oxidation state to another, plutonium can form appreciable concentrations of the IIl, IV, V and VI oxidation states even in the same solution, depending on the reaction conditions[14]. Plutonium tetrafluoride, like all of the actinide tetrafluorides, is also highly ionic in character as demonstrated by its melting point of 1037°C[30], whereas covalent tetrafluorides generally boil below ca. 300°C. Thus the lattice energy of PuF4 is very high, with a calculated value of ca. -2060kcal/mol[35], and is appreciably larger than the lattice energies of, for example, NaF (-216 kcal/mol) or MgF2 (-695 kcal/mol)[36]. Hence PuF4 is expected to show considerable stability and, consequently, is the preferred reduction product of PuF6. However, it should also be noted that under more vigorous reaction conditions, reduction of PuF6 to PuF3 can occur as shown in its reactions with BBr~ and with iodine [37]. In the related uranium and neptunium systems oxidation states greater than IV are relatively more stable than found in the plutonium system--the most stable oxidation states in aqueous solution are U(VI) and Np(V), respectively--so that similarities in the reaction patterns of the hexafluorides and other fluorides of these elements are to a large degree a reflection of the very high lattice energies of the tetrafluorides.
Comparative reactivity o[ plutonium hexafluoride and related hexafluorides As indicated in the Introduction to this paper previous studies in this series have compared the chemical reactivities of the higher actinide fluorides, PaF5 and UF6, with those of their formally analogous d-transition metal counterparts[7-10]. These studies demonstrated that PaF5 was a considerably stronger oxidant than NbF5 or TaFs, which exhibit similar chemical reactivities them-
Reactivity of transition metal fluorides--XII selves, but was a much weaker oxidant than its first row transition metal congener, VFs, Similarly, UF6 was shown to be a marginally stronger oxidant than MoF6 which was, in turn, somewhat stronger than WF6; however, like PaFs, UF6 was a very much weaker oxidant than its first row transition metal counterpart, in this case CrF~. Hence. when protactinium and uranium are considered in their historical positions as members of sub-groups V and VI, i.e. as d-transition elements, the chemical reactivities of PaF~ and UF6 are anomalous as there should be a marked decrease in reactivity of the higher fluorides with increase in atomic number and they should not be classified as typical d-transition metals. Consideration of the thermodynamic stabilities of the binarv fluorides of these elements also supports the above argument. Thus, while vanadium forms fluorides from the highly reactive VF~ to VF2, and niobium forms NbF,, NbF4 and the intermediate fluoride, NbF2.5--the two latter fluorides being produced only by high temperature reduction of the pentafluoride by the metal-there is no reliable evidence for the formation of any solid fluoride lower than TaFs. The very existence of PaF4 in addition to PaFs indicates that protactinium does not follow the trend expected for the d-transition metals, that is, an increase in the relative stability of the highest oxidation state down the group. Similarly, in sub-gruop V1, although chromium forms the unstable CrF6, simple binary or intermediate fluorides from CrF5 to CrF~ may be easily prepared, while molybdenum and tungsten form stable hexafluorides but less stable lower fluorides down to MoF, and WF4,respectively. Uranium, on the other hand, forms stable fluorides, including intermediate fluorides, ranging from UF6 to UF3. the relative chemical reactivities and thermodynamic stabilities of the higher fluorides therefore show that neither protactinium nor uranium behaves chemically as a d-transition element. This, it has been suggested, provides chemical evidence for differences in bonding between the respective actinide and d-transition metal fluorides and, consequently, within the context of Seaborg's actinide hypothesis[5,6], it has been proposed that 5f orbitals are available for bonding in these compounds. Hence for example, d2sf 3 hybrid orbitals could participate with d2sp 3 hybrids in the bonding of UF,[71. Importantly, just as the reactivity and thermodynamic stability of the higher fluorides of protactinium and uranium do not fit into the pattern shown by the d-transition elements, from a chemical viewpoint these elements differ markedly from their formal lanthanide counterparts. Thus, the chemistry of both praseodymium and neodymium is dominated by oxidation state llI. although compounds of Pr(IV) can be prepared with some difficulty. Before discussing the chemical reactivity of PuF6 in relation to its d-transition metal congeners and to the olher higher fluorides of the actinides, it is expedient to iook at various chemical and thermodynamic properties of the higher fluorides of the elements of sub-groups 11I, IV and VII. The chemical and physical properties of AcF3 closely resemble those of ScF3, YF3 and LaF3. Indeed, only differences in the appropriate cationic radii are required to account for variations in the properties of these compounds, and each element is assigned a ground-state electronic configuration of (n- 1) d~ns 2. For thorium and the elements of sub-group IV, therIN(
\tl] a l Nt! i* ]
1235
modynamic stability considerations indicate that, whereas tetrafluorides and lower fluorides are known for titanium and zirconium, hafnium and thorium form HfF4 and ThF4 only. Useful comparisons may also be extended to the other halides and here, it is noted that titanium, zirconium and hafnium form various simple and intermediate halides ranging from tetrahalides to monohalides, the lower halides being relatively easily prepared for titanium and zirconium, but tess so for hafnium. Thorium, on the other hand, forms only tetrahalides and, under very strongly reducing conditions, a di-iodide and, less-well characterized, a tri-iodide. However, these are not regarded as simple binary compounds but are metallic compounds with a formulation of Th(IV)I2e2 in the case of the di-iodide [38]. It is therefore apparent that oxidation state IV increases in stabilit} from titanium to thorium, and this is in accord with the assignment of a ground-state electronic configuration of (n-l)d2ns 2 for these elements. Moreover. thorium differs in a significant way from the lanthanide cerium, which has the accepted configuration 4f26s 2, and which exhibits oxidation states of III and IV in both the solid state and in solution. There is no reported solution chemistry of Th(III). At this point in the discussion we see in the actinide row typical d-transition metal behaviour for actinium and thorium but, as described above, for protactinium and, to a greater extent, for uranium, we begin to see in the series the emergence of properties different fi'om both the d-transition elements and the lanthanides. It would therefore appear that protactinium is just about at the energy crossover point of the 6d and 5f electrons as postulated by Seaborg. Proceeding to neptunium and the sub-group VII ,elements, even greater differences between this actinide ;and its formal d-transition metal congeners may be noted. For the d-transition elements increasing stability of oxidation state VII with increasing atomic number results in the existence of the stable binary fluoride. ReFT. If neptunium were considered a d-transition element then a stable heptafluoride should be formed even more readily than for rhenium. In fact, the highest binary fluoride of neptunium is NpF6, and easily attainable stable fluorides are formed down to the trifluoride. Furthermore, attempts to prepare even an oxide fluoride of Np(VII) have so far failed [33]. For plutonium and its d-transition metal congeners the differences are further magnified. Thus, there are no binary or oxide fluorides of plutonium that correspond to the oxidation state VIII and VII osmium species, OsO3F2, OsO2F3, OsOF5 and the rather unstable ()sF:,. Plutonium itself forms various binary, intermediate and oxide fluorides including the thermally unstable PuFf,. PuO2F,~ and the recently prepared PuOF4139]. and the very stable lower fluorides, PuF4 and PuF3. There is also evidence that an unstable and highly reactive pentafluoride exists, for which some supportive evidence was also obtained in this study. From a chemical point of view, consideration of the relative reactivities of the hexafluorides demonstrates the marked difference between plutonium and the trend established by its d-transition metal congeners. For the d-transition metal fluorides it has been shown that, while RuF6 and OsF6 oxidize PF3 (in anhydrous HF) and AsF, to their respective pentafluorides, RuF~ is reduced to Ru(IV) in both reactions but OsF6 gives Os(IV) and Os(V), respectively[2,3/. Furthermore, RuB, oxidizes
1236
R. C. BURNS et al.
SbF3 to SbF~, being reduced in turn to Ru(V), while
OsF6 does not react with this lower fluoride [2]. Osmium hexafluoride is, therefore, a considerably weaker oxidant than RuF6. Complementary evidence for this trend is provided by a comparison of the reactions of the hexafluorides with water in HF[40] and with boric oxide in an inert solvent (n-C6F~4)[41]. Ruthenium hexafluoride oxidizes both reagents, with reduction to Ru(V), while OsF6 undergoes simple oxygen-fluorine exchange to give OsOF4, with no change in the formal oxidation state of osmium. In comparison with the trend established above, PuF6 oxidizes PF3 (in anhydrous HF), AsF3 and SbF3 to their respective pentafluorides, with reduction of the hexafluoride to Pu(IV) in each case. However, PuF6 shows no reaction with BiF3 or C1F3, unlike RuF6 which reacts with both of these lower fluorides [3]. Based on its thermal instability, RuF6 should also be capable of oxidizing PuF4 to PuF6, as does PtF6142], although this reaction was not investigated in the present study. If this reaction does proceed then, conversely, PuF6 should not be able to oxidize RuF5 to RuF6, as shown in this work. Therefore, from a comparison of all of the above reactions, it is apparent that the order of oxidizing strength of the hexafluorides is RuF6 > PuF6 > OsF6, implying that PuF6 is considerably more reactive than would be expected if plutonium were regarded as a d-transition element. It could be argued, therefore, as will be discussed below, that 5f orbitals make some contribution to the covalent bonding in PuF6, just as previously suggested for PaF5 and UF6 on the basis of similar studies in this series[7, 8]. As is evident from the foregoing discussion, the thermodynamic stabilities of the higher fluorides decrease across the actinide row, and this is reflected in a corresponding increase in chemical reactivity. For example, PaF5 slowly oxidizes PF3 to PFs, being itself reduced to PaF4, but shows no reaction with AsF3. On the other hand, UF6 in a Kel-F reaction tube where heat dissipation is poor reacts readily with PF3 tO give PF5 and UF4. Reduction of UF6 in a metal reactor at -78°C produces/3-UF5143]. It is assumed that, under the thermal reaction conditions in a KeI-F reactor which is itself at room temperature, UF5 disproportionates to UF4 and UF6 which is then reduced by PF~. Uranium hexafluoride reacts to some extent with AsF3, but fails to oxidize either SbF3 or SF4 (also see [21]). Few reactions have been reported for NpF6, but it is known that this hexafluoride oxidizes PF3 tO PFs, with the formation of a-NpF5 under the conditions given above for preparation of/3-UF5144]. It reacts readily with bromine below room temperature to give NpF4 and BrF3145]. In contrast UF6 shows no reaction with bromine, or with BrF3, under similar conditions [30]. In fact, both NpF6 and OF6 may be prepared by oxidation of the respective tetrafluorides with BrF3 and BrF5 at high temperatures[45, 46]. As demonstrated in this study PuF6 oxidizes PF3 in anhydrous HF to PF~ and is reduced directly to the tetrafluoride. Plutonium hexafluoride also oxidizes AsF3, SbF3 and SF4 to their respective higher fluorides, with reduction to Pu(IV) in each case. Furthermore, PuF6 oxidizes both bromine and BrF3 to BrF5 with reduction of the hexafluoride to PuF4[12, 47]. Previously, reaction of metal and non-metal fluorides with CS2 has been ~own to provide a good indication of relative chemical reactivity as different products are formed depending on the oxidizing ability of the parti-
cular fluoride[I-3,48]. For the actinide hexafluorides it has been found that UF6 reacts with CSz at 25°C in the manner characteristic of a moderately strong oxidant, that is, by cleavage of one C-S link, to give UF4, SF4, (CF3)2Sz and (CF3)2S3, although at higher temperatures reaction is somewhat more vigorous and SF6 and CF4 are also formed. In contrast, PuF6 reacts with CS2 at room temperature to give PuF4, CF4, SF4 and SF6, characteristic of an extremely powerful oxidant; the reaction in this case involving the cleavage of all C-S bonds [22, 48]. Unequivocal evidence that PuF6 is a stronger oxidant than UF6 is shown by the fact that PuF6 completely oxidizes UF4 to UF6 above 200°C, with reduction of the hexafluoride to PuF4147]. For the higher fluorides of the actinides oxidant strength follows the order PuF6 > NpF6 > UF6 > PaFs, although comparison with the latter is less exact as it is involatile and the possibility of an homogeneous reaction is ruled out. The trend towards increasing chemical reactivity and the related decrease in thermal stability across the actinide row is further evidenced by the thermal instability of PuF6. For the post-plutonium elements the thermal stability of the higher fluorides decreases continuously across the series. No hexafluoride of americium has yet been prepared (but see [49]) and, although CmF4, BkF4 and CfF4 are known, from curium onwards the most stable oxidation state of the actinides is oxidation state III, As found in the lanthanides [50]. From studies on the relative chemical reactivities of the hexafluorides of plutonium, ruthenium and osmium, as discussed above, it has been shown that PuF6 is substantially more reactive than would be expected if plutonium were considered a d-transition element. These differences in chemical reactivity, which indicate that plutonium does not behave chemically as a d-transition metal, suggest differences in bonding between the hexafluoride of plutonium and its d-transition metal congeners. It is therefore proposed that 5f orbitals could make some contribution to the covalent bonding in PuF6, presumably involving hybrids such as d2sf 3, partially at least. Significantly, it should be noted that from a chemical viewpoint, plutonium does not resemble to any great degree samarium, its formal lanthanide counterpart, except, perhaps, for some aspects of oxidation state III. The chemistry of samarium is, like all of the lanthanides, dominated by this oxidation state, although divalent samarium can be stabilized in some compounds but is, in general, easily oxidized to Sm(III). It should also be noted that there is no evidence for the formation of divalent plutonium. Many authors have considered the possibility of 5[ orbital involvement in the bonding of the early actinides, and there is considerable physical and chemical evidence to support this suggestion. While it is beyond the scope of this paper to discuss such evidence, it is briefly noted that favourable overlap of hybrids which involve f orbitals has been theoretically demonstrated[51], while Coulson and Lester[52] have shown the 5[ orbital involvemen't in bonding in U(VI) species, such as UO22÷, is energetically favourable. Of considerable importance to the arguments described above is the thermodynamic stability and chemical reactivity of the higher fluorides and other halides of the d- and •-transition elements, which have been reviewed by O'Donnell[l]. Of course, it is to be appreciated that chemical studies such as those described above do not in themselves prove that 5f
Reactivity of transition metal tluorides--XlI orbitals are involved in bonding in these compounds. Rather, they serve to show that bonding in the higher actinide fluorides is very different from that in the dtransition metal fluorides where, for example, d2sp 3 hybridization is used to account for the bonding in the hexafluorides. One final point of interest concerns the relative ordering in chemical reactivity of the actinide higher fluorides with regard to their d-transition metal congeners across the Periodic Classification. Thus, while PaF5 is a considerably stronger oxidant than NbF5 or TaFs, UF6 is only slightly stronger than MoF6, which is a stronger oxidant than WF6. Little is known about the relative oxidizing abilities of NpF6 and TcF6, but they would appear to be of comparable reactivity and both are stronger oxidants than ReF~. Furthermore, neither neptunium nor technetium form heptafluorides, as does rhenium. However, in this study we have found that PuF6 is not a stronger oxidant than both of its dtransition metal congeners, but is, rather, a weaker oxidant than its second row congener, RuF6, but a stronger oxidant than its third row congener, OsF6. For any particular reaction, all things being equal, this presumably represents some disparity between the actinide and d-transition metal higher fluorides (including their reduction products) with regard to trends in the endothermic valence state promotion energy and the exothermic intrinsic bond energy (entropy terms are likely to be comparable in these systems, at least for the hexafluorides and their reactions) across the Periodic Classification, reflecting the closeness of the 7s, 7p, 6d and 5[ orbitals in the actinides. However, without data such as ionization potentials or, better, vale,ace state ionization energies as calculated from atomic spectra, together with other reliable physical data on both the actinides and d-transition metals, no quantitative comparisons can be made. For the post-plutonium elements the increasing inertness of the 5f orbitals, both energetically and because of their decreasing spatial extension as a result of poor shielding from the progressively increasing nuclear charge, militates against the formation of high oxidation state compounds so that we quickly see the progressive stabilization of oxidation state III in the remaining elements and even the onset of a chemistry of ~gxidation state 1I in the later members. Acknowledgements--We wish to acknowledge financial assistance from the Australian Atomic Energy Commission and the co-operation of the officers of the AAEC Research Establishment who provided the plutonium dioxide and made possible the glove-box manipulation of the radioactive materials. We also acknowledge receipt of an Australian Postgraduate Research Award (R.C.B.). REFERENCES
I. T. A. O'DonneIl, In Comprehensive Inorganic Chemistry (Edited by J. C. Bailar, H. J. Emel~us, R. S. Nyholm and A. F. Trotman-Dickenson), Vol. 2, pp. 1073-1106. Pergamon Press, Oxford (•973) and refs. therein. 2. R.C. Burns and T. A. O'Donnell, J. lnorg. Nucl. Chem. 42, 1285 (1980). 3, R.C. Burns and T. A. O'Donnell, J. lnorg. Nucl. Chem. 42, 1613 (1980). 4. N. Bartlett, Angew. Chem. Int. Edn. 7, 433 (1968). 5. G. T. Seaborg, In The Chemistry of the Transuranium Elements (Edited by G. T. Seaborg, J. J. Katz and W. M. Manning), Nat. Nucl. Energy Set., Vol. IV, 14B, pp. 14921524. McGraw-Hill, New York (1949). 6. G. F. Seaborg, Nucleonic,~ 5(5), 16 (1949).
1237
7. T. A. O'Donnell, D. F. Stewart and P. Wilson, Inorg. Chem. 5, 1438 (1966). 8. T. A. O'Donnell, A. B. Waugh and C. H. Randall. J. Inorg. Nucl. Chem. 39, 1597 (1977). 9. T. A. O'Donnell and D. F. Stewart, Inorg. Chem. 5, 1434 (1966). 10. J. H. Canterford and T. A. O'Donnell, lnorg. Chem. 5, 1!442 (1966). 11. J. H. Canterford and T. A. O'Donnell, In Technique of Inorganic Chemistry (Edited by H. B. Jonassen and A. Weissberger), Vol. VII, pp. 273-306, Wiley, New York (1968). 12. M. J, Steindler, In Proc. Rocky Flats Fluoride Volatility Conf. (Edited by J. M. Cleveland and M. A. Thompson). USAEC Report CONF-680610, pp. 2-17 (1968). 13. M. Vanka. In Plutonium Hexafluoride: Its Preparation and Properties, Czechoslovak Academy of Sciences, Nuclear Research Institute--Information Centre for Nuclear Energy Publication. pp. 1-62 (1970). 14. J. M. Cleveland, In The Chemistry of Plutonium. Gordon & Breach, New York (1970). 15. Y. A. O'Donnell, Anal. Chem. 43, 977 (1971). 16. J. Shamir and A. Netzer, 3. Sci. Instn., Series 2 1,770 (19,68). 17. D. F. Stewart, Ph.D. Thesis, University of Melbourne, 1965. 18. M. Baluka, N. Edelstein and T. A. O'Donnell, Submitted to lnorg. Chem 19. M. J. Steindler and D. V. Steidl, USAEC Rep. ANL-6183 (1960). 20. M, J. Steindler and D. V. Steidl, USAEC Rep. ANL-6231 (1960). 21. C. E. Johnson, J. Fischer and M. J. Steindler. J. Am. Chem. Soc. 83, 1620 (1961). 22. R. Wagner, Unpublished results quoted in M. J. Steindler, USAEC Rep. ANL-6753 (1963). 23. N. J, Howkins, Results cited in, USAEC Rep. KAPL-1273 (1955) (Classified). 24. J. H, Levy and P. W. Wilson, Austral. J. Chem. 26, 2711 (1973). 25. L. B. Asprey and R. T. Paine, J. Chem. Sot.. Chem. Commun. 920 (1973). 26. C. J. Mandleberg, H. K, Rae, R. Hurst, G. Long. D. Davies and K. E. Francis, Z lnorg. NucL Chem. L 358 (1956). 27. L. B. Asprey and R. A. Penneman, J, Am. Chem. Soc. 89, 172 (1967). 28. R. C, Burns and T. A. O'Donnell, Unpublished observations. 29. R, A. Penneman, G. D. Sturgeon, L. B. Asprey and F. H. Kruse, J. Am. Chem. Soc. 87, 5803 (1965). 30. D. Brown, In Halides of the Lanthanides and Actinides. Wiley-Interscience, London (1968). 31. W. W, Sabol and N. J. Hawkins, Results cited in, LISAEC Rep. KAPL-1216 (1954) (Classified). 32. N. J. Hawkins, USAECRep. KAPL-I159 (19541. 33. R. D. Peacock and N. Edelstein, .L lnorg. Nucl. Chem ,38, 771 (1976). 34. R. D. Peacock and N. Edelstein, Lawrence Berkeley Labs., Nucl. Chem., Ann. Rep. 1974, LBL-4000 (1975). 35. The lattice energy of PuF4 may be calculated using two Born-Haber cycles and data for the lattice energy of PuO2 (M.F.C. Ladd and W, H. Lee, J~ Inorg. Nucl. Chem. 23, 199 (1961)), AH~ 02 (g) and AH~ F (g) (Y. C. Waddington, In Advances in Inorganic Chemistry and Radiochemis,to, (Edited by H. J. Emel6us and A. G. Sharpe), Vol. 1, pp. 157-221. Academic Press, New York (1959)) and AH~ Pu,O~ (s) and AH~ PuF4 (s) (J. Fuger, In MTP International Review of Science, Lanthanides and Actinides, Inorganic Chemistry, Series 1 (Edited by K. W. Bagnall), Vol. 7, pp. 157-2]!0. Butterworths, London (1972)). 36. D. M. Adams, In Inorganic Solids. Wiley, London (1974l. 37. J. Moseley, Results quoted in Ref.[12] and in M. J. Steindler, USAEC Rep. ANL-6743 (1963). 38. L. J. Guggenberger and R. A. Jacobson, lnorg. ('hem. 7, 2257 (1968). 39. R. C. Burns and T. A. O'Donnell, lnorg. Nucl. Chem. t,ett. 13, 657 (1977). 40. H. Selig, W. A. Sunder, F. A. DiSalvo and W. E. Falconer. J. Fluorine Chem. 11, 39 (1978).
1238
R.C. BURNS et al.
41. R. C. Burns, T. A. O'Donnell and A. B. Waugh, J. Fluorine Chem. 12, 505 (1978). 42. B. Weinstock, J. G. Maim and E. E. Weaver, J. Am. Chem. Soc. 83, 4310 (1961). 43. R. Rietz, T. A. O'Donnell and S. Yeh, Unpublished observations. 44. M. Baluka, S. Yeh, R. Banks and N. Edelstein, Inorg. Nacl. Chem. Lett. 16, 75 (1980). 45. L. E. Trevorrow, In Proc. Rocky Flats Fluoride Volatility Conf. (Edited by J. M. Cleveland and M. A. Thompson), USAEC Rep. CONF-680610, pp. 140-160 (1968). 46. L. E. Trevorrow, T. J. Gerding and M. J. Steindler, J. lnorg. Nucl. Chem. 30, 2671 (1968). 47. B. Weinstock and J. G. Maim, J. Inorg. Nucl. Chem. 2, 380 (1956). 48. L. E. Trevorrow, J. Fischer and W. H. Gunther, Inorg. Chem. 2, 1281 (1963).
49. So far, all attempts to prepare AmF6 (see, for example, Ref.[47] and S. Tsujimura, D. Cohen, C. L. Chernick and B. Weinstock, J. Inorg. Nucl. Chem. 25, 226 (1963)) have failed, but this may be due to radiation decomposition caused by the intense a-emission of 241Am.Use of the longer-lived isotope 243Am may allow AmF6 to be prepared. It should be noted that a similar effect is actually observed for CmF4, which can only be prepared with 244Cm and not with the shorter-lived isotope 242Cm(K. W. Bagnall, In The Actinide Elements. Elsevier, Amsterdam (1972)). 50. There is evidence to suggest that No(II) is considerably more stable than No(Ill) in aqueous solution, only being oxidized to the latter with difficulty (see R. J. Silva, T. Sikkeland, M. Nurmia, A. Ghiorso and E. K. Hulet, J. Inorg. Nucl. Chem. 31, 3405 (1969). 51. J. C. Eisenstein, J. Chem. Phys. 25, 142 (1956). 52. C. A. Coulson and G. R. Lester, J, Chem. Soc. 3650 0956).
0022-1902181/061231~98502.00/0 Copyright © 1981 Pergamon Pres!, ltd
REACTIVITY OF TRANSITION METAL FLUOR1DES--XII PLUTONIUM HEXAFLUORIDEt R. C. BURNS and T. A. O'DONNELL* Department of InorganicChemistry,Universityof Melbourne. Parkville,Victoria 3052, Australia and C. H. RANDALL AustralianAtomic Energy CommissionResearch Establishment,Lucas Heights, N.S.W. 2232, Australia
(Received 12 September 1980; receivedfor publication 16 October I980) Abstract--In a continuing study of the relative chemical reactivities of the higher fluorides of the d- and f-transition elements, oxidation-reductionreactions of plutonium hexafluoridewith a selected series of non-metal fluorides in low oxidationstates have been examined.This investigationhas also includedreactions with hydrogen and with various covalent chloridesand bromides.The results of this study have been correlated with those from previous investigations,and the reactivity of plutonium hexafluorideis considered in the light of plutonium as an f-transition element. Specialtechniques were developedfor the study of compoundsof this radioactiveelement,and a new procedurefor the preparationof plutoniumhexafluoridehas been devisedwhich requiresapparatus very much simpler than previouslyreported. INTRODUCTION Many studies of the relative chemical reactivities of the higher fluorides of the d-transition metals have established that the oxidant strength of a particular fluoride increases with increasing atomic number across any row of the Periodic Classification, but decreases with increasing atomic number in any group [1-3]. In addition, it has also been shown that significant differences develop progressively between successive pairs of second and third row elements moving from zirconium and hafnium across the transition metals series. The above trends, as established by chemical studies, are in agreement with trends in thermodynamic and other physical properties of the compounds [1-4]. Historically, the early members of the actinides were thought to form the beginnings of a new d-transition metal row because of the similarities of the formulae of the known compounds and many of their chemical properties to those of their respective d-transition metal congeners. Thus, for example, uranium was originally regarded as a member of sub-group VI. However, after the preparation of the post-uranium elements and publication of Seaborg's [5, 6] actinide hypothesis this view was not widely held. On the basis of the available physical and chemical evidence, the latter being very limited at the time, Seaborg argued that for actinium fZ = 89) and thorium (Z = 90) the binding energies of the 6d electrons were lower than for the 5f electrons but that at about protactinium ( Z = 91) a crossover occurs, so that for uranium (Z = 92) and the post-uranium elements the order is reversed leading to the progressive development of f-transition metal character in the actinide series. However, there are major differences in chemical behaviour between the earlier actinides and the formally corresponding lanthanides. In the lanthanides themselves the corresponding crossover (Sd vs 4f orbital energies) occurs at the beginning of the lanthanide series, *Author to whom correspondenceshould be addressed. +Part XI: J. lnorg. Nucl. Chem. 42. 1613 (1980). 1231
and more abruptly than is postulated for the actinides. The 4f orbitals do not extend very far from the nuclei by comparison with the s, p and d orbitals of higher principal quantum number. They are very effectiw~ly screened by these outer orbitals and take little or no part in the bonding of compounds of these elements. It is postulated that, for the actinides, relative spatial exte,nsion of the 5f orbitals is much greater than for the 4f and that 5f orbitals are available for compound formation. In earlier studies in this series selected oxidationreduction and halogen-exchange reactions of UF6 and PaF5 were studied[7,8], and comparisons with the formally analogous higher fluorides of the elements of subgroups VI (Cr, Me and W) and V (V, Nb and Ta), respectively, were examined[9,10]. For UF6 and the higher fluorides of sub-group VI the order of oxidant strength was found to be CrF6 ~ UF6 -> MoF6 > WF6 and it was argued that if uranium were a typical d-transition element the order should be CrF6 ~>MoF6 > WF~ > [-JF6. It was proposed that the experimental order of reactivity showed differences in bonding between UF6 and the other hexafluorides, providing chemical evidence for participation of 5f orbitals in the covalent bonding in UF6 and for the classification of uranium as an f-transition element. Similarly, for PaF5 and the higher fluorides of sub-group V, the experimental evidence indicates that while protactinium behaves somewhat like a d-transition element, the observed differences indicate that this element, like uranium, forms part of a series with properties different from both the d-transition elements and the lanthanides. In order to study further the effects of a d- to fcrossover and the emergence of f-transition metal character in the post-uranium elements, within the con-. text of studies of the type previously undertaken, it was decided to compare the chemical reactivity of PuF6 with that of the actinide fluorides UF6 and PaF5 and with OsF,, and RuF6, the formally analogous d-transition metal congeners of PuF6, which were also recently examined [2, 3]. In most of the earlier work on the study of tlhe relative chemical reactivities of the higher metal
1232
R. C. BURNS et al.
fluorides conducted at Melbourne, reactions were studied using fairly large quantities of fluorides (0.5-1g), employing apparatus and procedures generally following those described by Canterford and O'Donnell[ll]. However, in this study all work was carried out in glove-boxes because of the radioactive nature of the compounds and, as a safety requirement, reactions were studied using only about 50mg of hexafluoride. Thus special apparatus and procedures were developed for use in this investigation and, moreover, a simple method for the preparation of PuF6 was devised. The reliability of the procedures and apparatus was checked prior to the current study by investigating the reactivity of RuF6 which, at the time, was expected to be fairly similar to PuF6 in its chemical reactions with those selected reagents used in these investigations. When the technological importance of PuF6 in nuclear fuel processing and related contexts is considered, it is surprising that so little of the chemistry of this hexafluoride has been reported. Although its decomposition reaction and its interaction with alkali and alkaline earth metal fluorides are well-documented, few other reactions have been studied in any systematic way. A summary of reported results can be found in reviews by Steindler [12] and Vanka [13], and also as discussed by Cleveland [14]. EXPERIMENTAL
Apparatus and materials. All work was performed at the laboratories of the Australian Atomic Energy Research Establishment, Lucas Heights, N.S.W., on a vacuum line constructed from monel, stainless steel, and KeI-F. In general, the apparatus was similar to that described by Canterford and O'Donnell[ll], but had been adapted for use under glove-box conditions, as indicated above. Exentsive use was made of Whitey SS-1KS4, SS-IVS4, SS-14DKM4-S4 valves and valves constructed from Kel-F[15], while Whitey Sample Cylinders were used for the storage and handling of fluorine, and in the preparation of PuF6. Reactions were carried out in small, specially designed KeI-F reaction tubes (of capacity 3 cm3),which could, be fitted with an expansion bulb (ca. 25 cm3) if a high pressure was anticipated as a result of reaction, or for handling gases such as hydrogen as reactants. Preparation of plutonium hexafluoride. The apparatus used for preparation of this thermally unstable hexafluoride has been described previously[3] where it was used successfully for the preparation of RuF6. In principle, the method of preparation involves immediate cooling of the hexafluoride upon formation as in earlier procedures which relied on very complex, costly and sophisticated equipment but, unlike those earlier methods, requires only simple apparatus based on a stainless steel cylinder, the base of which can be heated and which is provided with an external cooling trough which can be maintained at -78°C fitted to the upper part of the cylinder (for details see Ref. [3]). In this work PuO2 (ca. 0.8g, supplied by the AAEC; >96% 239pu) was first treated in the reaction cylinder with excess hydrogen fluoride at 450-550°C for 2 hr to produce a mixture of lower fluorides of plutonium (predominantly PuF3), water and oxygen. The excess hydrogen fluoride and other products were then removed by evacuation. This procedure was repeated once. Fluorine, slightly in excess of the amount required to form 50 mg of PuF6, was admitted to the cylinder and PuF6 was prepared by heating the lower fluoride(s) and fluorine at 500-550°C for 1 hr. Upon formation, the PuF6 condensed in the cooled (-78°C) region of the reactor and, after cooling of the cylinder and removal of the excess fluorine, the coolant was removed from the external trough and the hexafluoride used directly from the vessel. Waste hydrogen fluoride and fluorine were passed through a granulated alumina tower, to give water and oxygen, prior to disposal. Other reagents. Anhydrous hydrogen fluoride (AHF) Matheson, 99.8% min) was distilled on a column similar to that
described by Shamir and Netzer[16], while C1F3 (Matheson, 98% rain) was doubly distilled in a vacuum before use. The perfluoroalkanes, n-C6Fi4 and n-CTFI6 (P.C.R.), were stored over P4Oto and were distilled directly before use. All other reagents were prepared and/or purified as described previously[9], or were of reagent grade purity. Reaction procedures. All volatile reagents were transferred by distillation in a vacuum, while solid materials were handled in a glove bag filled with dry argon. Reactions were carried out by condensing excess reductant, if volatile, onto the hexafluoride at -196°C, whereas for solid reductants the hexafluoride was condensed directly onto excess reagent at -196°C. In either case the reactants were then allowed to warm slowly to room temperature during which time any reaction was noted. Previous studies involving OsF6, IrF6 and RuF6 have shown that some reactions of these hexafluorides are best studied under controlled conditions at low temperatures, using an inert solvent as a "thermal sink';, because of their extremely exothermic nature[2, 3]. Now PuF6 is thermally unstable at room temperature and so would be expected to be extremely reactive[l]. Therefore the reaction of PuF6 with PF3 was conducted in liquid AHF and the majority of the reactions of PuF6 with CCI4, BCI3 and BBr3 were conducted in an n-perfluoroalkane (either n-C6FI4 or n-CTFI6) solvent. The procedure employed for this type of reaction has been described previously [21. Identification o/products. Upon completion of a reaction the volatile species were removed and, after fractionation if necessary, were identified by their physical properties and IR spectra. Gas-phase IR spectra were recorded on either a Perkin-Elmer 225 grating IR spectrophotometer or a Perkin-Elmer Model 21 double beam IR spectrophotometer, the samples being contained in a monel-bodied cell fitted with detachable AgC1 windows. Solid products of reaction were identified by X-ray powder photography and by their physical properties. Samples for X-ray analysis were loaded into ca. 0.4 mm o.d. Pyrex capillaries under anhydrous conditions and photographs were taken using a Philips Debye-Scherrer camera of diameter 11.46cm, on a Philips Industries Ltd. X-ray generator, Model PW 1008/30. Nickel-filtered Cu-K~ radiation (Ks = 1.5418,~) was employed for all samples and exposures required about 12 hr. RESULTS AND DISCUSSION
A summary of results of the reactions of PuF6 that were studied in this investigation is given in Table 1. As observed in previous studies of other hexafluorides, the reactions of PuF6 with the lower fluorides of the nonmetals, when they occurred, involved simple oxidationreduction, with oxidation of the lower fluoride to the corresponding higher fluoride and reduction of the hexafluoride to oxidation state IV. Also, in some reactions evidence was found for adduct formation between the products of reaction and this is discussed in detail below. Other reactions of PuF6, for example those with hydrogen and the covalent chlorides and bromides, also involved oxidation-reduction and, where possible, comparisons are made with the related uranium and neptunium systems.
Oxidation-reduction reactions o[ plutonium hexafluoride In anhydrous HF, PuF6 was found to oxidize PF3 to PF5 well below room temperature, with reduction of the hexafluoride to PuF4 as shown by X-ray powder photography. Arsenic trifluoride also reacted with PuF6 below room temperature, but in this case no free AsF5 could be detected as a gaseous product. Furthermore, the solid product of reaction was pale orange in colour, whereas PuF4 is light brown to pale pink depending on the method of preparation. X-ray examination of the solid product did show, however, a rather weak and diffuse pattern of PuF4. The solid product exhibited no great solubility in anhydrous HF and showed no evidence of
Reactivityof transition metal fluorides--XII Fable I. Oxidation-reductionreactions of plutonium hexafluoride Reactant
Products
I>F3:AHF (a)
P u F . , PF~
AsI:,
PuFf. PuF..AsF~ adduct
Si~F
PuF4, SbF5
BiF,
No reaction
C!F~
No reaction
R~F.
No reaction
!t2/AEF
Pu(IV),
Ci~I,,,'PFA
PuF., CF3CI, CF2C12, C12
BI?!~/PFA
PuF.,
BBr~/PFA
PuFf, BF3 (c) , Br2
gBr,
PuFf, BF3, Br2
Pu(V)
(b)
BF3 (c) , C12
(a)
Key:
AHF = anhydrous hydrogen fluoride;
(b)
For a dlscussJon of this reaction,
~c>
No mixed halides of boron were detected.
PFA = n-perfluoroalkane
(n-C6Fx~ or n-CTF~6). see text.
liberating AsFs on heating at 90°C for l hr. A similar result was also obtained when the reaction was repeated using approximately stoichiometric quantities of reactants. Interestingly, it should also be noted that the reaction between UF6 and AsF3, although it does not proceed readily, gives UF4 as one product, and, like the above reaction, no evidence of free AsF5 as a gaseous product [7, 17]. Recently, it has been shown that UF4 may be solubilized to some extent in anhydrous HF by reaction with AsF5 over many days[18]. After decantation and removal of the solvent a solid product of empirical composition approximating to UF4'2AsF5 remains and it is suggested that the residue might be UF2(AsF,)2[18]. It is assumed here that an involatile adduct is formed between PuF4 and AsFs, but its composition is unknown. The reaction of PuF6 and SbF3 occurred just below room temperature to give PuF4 and SbFs. The formation of both products was inferred by identification of the salts /3:K2PuF6 (isomorphous with /3,-K2UF6 and ~,K2NpF6) and KSbF6 by X-ray powder photography, after conversion of the reaction products to the above salts by treatment with potassium fluoride in anhydrous H F In view of the probable complexity of the actual reaction mixtures, no attempt was made to investigate if there was any adduct formation between the products of reaction and excess SbF3, and/or between the products themselves, although the latter seems highly probable, following the AsF3 reaction. Ihe two remaining lower fluorides, BiF3 and C1F3, showed no reaction with PuF6 on standing for several hours. In the latter case this is not surprising as CIF3 may be used to prepare PuF6 by fluorination of PuF4 at high temperatures[12, 14]. Also, no reaction was observed to occur between PuF~ and the highest thermally stable fluoride of ruthenium, RuFs. In previous investigations in this series the reactions of a particular fluoride with the reagents SF4, SeF4 and CS2 have also been used as a guide to relative chemical
1233
reactivity. However, it should be noted that prior to t]his study PuF6 has been shown to oxidize SF4 to SF6, with reduction of the hexafluoride to PuF4 [19-21], while CS~, and PuF6 react to give PuF4, with CF4, SF4 and SF6 as gaseous products[22]. As SeF4 has been found to be comparable to SF4, but stronger than SbF3, in reducing ability towards related hexafluorides, reaction with this potential reductant was not investigated in this study. The reaction between PuF6 and hydrogen has previously been studied by Hawkins[23], who observed that reaction proceeded slowly at room temperature to give HF and a solid which was presumed to be PuF4, perhaps with the intermediate formation of the uncharacterized pentafluoride. In two related studies, Levy and Wilson[24] and Asprey and Paine[25], found that UF6 and hydrogen react in anhydrous HF to give /3-UF5 but, in the presence of platinum foil or UVradiation, UF4 is formed. In the present study the reaction between PuF6 and hydrogen was investigated in anhydrous HF. Reaction between PuF6 and an excess of hydrogen occurred slowly at room temperature to give an insoluble white to pale pink solid, the reaction being allowed to continue until all of the PuF6 had been consumed. The solid reaction product was somewhat sensitive to moisture, and gave some Pu(VI) (as PuO:~~+) on addition to weakly acidic or neutral solution, in addition to insoluble PuF4. This suggests that some Pu(V) formed on reduction which, like UFs, disproportionates into the dioxo-cation (VI) and the insoluble tetrafluoride in aqueous solution in the presence of F ion, a strong complexing agent. Unfortunately, the solid was amorphous to X-rays. The reaction was repeated, with identical results, and on addition of a solution of excess rubidium fluoride in anhydrous HF to a sample of the reaction product a light-pink solid was isolated. This material turned brown on exposure to moisture, with evidence for some Pu(VI) on dissolution in aqueous solution, which is uncharacteristic of any known rubidium (or other alkali metal) complex fluoride of Pu(IV). In another reaction, using approximately stoichiometric quantities of reagents, the reaction was terminated with a little PuF~ still present, and a small amount of an involatile crimson-red material was also observed, which exhibited properties similar 1:o those described above. The formation of a brick-red material has previously been observed during fluorination of PuF4 and was tentatively identified as Pu4F~v by chemical analysis and X-ray powder photography[26]. Although no conclusive evidence for PuF~ was obtained in this study it would appear that some plutonium is in oxidation state intermediate between VI and IV on reduction of PuF6 by hydrogen, although Pu(IV) appears to be the primary product, particularly in the presence of excess reductant. Moreover, it is suggested that this reaction might be strongly controlled by kinetic factors such that any Pu(V) that is formed may be reduced at a comparaNe rate to that of PuF6. Thus the reaction could proceed through the intermediate formation of mixed oxidation state products, analogous to U2F9and U , F : . ultimately leading to PuF4. The formation of Pu(V), either as PuF~ or as a complex fluoride, was also attempted by fluorine oxidation of PuF4 in anhydrous HF or in an RbF. 2HF melt at 90°C over several hours. However, neither reaction proved successful. It should be noted that under similar conditions UF4 in anhydrous HF is oxidized to /3-UF~ at room temperature, with slow formation of UF~ on
1234
R. C. BURNS et al.
standing, and forms salts oI composition MUF6 (M = Na, Rb, Cs) in the respective melts [27, 28]. In contrast, NpF4 is not oxidized in anhydrous HF at 25°C but does give CsNpF6 in a CsF.2HF melt at 70°C[27]. These results reflect the trend toward progressively decreasing stability, and hence difficulty of attaining, high oxidation states across the actinide row from uranium onwards. Thus complex fluorides of Pu(V), containing rubidium and cesium, have only been prepared by direct fluorination of the respective Pu(IV) salts at 300--400°C for several hours [29].
Halogen-exchange reactions of plutonium hexafluoride In their extensive survey of the reactions of UF6 with various covalent chlorides (and bromides), O'Donnell et al.[7] found that the formation of different products could be rationalized on the basis of a competition between two reactions, depending on their kinetic involvement, viz. simple halogen exchange of UF6 to give UC16 and the appropriate exchange products, and the secondary reaction of UF6 and UCI6 to give UF4 and chlorine (for which AG25oc= -29 kcal/mol of UF6). For excess of the halogen-exchange reagent the final product obtained in each reaction was found to depend on the comparative rates of the two reactions, such that when halogen exchange was fast, e.g. with BCI3 and AIC13, only UC16 was produced but, when halogen exchange was slow, e.g. with CC14 and SiCI4, the UF6 was effectively in excess and reacted with the UC16 as it formed to give UF4. For TiCI4 both products were observed so that the two reactions have comparable rates in this case. For excess of UF6, only UF4 was observed as a reduction product, in accord with the above arguments. On reaction of UF6 with the covalent bromides BBr3 and PBr5 only UF4 was observed, and it was concluded that in these reactions UF6 was reduced directly to UF4, reflecting the thermodynamic instability of the higher uranium bromides [30]. In this analogous study the reactions of PuF6 with the covalent chlorides, BCI3 and CCI4, reagents which had shown fast and slow reactions with UF6, respectively, and with the covalent bromide, BBr3 were investigated. The reaction of PuF6 and CC14 in the gas phase has previously been studied by Sabol and Hawkins[31,32], who observed that reaction proceeded rapidly at room temperature to form CFCI3 and CF2Clz, while PuF6 and CFCI3 react to give CF2C12 and CF3C1. For neither reaction was the reduction product reported, but was presumed to be PuF4 in each case. In this study, reaction between PuF6 and CC14 was observed just below room temperature in n-perfluoroalkane to give CF2C12, CF3CI and chlorine as gaseous products, while the solid product was identified as PuF4 by X-ray powder photography. Reaction of PuF6 with either BCI3 or BBr3 in nperfluoroalkane occurred well below room temperature to give PuF4 and BF3 in each reaction, and chlorine or bromine, respectively. Also, reaction of PuF6 and BBr3 in the absence of n-perfluoroalkane gave a violet-blue solid which, based on its colour, was presumably PuF3. It would appear, therefore, that all of the reactions of PuF6 with the covalent chlorides and bromides under conditions where the heat of reaction can be dissipated effectively involve simple oxidation-reduction, such as, 3PuF6 + 2BX3 P F A > 3PuF4+2BF3+3X2; (X = CI, Br).
Under less controlled conditions, a typical reaction is PuF6 + BBr3 ~
3 PuF3 + BF3 + ~ Br2.
The highest thermodynamically stable solid chlorides and bromides of plutonium are the trihalides, although PuC14 has been observed in the gas phase at high temperatures[14], while both tetrahalides are found at room temperature in solid complexes with various amides, dimethyl sulphoxide or triphenylphosphine oxide [14]. Hence, the formation of lower fluorides in the above reactions, probably by direct reduction of the hexafluoride, is consistent with the thermodynamic instability of the "higher" chlorides and bromides of plutonium. It should be noted that Peacock and Edelstein[33, 34] have reported that NpF6 and BC13react to give NpF4, BF3 and chlorine, so that an analogous argument is applicable to the neptunium system as well. In the reactions of PuF6 with the reagents listed above, and others[12], it should be noted that the hexafluoride is reduced almost exclusively to the tetrafluoride. The formation of a solid product, like PuF4, in such a wide variety of reactions may be attributed to either or both of the following factors; a very stable oxidation state or, for a compound which is predominantly ionic in character, a very high lattice energy. For plutonium, it would appear that a combination of both factors allows the product PuF4 to dominate the reduction reactions of the hexafluoride. In general, Pu(IV) is the most stable oxidation state (see also below) in aqueous solution, although because of the various equilibria and kinetics of conversion from one oxidation state to another, plutonium can form appreciable concentrations of the IIl, IV, V and VI oxidation states even in the same solution, depending on the reaction conditions[14]. Plutonium tetrafluoride, like all of the actinide tetrafluorides, is also highly ionic in character as demonstrated by its melting point of 1037°C[30], whereas covalent tetrafluorides generally boil below ca. 300°C. Thus the lattice energy of PuF4 is very high, with a calculated value of ca. -2060kcal/mol[35], and is appreciably larger than the lattice energies of, for example, NaF (-216 kcal/mol) or MgF2 (-695 kcal/mol)[36]. Hence PuF4 is expected to show considerable stability and, consequently, is the preferred reduction product of PuF6. However, it should also be noted that under more vigorous reaction conditions, reduction of PuF6 to PuF3 can occur as shown in its reactions with BBr~ and with iodine [37]. In the related uranium and neptunium systems oxidation states greater than IV are relatively more stable than found in the plutonium system--the most stable oxidation states in aqueous solution are U(VI) and Np(V), respectively--so that similarities in the reaction patterns of the hexafluorides and other fluorides of these elements are to a large degree a reflection of the very high lattice energies of the tetrafluorides.
Comparative reactivity o[ plutonium hexafluoride and related hexafluorides As indicated in the Introduction to this paper previous studies in this series have compared the chemical reactivities of the higher actinide fluorides, PaF5 and UF6, with those of their formally analogous d-transition metal counterparts[7-10]. These studies demonstrated that PaF5 was a considerably stronger oxidant than NbF5 or TaFs, which exhibit similar chemical reactivities them-
Reactivity of transition metal fluorides--XII selves, but was a much weaker oxidant than its first row transition metal congener, VFs, Similarly, UF6 was shown to be a marginally stronger oxidant than MoF6 which was, in turn, somewhat stronger than WF6; however, like PaFs, UF6 was a very much weaker oxidant than its first row transition metal counterpart, in this case CrF~. Hence. when protactinium and uranium are considered in their historical positions as members of sub-groups V and VI, i.e. as d-transition elements, the chemical reactivities of PaF~ and UF6 are anomalous as there should be a marked decrease in reactivity of the higher fluorides with increase in atomic number and they should not be classified as typical d-transition metals. Consideration of the thermodynamic stabilities of the binarv fluorides of these elements also supports the above argument. Thus, while vanadium forms fluorides from the highly reactive VF~ to VF2, and niobium forms NbF,, NbF4 and the intermediate fluoride, NbF2.5--the two latter fluorides being produced only by high temperature reduction of the pentafluoride by the metal-there is no reliable evidence for the formation of any solid fluoride lower than TaFs. The very existence of PaF4 in addition to PaFs indicates that protactinium does not follow the trend expected for the d-transition metals, that is, an increase in the relative stability of the highest oxidation state down the group. Similarly, in sub-gruop V1, although chromium forms the unstable CrF6, simple binary or intermediate fluorides from CrF5 to CrF~ may be easily prepared, while molybdenum and tungsten form stable hexafluorides but less stable lower fluorides down to MoF, and WF4,respectively. Uranium, on the other hand, forms stable fluorides, including intermediate fluorides, ranging from UF6 to UF3. the relative chemical reactivities and thermodynamic stabilities of the higher fluorides therefore show that neither protactinium nor uranium behaves chemically as a d-transition element. This, it has been suggested, provides chemical evidence for differences in bonding between the respective actinide and d-transition metal fluorides and, consequently, within the context of Seaborg's actinide hypothesis[5,6], it has been proposed that 5f orbitals are available for bonding in these compounds. Hence for example, d2sf 3 hybrid orbitals could participate with d2sp 3 hybrids in the bonding of UF,[71. Importantly, just as the reactivity and thermodynamic stability of the higher fluorides of protactinium and uranium do not fit into the pattern shown by the d-transition elements, from a chemical viewpoint these elements differ markedly from their formal lanthanide counterparts. Thus, the chemistry of both praseodymium and neodymium is dominated by oxidation state llI. although compounds of Pr(IV) can be prepared with some difficulty. Before discussing the chemical reactivity of PuF6 in relation to its d-transition metal congeners and to the olher higher fluorides of the actinides, it is expedient to iook at various chemical and thermodynamic properties of the higher fluorides of the elements of sub-groups 11I, IV and VII. The chemical and physical properties of AcF3 closely resemble those of ScF3, YF3 and LaF3. Indeed, only differences in the appropriate cationic radii are required to account for variations in the properties of these compounds, and each element is assigned a ground-state electronic configuration of (n- 1) d~ns 2. For thorium and the elements of sub-group IV, therIN(
\tl] a l Nt! i* ]
1235
modynamic stability considerations indicate that, whereas tetrafluorides and lower fluorides are known for titanium and zirconium, hafnium and thorium form HfF4 and ThF4 only. Useful comparisons may also be extended to the other halides and here, it is noted that titanium, zirconium and hafnium form various simple and intermediate halides ranging from tetrahalides to monohalides, the lower halides being relatively easily prepared for titanium and zirconium, but tess so for hafnium. Thorium, on the other hand, forms only tetrahalides and, under very strongly reducing conditions, a di-iodide and, less-well characterized, a tri-iodide. However, these are not regarded as simple binary compounds but are metallic compounds with a formulation of Th(IV)I2e2 in the case of the di-iodide [38]. It is therefore apparent that oxidation state IV increases in stabilit} from titanium to thorium, and this is in accord with the assignment of a ground-state electronic configuration of (n-l)d2ns 2 for these elements. Moreover. thorium differs in a significant way from the lanthanide cerium, which has the accepted configuration 4f26s 2, and which exhibits oxidation states of III and IV in both the solid state and in solution. There is no reported solution chemistry of Th(III). At this point in the discussion we see in the actinide row typical d-transition metal behaviour for actinium and thorium but, as described above, for protactinium and, to a greater extent, for uranium, we begin to see in the series the emergence of properties different fi'om both the d-transition elements and the lanthanides. It would therefore appear that protactinium is just about at the energy crossover point of the 6d and 5f electrons as postulated by Seaborg. Proceeding to neptunium and the sub-group VII ,elements, even greater differences between this actinide ;and its formal d-transition metal congeners may be noted. For the d-transition elements increasing stability of oxidation state VII with increasing atomic number results in the existence of the stable binary fluoride. ReFT. If neptunium were considered a d-transition element then a stable heptafluoride should be formed even more readily than for rhenium. In fact, the highest binary fluoride of neptunium is NpF6, and easily attainable stable fluorides are formed down to the trifluoride. Furthermore, attempts to prepare even an oxide fluoride of Np(VII) have so far failed [33]. For plutonium and its d-transition metal congeners the differences are further magnified. Thus, there are no binary or oxide fluorides of plutonium that correspond to the oxidation state VIII and VII osmium species, OsO3F2, OsO2F3, OsOF5 and the rather unstable ()sF:,. Plutonium itself forms various binary, intermediate and oxide fluorides including the thermally unstable PuFf,. PuO2F,~ and the recently prepared PuOF4139]. and the very stable lower fluorides, PuF4 and PuF3. There is also evidence that an unstable and highly reactive pentafluoride exists, for which some supportive evidence was also obtained in this study. From a chemical point of view, consideration of the relative reactivities of the hexafluorides demonstrates the marked difference between plutonium and the trend established by its d-transition metal congeners. For the d-transition metal fluorides it has been shown that, while RuF6 and OsF6 oxidize PF3 (in anhydrous HF) and AsF, to their respective pentafluorides, RuF~ is reduced to Ru(IV) in both reactions but OsF6 gives Os(IV) and Os(V), respectively[2,3/. Furthermore, RuB, oxidizes
1236
R. C. BURNS et al.
SbF3 to SbF~, being reduced in turn to Ru(V), while
OsF6 does not react with this lower fluoride [2]. Osmium hexafluoride is, therefore, a considerably weaker oxidant than RuF6. Complementary evidence for this trend is provided by a comparison of the reactions of the hexafluorides with water in HF[40] and with boric oxide in an inert solvent (n-C6F~4)[41]. Ruthenium hexafluoride oxidizes both reagents, with reduction to Ru(V), while OsF6 undergoes simple oxygen-fluorine exchange to give OsOF4, with no change in the formal oxidation state of osmium. In comparison with the trend established above, PuF6 oxidizes PF3 (in anhydrous HF), AsF3 and SbF3 to their respective pentafluorides, with reduction of the hexafluoride to Pu(IV) in each case. However, PuF6 shows no reaction with BiF3 or C1F3, unlike RuF6 which reacts with both of these lower fluorides [3]. Based on its thermal instability, RuF6 should also be capable of oxidizing PuF4 to PuF6, as does PtF6142], although this reaction was not investigated in the present study. If this reaction does proceed then, conversely, PuF6 should not be able to oxidize RuF5 to RuF6, as shown in this work. Therefore, from a comparison of all of the above reactions, it is apparent that the order of oxidizing strength of the hexafluorides is RuF6 > PuF6 > OsF6, implying that PuF6 is considerably more reactive than would be expected if plutonium were regarded as a d-transition element. It could be argued, therefore, as will be discussed below, that 5f orbitals make some contribution to the covalent bonding in PuF6, just as previously suggested for PaF5 and UF6 on the basis of similar studies in this series[7, 8]. As is evident from the foregoing discussion, the thermodynamic stabilities of the higher fluorides decrease across the actinide row, and this is reflected in a corresponding increase in chemical reactivity. For example, PaF5 slowly oxidizes PF3 to PFs, being itself reduced to PaF4, but shows no reaction with AsF3. On the other hand, UF6 in a Kel-F reaction tube where heat dissipation is poor reacts readily with PF3 tO give PF5 and UF4. Reduction of UF6 in a metal reactor at -78°C produces/3-UF5143]. It is assumed that, under the thermal reaction conditions in a KeI-F reactor which is itself at room temperature, UF5 disproportionates to UF4 and UF6 which is then reduced by PF~. Uranium hexafluoride reacts to some extent with AsF3, but fails to oxidize either SbF3 or SF4 (also see [21]). Few reactions have been reported for NpF6, but it is known that this hexafluoride oxidizes PF3 tO PFs, with the formation of a-NpF5 under the conditions given above for preparation of/3-UF5144]. It reacts readily with bromine below room temperature to give NpF4 and BrF3145]. In contrast UF6 shows no reaction with bromine, or with BrF3, under similar conditions [30]. In fact, both NpF6 and OF6 may be prepared by oxidation of the respective tetrafluorides with BrF3 and BrF5 at high temperatures[45, 46]. As demonstrated in this study PuF6 oxidizes PF3 in anhydrous HF to PF~ and is reduced directly to the tetrafluoride. Plutonium hexafluoride also oxidizes AsF3, SbF3 and SF4 to their respective higher fluorides, with reduction to Pu(IV) in each case. Furthermore, PuF6 oxidizes both bromine and BrF3 to BrF5 with reduction of the hexafluoride to PuF4[12, 47]. Previously, reaction of metal and non-metal fluorides with CS2 has been ~own to provide a good indication of relative chemical reactivity as different products are formed depending on the oxidizing ability of the parti-
cular fluoride[I-3,48]. For the actinide hexafluorides it has been found that UF6 reacts with CSz at 25°C in the manner characteristic of a moderately strong oxidant, that is, by cleavage of one C-S link, to give UF4, SF4, (CF3)2Sz and (CF3)2S3, although at higher temperatures reaction is somewhat more vigorous and SF6 and CF4 are also formed. In contrast, PuF6 reacts with CS2 at room temperature to give PuF4, CF4, SF4 and SF6, characteristic of an extremely powerful oxidant; the reaction in this case involving the cleavage of all C-S bonds [22, 48]. Unequivocal evidence that PuF6 is a stronger oxidant than UF6 is shown by the fact that PuF6 completely oxidizes UF4 to UF6 above 200°C, with reduction of the hexafluoride to PuF4147]. For the higher fluorides of the actinides oxidant strength follows the order PuF6 > NpF6 > UF6 > PaFs, although comparison with the latter is less exact as it is involatile and the possibility of an homogeneous reaction is ruled out. The trend towards increasing chemical reactivity and the related decrease in thermal stability across the actinide row is further evidenced by the thermal instability of PuF6. For the post-plutonium elements the thermal stability of the higher fluorides decreases continuously across the series. No hexafluoride of americium has yet been prepared (but see [49]) and, although CmF4, BkF4 and CfF4 are known, from curium onwards the most stable oxidation state of the actinides is oxidation state III, As found in the lanthanides [50]. From studies on the relative chemical reactivities of the hexafluorides of plutonium, ruthenium and osmium, as discussed above, it has been shown that PuF6 is substantially more reactive than would be expected if plutonium were considered a d-transition element. These differences in chemical reactivity, which indicate that plutonium does not behave chemically as a d-transition metal, suggest differences in bonding between the hexafluoride of plutonium and its d-transition metal congeners. It is therefore proposed that 5f orbitals could make some contribution to the covalent bonding in PuF6, presumably involving hybrids such as d2sf 3, partially at least. Significantly, it should be noted that from a chemical viewpoint, plutonium does not resemble to any great degree samarium, its formal lanthanide counterpart, except, perhaps, for some aspects of oxidation state III. The chemistry of samarium is, like all of the lanthanides, dominated by this oxidation state, although divalent samarium can be stabilized in some compounds but is, in general, easily oxidized to Sm(III). It should also be noted that there is no evidence for the formation of divalent plutonium. Many authors have considered the possibility of 5[ orbital involvement in the bonding of the early actinides, and there is considerable physical and chemical evidence to support this suggestion. While it is beyond the scope of this paper to discuss such evidence, it is briefly noted that favourable overlap of hybrids which involve f orbitals has been theoretically demonstrated[51], while Coulson and Lester[52] have shown the 5[ orbital involvemen't in bonding in U(VI) species, such as UO22÷, is energetically favourable. Of considerable importance to the arguments described above is the thermodynamic stability and chemical reactivity of the higher fluorides and other halides of the d- and •-transition elements, which have been reviewed by O'Donnell[l]. Of course, it is to be appreciated that chemical studies such as those described above do not in themselves prove that 5f
Reactivity of transition metal tluorides--XlI orbitals are involved in bonding in these compounds. Rather, they serve to show that bonding in the higher actinide fluorides is very different from that in the dtransition metal fluorides where, for example, d2sp 3 hybridization is used to account for the bonding in the hexafluorides. One final point of interest concerns the relative ordering in chemical reactivity of the actinide higher fluorides with regard to their d-transition metal congeners across the Periodic Classification. Thus, while PaF5 is a considerably stronger oxidant than NbF5 or TaFs, UF6 is only slightly stronger than MoF6, which is a stronger oxidant than WF6. Little is known about the relative oxidizing abilities of NpF6 and TcF6, but they would appear to be of comparable reactivity and both are stronger oxidants than ReF~. Furthermore, neither neptunium nor technetium form heptafluorides, as does rhenium. However, in this study we have found that PuF6 is not a stronger oxidant than both of its dtransition metal congeners, but is, rather, a weaker oxidant than its second row congener, RuF6, but a stronger oxidant than its third row congener, OsF6. For any particular reaction, all things being equal, this presumably represents some disparity between the actinide and d-transition metal higher fluorides (including their reduction products) with regard to trends in the endothermic valence state promotion energy and the exothermic intrinsic bond energy (entropy terms are likely to be comparable in these systems, at least for the hexafluorides and their reactions) across the Periodic Classification, reflecting the closeness of the 7s, 7p, 6d and 5[ orbitals in the actinides. However, without data such as ionization potentials or, better, vale,ace state ionization energies as calculated from atomic spectra, together with other reliable physical data on both the actinides and d-transition metals, no quantitative comparisons can be made. For the post-plutonium elements the increasing inertness of the 5f orbitals, both energetically and because of their decreasing spatial extension as a result of poor shielding from the progressively increasing nuclear charge, militates against the formation of high oxidation state compounds so that we quickly see the progressive stabilization of oxidation state III in the remaining elements and even the onset of a chemistry of ~gxidation state 1I in the later members. Acknowledgements--We wish to acknowledge financial assistance from the Australian Atomic Energy Commission and the co-operation of the officers of the AAEC Research Establishment who provided the plutonium dioxide and made possible the glove-box manipulation of the radioactive materials. We also acknowledge receipt of an Australian Postgraduate Research Award (R.C.B.). REFERENCES
I. T. A. O'DonneIl, In Comprehensive Inorganic Chemistry (Edited by J. C. Bailar, H. J. Emel~us, R. S. Nyholm and A. F. Trotman-Dickenson), Vol. 2, pp. 1073-1106. Pergamon Press, Oxford (•973) and refs. therein. 2. R.C. Burns and T. A. O'Donnell, J. lnorg. Nucl. Chem. 42, 1285 (1980). 3, R.C. Burns and T. A. O'Donnell, J. lnorg. Nucl. Chem. 42, 1613 (1980). 4. N. Bartlett, Angew. Chem. Int. Edn. 7, 433 (1968). 5. G. T. Seaborg, In The Chemistry of the Transuranium Elements (Edited by G. T. Seaborg, J. J. Katz and W. M. Manning), Nat. Nucl. Energy Set., Vol. IV, 14B, pp. 14921524. McGraw-Hill, New York (1949). 6. G. F. Seaborg, Nucleonic,~ 5(5), 16 (1949).
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49. So far, all attempts to prepare AmF6 (see, for example, Ref.[47] and S. Tsujimura, D. Cohen, C. L. Chernick and B. Weinstock, J. Inorg. Nucl. Chem. 25, 226 (1963)) have failed, but this may be due to radiation decomposition caused by the intense a-emission of 241Am.Use of the longer-lived isotope 243Am may allow AmF6 to be prepared. It should be noted that a similar effect is actually observed for CmF4, which can only be prepared with 244Cm and not with the shorter-lived isotope 242Cm(K. W. Bagnall, In The Actinide Elements. Elsevier, Amsterdam (1972)). 50. There is evidence to suggest that No(II) is considerably more stable than No(Ill) in aqueous solution, only being oxidized to the latter with difficulty (see R. J. Silva, T. Sikkeland, M. Nurmia, A. Ghiorso and E. K. Hulet, J. Inorg. Nucl. Chem. 31, 3405 (1969). 51. J. C. Eisenstein, J. Chem. Phys. 25, 142 (1956). 52. C. A. Coulson and G. R. Lester, J, Chem. Soc. 3650 0956).