Binary Fluorides of Noble-Gases and Their Compounds Boris Žemva

Home , Fluorine, Hexafluoride, Krypton difluoride, Perbromate, Xenon difluoride, Xenon hexafluoride

CROATICA CHEMICA ACTA CCACAA 61 (1) 163-187 (1988)

YU ISSN 0011-1643 CCA-1789 UDC 546.29 Author's review Binary Fluorides of Noble-Gases and Their Compounds Boris Žemva

»Jožei Stejosi« Institute, »Edvard Kardelj« University of Ljubljana, 61000 Lj'j,bljana, Yugoslavia

Received August 26, 1987 The synthetic work done during the last 25 years in the field of binary fluorides of noble gases (especially xenon and krypton) and the ir compounds is crttically reviewed. Besides, a historical introduction to the noble gas chemistry, its challenges in the future and possible uses of the noble gas compounds are also given.

INTRODUCTION 1.1. General Introduction Since 1962, when the first xenon compound was reported", many books and reviews treating the developments in the noble-gas chemistry have been published. Up to now the most extensive one, dealing with all aspects of the noble-gas chemistry, was written by Bartlett and Sladky-. Recently, a review of cationic and anionic complexes of the noble gases was added by Selig and Holloway", The purpose of this review is to summarize the synthetic work done during the last two and a half decades in the field of bin ary fluorides of noble gases (especially xenon and krypton) and their compounds.

1.2. Historical Introduction Argon, the most plentiful of the group of noble gases, was isolated and announced as a new element in 1894 by Lord Rayleigh and Sir W. Ramsay. They tried various chemical reagents to form compounds of the new element but without success. Ramsay also sent a 10 ml sample of the gas to his friend H. Moissan, the discoverer of fluorine, so that he should attempt to prepare a fluoride. In 1895, Moissan reported" that the reaction between arg on and fluorine did not proceed. This experiment was considered the ultimate test for chemical inertness of argon. Some years later, the electronic theory of valenceš-" supported the proposed chemical inertness of noble gases. The noble-gas valence electron configuration was clearly defined as the configuration to which other elements tended in their chemical bonding.

In 1933, some new reports7,8,9,lO of unsuccessful experiments to make noble gas compounds were published simultaneously with some theoretical specu- lations that noble-gas compounds should existl1,12. Among these experiments the most promising experiment was made by Yost and Kaye on a suggestion 164 B. ZEMVA of Pauling, who was convinced that, at least, xenon and fluorine should react. This failure was really unfortunate since a similar experiment, carried out 32 years later but replacing the electric discharge of the early experiment with sunlight13,14 provided a convenient preparative method for xenon difluoride. There is no doubt that if xenon had been as abundant as argon we should not have had to wait more than 60 years for the noble-gas chemistry. The legend that the inert gases do not form stable compounds was a rather firmly established chemical dogma until the spring of 1962. At that time Bartlett and Lohmann discover ed an oxyfluoride of platinum which proved to be the salt 02+PtF6- 15. This formulation indicates that PtF6 is an oxidizer of unexpected power. Bartlett noted that the ionization potential of xenon atoms is almost identical to that of oxygen molecules (02 = 12.2 eV, Xe = 12.10 eV). He tried reaction of xenon gas with PtF6 vapours. Yellow-o range xenon fluoroplati- nate(V) was obtained--P. This and the subsequent discovery of xenon fluo- rides17,18,19 started a new era of inorganic chemistry.

2. XENON CHEMISTRY The normal oxidation states are even and range from + 2 to + 8. The oxidation state + 8 is known in the oxide Xe04, the perxenates Xe064- and xenon trioxide difluoride. XeFs is not known. The fluorides are readily preparable from the elements and are thermodynamically stable. Among the other binary compounds with the halogens only XeCl2 and XeBr2 are known.

2.1. Xenon Fluorides

Three binary fluorid es of xenon are known: XeF2, XeF4 and XeF6. Xenon difluoride and xenon hexafluoride form numerous stable complexes with fluoride-ion acceptors. The fluoride ion donor ability of XeF4 has been shown to be less than that of XeF2 and XeF6: XeF6> XeF2 » XeF420 These findings are in agreement with the enthalpies of ionization derived from photoioni- zation studies'". HO(eV)

XeF2(g) -+ XeF+(g) + F-(g) 9.45

XeF4(g) -+ XeF3+(g) + F-(g) 9.66

XeF6(g) ->- XeFs+(g) + F-(g) 9.24 The enthalpy of ionization of XeF6 is more than 0.6 eV less than anti- cipated on the basis of XeF2 and XeF4 data. This is compatible with the greater stability of XeFs+ salts compared with XeF3+ and XeF+. It seems that the pseudo octahedral geometry of XeFs+ ion is especially favourable. Although there is no doubt that XeF6 is the strongest fluoride ion donor, there is disagreement as to the relative fluoride ion donor abilities of XeF2 versus XeF4. On the basis of the length, direction and number of fluorine bridges in the XeF+ and XeF3; salts it was concluded that the order of fluoride ion donor ability is XeF2 < XeF4 < XeF6 22. The facts that very few XeF3+ salts have been isolated and that they are thermally not so stable as those of XeF2 do not confirm the last statement. NOBLE-GAS COMPOUNDS 165

The simple cation XeF+,_l derived from each of the bin ary fluorides has been in each case crystallographically defined. Removal of F- from XeFx in each case leaves one short Xe-F bond trans to the site vacated by F-. The detailed structural arrangements suggest that the non-bonding xenon valence- -electron pairs are distributed as shown in Figure 1, with anionic species making their closest approaches to the cations in the directions indicated by the arrows.

o;:\\) ~ •. t

Figure 1. Shapes of the XeF\_l ions based on steric activity of the non-bonding xenon valence-electron pairs. (From D. E. McK e e, A. Za 1k i n, and N. Bar tle t t, Inorg. Chem. 12 (1973) 1713.)

2.1.1. Xenon Difluoride Xenon difluoride was first prepared independently in two laboratories-š-". After that, several effective syntheses were reported24,25,2G. If the thermal method of preparation is used27,28, a large excess of xenon over fluorine should be employed in order to minimize formation of higher xenon fluorides. A convenient preparation using very simple apparatus, affording a pure product, is the exposure of xenon-fluorine gas mixture to sunlightl3,14,29. Very pure

XeF2 in bigger quantities (up to 1 kg) can be best prepared by UV irradiation of a gaseous xenon-fluorine mixture in the mole ratio 1: 2 with addition of about 1 mol ofo of RF in fluor1ne as catalyst. The rate of XeF2 formation is thereby increased by about four times!", The thermal reaction between xenon and fluorine to form xenon difluoride is heterogenous and it takes place on the prefluorinated walls of the reaction vessel and/or the surface of added fluoride~1-3G.Among the examined fluorides, nickel difluoride proved to be an excellent catalyst for the xenon-fluorine reaction. In its presence it is possible to synthetize XeF6 even at 120 ce, starting from the gaseous mixture with the mole ratio Xe : F2 = 1 : 537.Fur- ther, it was shown'" that under certain conditions the reaction between xenon and fluorine could also be homogeneous. 166 B. 2EMVA

Although xenon difluoride is potentially a strong oxidizer, it is frequently unreactive for kinetic reasons. Its stability in aqueous solution (XeF2 may be recovered by CC14 extraction or fractional distillation) is typical of this kinetic inertness. Neutral or acid solutions decompose rather slowly (the half life being ~ 7h at O 0c) while in basic solution the decomposition is very fast, the base catalytic effect being roughly in the order of base strength. The hydrolysis products are xenon, oxygen, fluoride ions and hydrogen peroxide. The interaction of XeF2 with water is not only catalyzed by base but also by species which have an affinity for fluoride ions. Some physical properties of xenon .fluorides are given in Table 1.

2.1.2. Xenon Tetrafluoride The first discovered bin ary compound of noble gases was xenon tetrafluoride'". It was prepared by heating a gaseous mixture of xenon and fluorine in the mole ratio 1: 5 and total pressure of 0.6 MPa at 400°C in a sealed nick el can. These conditions are close to the optimum for XeF4 preparation. As Weinstock et al.47 have demonstrated, it is not possible to prepare pure

XeF4 by the thermal method. Low temperatures and high fluorine pressures favour XeF6 formation while low fluorine pressures and high temperatures favour XeFz formation. If high purity material is desired, it is necessary to purify it by chemical methods because it is very difficult to separate xenon tetrafluoride from the difluoride by physical means (the two have similar vapour-pressure relationships and also form a 1 : 1 adduct). Chemical purification could be done from a mixture of xenon fluorides so prepared that the content of XeFz is minimal and XeF6 is the major impurity. XeF6 could be removed by treating the mixture with sodium fluoride, which forms a complex with it at room temperatureš". A more general chemical purification'", which effectively eliminates XeF2 and XeF6 simultaneously ,exploits the inferior fluoride ion donor ability of XeF4 compared to XeF2 or XeF6. This purification is appropriate for a XeF4 sample which contains significant quantities of XeF2 and XeF6.

XeF'j o 'C in static vacuum -AsF5• -BrF5 XeF + AsFs (excess) XeF4 + (XeZF3) (AsF6t+ 4 ------+ XeF in BrFs soln. (XeFst (AsF6t 6 vacuum ---+, XeF4 20 'C

Recently, three new methods for the preparation of XeF4 have been deve- 49 52 loped - which give very pure XeF4 and additional purification is not necessary: - photochemical reaction between xenon and excess fluorine in the presence of NiF2 as catalyts49,50 thermal decomposition of XeF6 in the presence of NaF. (This method can also be used for the preparation of pure fluorineš')

pyrolysis of XeF+PtF6- or XePd2FlO at 150"c 52: TABLE I

Some Physical Properties of Xenon Fluorides

XeF2 XeF4 XeF6

Triple point/K 402.18" 390.25a 322.38 (P = 177.51 mmfIg)" l'.H(,ub) (kJ/mol at 298.15 K 55.71 ± 0.07' 60.92 ± 0.16" 59.1 ± 1.2° z Mi, (kJ/mol)', solid o at 298.15 K -162.76 ± 0.88 -267.11 ± 0.88 -338.2 ± 2.2 to g g fu Solid density/g/cm" 4.32 (calcd.) 4.04 ,,,(calcd.) 3.411 ± 0.015 at 297.55 K" 6 -3057.67 - 3226.21 ~ Vapor pressure/mm Hg log pa = log pa = log pb = -(3313.5/T) + 12.5923 o T T o -1.23521 log T -0.43434 log T ~ + + solid XeF6 between 291.8 and 322.38 K o q log P = - (3093.9/T) 11.8397 + 13.969736 + 12.301738 + solid XeF6 between 291.8 and 322.38cK u~l log P = - (6170.88/T)- 23.67815 log T valid between 273 and valid between 275 and + 80.77778 380 K 390.25 K + liquid XeF6 between 322.38 and 350 K Mean thermochemical bond energy (kJ /mol) at 298.15 Kr 132.30 ± 0.67 130.33 ± 0.54 125.27 ± 0.67

"Ref. (39); "Ref. (40);

f-' o'l -:) 168 B. ZEMVA

2 XeFPtF6(c) ->- XeF4(g) + XePt2FlO(c)

XePd2FlO(c) ->- XeF4(g) + Pd2F6(c)

XeF4 can be kept in thoroughly dried glass or quartz and can be sto red indefinitely in Kel-F, nickel or monel containers. The ultimate products of hydrolysis are Xe, O2, HF and Xe0353. The low bond energy of XeF4 causes it to be a strong oxidative f1uorinator (comparable with BrF3) but it has high kinetic inertness like XeF2.

2.1.3. Xenon Hexafluoride The preparation of XeF6 was first described in four independent and almost simultaneous reports19,54,55,56A. ll preparations are carried out in nickel or monel reaction vessels. From the equi1ibrium constants for the formation of xenon fluorides'" it can be seen that an excess of f1uorine (usually Xe : F2 = = 1 : 20) is necessary to minimize the formation of xenon tetraf1uoride. In addition, the reaction temperature should be as low as possible, yet consistent with a reasonable reaction rate. In the presence of nickel difluoride as catalyst'". the activation energy for the xenon-f1uorine reaction is lowered and the reaction proceeds rapidly at 120 CC, in contrast with 250 aC without a catalyst'".

The equilibrium constant for the reaction XeF4 + F2--+ XeF6 is much higher at 120°C (K = 330 atm ") than at 250 aC (K = 0.94 atrrr") 47. To purify XeFs, the crude material should be condensed on to NaF which forms a complex o with XeF6 at 50 C. Impurities (XeOF4, XeF2, XeF4) may be removed by pumping on the NaF mixture to constant weight at temperatures up to 50oC. Na2XeF8 is then decomposed at 125 aC under vacuum and pure XeF6 is col- lected in a cold trap". Small impurities of XeF4 in XeF6 can also be removed by treating it with excess of KrF2 at room temperature or a little higher=. The xenon hexaf1uoride is much more volatile-? than either XeF2 or

XeF4, although much less so than other hexaf1uorides. Crystalline XeF6 is effectively XeFs+Y-. The geometry of the XeFs+ and F- clusters59,60,61suggests that the nonbonding xenon(VI) valence electron pair is located on the four-fold axis of the cation, trans to the axial He-F bond. Thus described HeFs+ is pseudo-octahedral. The XeFs+ unit in (XeF6)x polymeric groups is almost indi- stinguishable from that observed in the XeFs+ salts. XeF6 is a much more powerful oxidizer and f1uorinator than the lower xenon f1uorides. It also has lower kinetic stability than XeF2 and XeF4• Thus, XeF6 cannot be stored in glass or quartz but reacts with it exchanging fluorine for oxygen to yield initially XeOF4 and finally xenon trioxide, with silicon tetraf1 uoride elimina tion:

XeF6 + x/2 Si02 ->- XeOxF 6-2x + x/2 SiF4 2.2. Xenon(II) Compounds All known xenon(II) compounds including molecular adducts are listed in Table II. Strong f1uoride ion acceptors, such as SbFs, AsFs and metal pen- taf1uorides, some metal tetrafluorides and BF3 can withdraw a f1uoride ion from XeF2 forming XeF+ species. The increase in the Xe-F bond energy in the cation formation (195.9 kJ/mol in the cation'", 133.9 kJ/mol in XeF2 (;2) contributes to fluoride ion donor ability of the dif1uoride. Furthermore, one NOBLE-GAS COMPOUNDS 169

or both of the fluorine ligands of XeF2 may be substituted by O (e. g. 99,104) or N (100) atoms rendered highly electronegative by pendant electro n with- drawing group s forming F-Xe-L or L-Xe-L (with L being -OS02F, -0C(0)CF3, -OPOF2, OSeFs, OTeFs, OCI03 -N(OS02F)z etc.). Xenon difluoride forms also molecular adducts with other hypervalent species in which the negatively charged F ligands are attracted to the positive centre of the partner molecule of the adduct.

2.2.1. [XeF+] and [Xe2F/] XeF2 complexes can be prepared by fusing stoichiometric amounts of the component fluorides under an atmosphere of a dry inert gas or by dissolving them in a non oxidizing or reducing solvent. One cornponent can also be used in excess and removed by vacuum evaporation to constant weight. Some reactions between binary fluorides in lower oxidation states and XeF2 have also been carried out in the melt of XeF2. Under these conditions, the oxidizing as well as the fluorobase properties of the XeF2 are revealed and Xe(II) complexes of higher oxidation state fluorides are formed:

MnF2 + nXeF2 -* XeF2 • MnF4 + Xe + (n - 2) XeF286. Xenon(II) compounds have also been prepared from hydrazinium or ammonium fluorometalates instead of binary fluorides'", This synthetic route has been extremely useful in the preparation of new xenon(VI) fluorometa- lates94,95,9G but it has not yielded xenon(II) fluorometalate which could not be prepared using other synthetic routes. The reaction system XeF2 - ammonium fluorometalates offers a unique synthetic route for the preparation of ammonium fluorometalates with the metal in higher coordination and oxidation states (e. g. NH4VF697, (NH4)NF893, (NH4)4UFlO 98). Xenon difluoride forms with pentafluorides three types of compounds: 2 XeF2 • MFs, XeF2 • MFs and XeF2 • 2 MFs. It is also possible that compositions poorer in XeF2 content could exist'". Xenon difluoride forms with tetrafluorides three types of compounds: 2 XeF2 . MF4, XeF2 • MF4 and XeF2 • 2 MF4• The 2 : 1 compounds between xenon difluoride and tetrafluorides do not exist in a dynamic vacuum at room temperature. Only the 1 : 1 and 1 : 2 compounds are vacuum stable. All appear as the XeF+ salts with polymeric anions (MFs-)xand (M2F9-)x' With trifluorides only 2 XeF2 . BF3 and XeF2 . BF3 are known. In XeF2 rich systems the V-shaped Xe2F3+cation usually occurs. It was first characterized by Bartlett and his coworkers'" and is plan ar and symmetrical about the bridging fluorine atom. It is characterized by strong Raman bands in the Xe-F stretch regi on at cca. 593 and 580 cm'" and in the bend region at 160 cm-I. The XeF+ ion was first defined structurally by Peacock & Russell+" and is characterized by strong vibrational bands in the 600-610 cm:" region'".

2.2.2. Molecular Adducts of XeFz

The first isolated molecular adduct of XeFz was XeF2 • XeF4 which was initially thought to be a second crystalline modification of XeF4. Since then many other adducts of XeF2 have been prepared. Some of them are listed in •..... -J o

TABLE II

Xenon(II) Compounds

Ti V Cr Mn Fe Co Ni 2XeF2· TiF 479,65,a,I> 2XeF2' VF58t,d XeF2·2CrF483 XeF2·MnF486 3XeF2' 2TiF 480 XeF2'VFs02 XeF2' 2MnF 486,87 XeF2· TiF 480 XeF2 .2TiF 480

Zr Nb Mo Tc Ru Rh Pd XeF · MoF 84,r 16 52 2XeF2· ZrF 4"°,0 XeF2' NbF582 2 o 2XeF2· RuF567 XeF2·RhF4 4XeF2·PdF4 ,1 b:! XeF2·2NbF582 XeF2' RuF567.82 2XeF2' RhF589 3XeF2' PdF452 XeF2·2RuF567.82 89 52 No XeF2· RhF5 2XeF2' PdF 4 M XeF2·2RhF589 XeF2' PdF452 ~ 52 Hf Ta W Re Os Ir Pt o 2XeF2·HfF4"5. 2XeF2'TaFs 73·<' XeF2' WFn8,I,f 2XeF2·OsF567 2XeF2·IrF567 2Xe2' PtF567 XeF2· TaF582 XeF2'OSF567,h XeF2·IrF5G7 Xe2·PtF567 XeF2' 2TaF5 71,82 XeF2· 2IrF567 Xe2·2PtF567 Xe(PtF6),t.16,j XePt2' F1016 U XeF2' UF685.•

Table II to be continued Table II continued

B C N O F Ne 2XeF2' BFs63.k XeF2' BF363.k Si P S Cl Ar 64 1 Al XeF25iF6 . 2XeF2' PFS36,k XeF2' PFSS6,k z Cu Zn Ga Ge As Se Br Kr O 2XeF2' AsFs67 XeF2·2BrFs74.1' oi XeF2' AsFs67," r;; 2XeF2(XeFsAsF6)91 6 68 ;J> XeF2' XeFs· AsF6 ,91 ul XeF2' 2(XeFsAsF6)68,91 o o ~ Ag Cd In Sn Sb Te I Xe "o 65 o 2XeF2' SnF 4 ,m 2XeF2' SbFs60,O XeF2' IFs 75,76 XeF2' XeF 478 ~ 3XeF2· 4SnF 466 XeF2' SbFs60,70 XeF2' 2IFs 770' Z 66 tj XeF2' 2SnF 4 2XeF2·3SbFs60,70 ul XeF2' 2SbFs 71 XeF2' 3SbFs69 XeF2' 6SbFs 76

Au Hg TI Pb Bi Po At Rn 2XeF2' AuFs90 2XeF2' BiFs 72 XeF2' BiFs 72 XeF2' 2BiFs 72

Foosnotes on the next side

f-" -J t-' 172 B. ZEMVA

TabLe II continued: Footnotes

e Ref. 79. The ečistence of the compound was established only by measurements of the electrical conductivity of XeF2/TiF4 melt as a function of the composition. h Ref. 65. The isolation of compound 2 XeF2·TiF4 was not conf irmed.š? The reportecl light yellow colour of 2 XeF2·TiF4 might be due to the parti al hydrolysis of the isolated material because all čenonrl I) and xenon(VI) fluorotitanates(IV) are white solids.š? C The isolation of 2XeF2·ZrF4 and 2XeF2·HfF4 was not conf'irrned'" even using the efficient route of hydrazinium fluorometalates and XeF2 (thus preparing tetrafluorides in situ in the very reactive, so called »molecular« form.). d The isolation of 2 XeF2• VF5 was not confirrned'". Stable 2: 1 compounds were not obtained either in the NbF5/XeF2 or TaF5/XeF2 systems . • Ref. 73. The way the cornpound was isolated and characterized does not support the existence of the 2 : 1 compound, which was neither confirmed by others'". r The existence of adducts XeF2· MoF6 and XeF2• WF6 was established from the melting point - composition diagrarn. g The existence of XeF2• UFe, was established from the melting point - composition diagrarn. h Compound XeF2' 20sF5 was not prepared. Even the 1 :'1 compound is thermally unstable at room temperature and decomposes to 2 XeF2' OS05

3 XeF+OsF- --r Xe2F30sF6- + Xe + 2 OSF6 o I The interaction of liquid XeF2 at 140-150 C with PdF4 (or Pd2F6) yielded a diamagnetic yellow material. Changes in slope in a weight-Ioss time curve indicated that 4: 1, 3: 1, 2: 1 and 1: 1 XeF2: PdF. complexes probably all occur but the 1 : 1 complex los es XeF2 at or below 20 =c. J In the historical oxidation of xenon by platinum bexafluoride it was found ' that the stoichiometry of the product of spontaneous reaction varied between XePtF ~ and Xe(PtF6h. Chemical and physical evidence indicated that the oxidation state of platinum in Xe(PtF6)xwas +5. Pyrolysis of Xe(PtF6)x, at 1650C, yielded XeF4 as the only volatile product and the solid, of composition XePt2FIO, was diamagnetic, thus suggesting a Pt(IV) compound. The product of spontaneous reaction between xenon and PtF6 at 200C contains XeF+PtF6-, as it was shown by X-ray powder photography. When the material of composition Xe(PtF6l2 was warmed to 60oC, the product [XeF"] [Pt2Fll1- was obtained thus indicating that, in addition to Xe+FPtF6-, amorphous PtF5 was also present. k Adduct 2 XeF2' PF;;, XeF2' PF5, 2 XeF 2' BF3 and XeF2' BF3 were not established with certainty and some additional work should be done in these two systems. I The reported compound Xe2SiF6 is probably erroneous. The authors claimed to have prepared it from a gaseous mixture of xenon and fluorine in the mole ratio 2 : 1 in a glow discharge. m The existence of the 2 : 1 compound was not confirrned?". The existence of 2 XeF~' 'SnF4 was not noticed in the XeF2/SnF4 system even during a very careful recording of the course of pumping off the excess of xenon difluoride at O aC. It is possible that the compound claimed to be 2XeF2·SnF4 was only XeF2·2SnF4 with some impurities. In favour of this conclusion are the following facts: the published infrared spectrum of the compound reported to be 2 XeF2' SnF 4 is practically identical with the infrared spectrum of the 1 : 2 compound'" and the colour of the 2 : 1 compound is reported to be lemon yellow, whereas pure xenon (II) and xenon(VI) fluorostannates are all white solid materials66•110 n All efforts to prepare a phase richer in AsF;; than XeF2· AsF5 failed and even the 1 : 1 compound loses AsF5 at room temperature to yield 2 XeF2· AsF5 67. o Ref. 69 and 70. In the system XeF2/SbF5 a whole set of compounds was postulated on the basis of the melting point - composition diagrarn. p The adduct XeF2' 2BrF5 was postulated on the basis of the melting point - composition diagram. r Ref. 77. It was reported that xenon tetrafluoride interacts with IF5 to yield IF, and an adduct XeF2· 2IF5. The authors?" did not find evidence for an adduct of that composition and, furthermore, they found that XeF4 dissolves easily in IF5 but does not form a thermally stable adduct with it. At ambient temperatures there was no evidence of evolution of xenon or of iodine heptafluoride formation. NOBLE-G.AS COMPOUlNDS 173

Table II. In these adducts the individual molecular species have essentially the same size and shape as in a crystal of the pure component. In all of them the F ligands of one molecule show maximum avoidance of other F ligands and are directed towards the positively charged central atom of the other molecule. This provides the best lattice energy for the semi-ionic assembly. The existence of XeF2· IF575,76 and XeF2· XeF468 and the isoelectronic relationship and near identity of shape of XeF5+ to IF5 and XeOF4 suggested that XeF2 might also make adducts with XeF5" salts. 2 XeF2 . XeF5+AsF6- ul, XeF2 . XeFs+AsF6- and XeFz . (XeF5+AsF6-)z68,91 are typical and are conveniently prepared by mixing the component in the appropriate molar proportions either by fusing the components or dissolving them in a suitable solvent. Such adducts are limited to mono anions, in which the fluorine-ligand charge is suf- ficiently low, so that XeF2 can effectively compete with the anion in interaction with the XeF5+'

2.3. Xenon(IV) Compounds Xenon tetrafluoride is the weakest fluoride ion donor among xenon fluorides and only adducts with the strongest Lewis acids (SbF5 and BiF5) have been isolated at room temperature (see Table III). Such adducts were first presented as XeF3+ salts on the basis OI the crystal structure of XeF3Sb2FI1102.

The existence of an adduct XeF4 • AsF5 has been claimed-'" at low temperatures and in the presence of an excess of AsF5.

TABLE III

Xenon(IV) Compounds

XeF4' AsF5101 XeF 4' SbF5102,103 XeF 4' 2SbF5102,10o XeF 4' BiFs 72 XeF4·2BiFs72

The known xenon(IV) compounds have been prepared by fusing the corn- ponents or by dissolving them in anhydrous RF, either in stoichiometric amounts or by using excess of one component and removing the excess under vacuum to constant weight. These complexes are best described in terms of a T-shaped cation, (XeF31) coordinated to the anion via fluorine bridges, the lengths and polar character of which are determined by the Lewis acidity of the anion. The totally syrn- metric stretching vibration VI of XeF3+is ~ 643 cm-lo

2.4. Xenon(VI) Compounds A number of adducts involving XeF6 in combination with recognized F- acceptors (Table IV) and recognized F- donors have been reported (Table VJ.

High conductivity of XeF6 in RF solutions, in comparison with those of XeFz or XeF4, indicates the possibility that XeF6 is extensively ionized in such solutions and is in fact [XeF5+]. F-(solv). The adducts with fluoride-ion acceptors all appear to be either XeF5' or XezFIl+ salts, the prototype structures of which were XeF5+PtF6-128 and Xe2FI1+AuF61-05. These cations are the Xe(VI) analogues of XeF+ and Xe2F3+' ~ •..... , 4 -:J TABLE IV *'" Xenon(VI) Compounds with p- Acceptors Li Be

Na Mg

K Ca Se Ti V Cr Mn Fe Co 116 XeF6' SeF3 ,a 4XeF6'TiF48o,d 2XeF6'VF5118 XeFo'CrF4121 4XeF6' MnF 487 XeFo' FeF394,g XeFo' COF3V4 XeF o' TiF 480 XeF6'VF588 2XeF6'MnF487 XeF6' 2TiF 480 XeF6' 2VF588 XeF6'MnF487 XeF6' 2MnF 487

Rb Sr Y Zr Nb Mo Te Ru Rh ~ 122 3XeF6' YF3117,b XeF6'ZrF495,e 2XeF, 6NbF5119 2XeFo' MoF6 " 2XeF6' RuF5105,h N 124 t

La and Cs Ba Lantanides Hf Ta w Re Os Ir 4XeF6'CeF4117,C XeF6'HfF495,e 2XeF6' TaF5120 2XeF6' IrF520,i XeF6' 2CeF 4117 XeF6' TaF5120 XeFo' IrF520 4XeF6'PrF4117 XeF6'2PrF4117 XeF6'4PrF4117 4XeF6,TbF4117 XeF6'2TbF4117 3XeF6' DyF3117 3XeF6' HoF3117

U 2XeF6' UF5123 XeF6' UF5123 Table IV to be continued Tabte IV continueđ ~ B C N O F Ne XeFo' BF3107,163

Al Si P S Cl Ar XeFo' 2AIF396,m 2XeFo' PF5111

Ni Cu Zn Ga Ge As Se Br Kr 4XeFo' NiF4126,j XeFo' GaF396,m 4XeF6' GeF4108 2XeF6' AsF5111 2XeFo' NiF 4126 2XeFo' GeF4108 XeFo' ASF5107,163 XeFo'GeF4108 Z o to Pd Ag Cd In Sn Sb Te I Xe r- M 4XeF o' PdF 4127 XeFo' AgF3120,k 4XeFo' SnF 4109," 2XeFo' SbF5112 t 2XeF 6' PdF 4127 2XeFo'SnF4109 XeFo' SbF5112 ;pO 4XeFo'3SnF4110 XeFo' 2SbF5112 u: () 3XeFo'4SnF4110 o XeFo' 4SnF 4110 ~ o c:: Pt Au Hg TI Pb Bi Po At Rn Z IJ 2XeFo' PtF5128 2XeFo'AuF5t30 4XeF6' PbF411O," 2XeFo' BiF5113 u: XeFo' PtF5128 XeFo'AuF5130 3XeF6'4PbF4110 XeF6' BiF5114 XeFo' AuF3165,\ XeFo'4PbF4110

Foosnotes on the next siđe •..... -J en 176 B. ZEMVA

Table IV continued: Footnotes n XeF6' SCF3 has been prepared only by the reaction between N2H5' ScF 4 and excess of XeF6 at room temperature. The adduct is stable at O ° but it decomposes at 200C to SCF3 and XeF6. The reaction between so obtained SCF3 and XeFo to form XeF6' SCF3 is not reversible. b The compound 3 XeF6' YF3 is the first known monomeric salt of XeF5+ with a metal in the +3 valence state. Raman spectroscopy shows that the compound is an XeF5+ salt. It can be formulated as (XeF513YF63-. It has been prepared by the reaction between Y203 or N2H5YF4 and excess XeF6.

c With the lanthanides, two types of XeF5 salts have been synthetized. If the lanthanide is oxidized to the +4 ooidation state, the 4: 1 and 1: 2 compounds are formed (in the case of praseodymium also '1 : 4). Only the 1: 2 compounds are thermally stable at room temperature (1: 4 in the case of praseodymium). If the lanthanide is not oxidized to the + 4 oxidation state, the 3 : 1 compounds are obtained when the lanthanide trifluoride has the YF3 structure. When the trifluoride had the LaF3 structure, no xenon(VI) compounds were preparable. d The evidence for the 2 : 1 compound was found only durmg the thermal analysis of the compound with the high xenon hexafluoride content. The small difference between the decomposition temperatures of the 4 : 1 and 1 : 1 cornpounds hampers the isolation of the 2 : 1 compound. e Compounds XeF6' MF 4 (M = Zr, Hf) were isolated by using hydrazinium fluorometalates as starting materials for the reaction with XeF6. The reaction between the corresponding tetrafluorides and XeF6 did not proceed.

r The existence of 2 XeF6' MoF6 and XeF6' MoF6 was established only from the melting point - composition di agram. • XeF6' FeF3 is another example where the use of a hydrazinium salt is essential for the synthesis of an XeF5+ salt. Raman spectra establish that it is an XeF5' salt. The anion is probably polymeric (FeF4)nn-. h The synthesis of 2 XeF6' RuF5 has not been described. However, the compound is isostructural with the other 2 XeF6' MF5 where M is Ru, Ir, Pt and Au.

1 The syntheses of 2 XeF6' IrF5 and XeF6' IrF5 have not been described. For the 2 : 1 compound see h). The 1 : 1 compound is iso structural with XeF6' MF5 where M = Ru, Pt, Nb.

j Both nickel compounds have been prepared using XeF6 and KrF2 simultaneously, The 4 : 1 compound decomposes at room temperature yielding XeF6 and the 2 : 1 compound, k The XeF6' AgF3 compound was synthetized only by simultaneous use of XeF6 and KrF2.

1 The XeF6'AuFg compound represents with XeF6AuFG the first instance of XeF5 occurring with two different valence states of a metal. no Compounds XeF6' GaF3, XeF6' 2 AlF3 and XeF6' InF3 are additional examples of the synthetic route using hydrazinium fluorometalates as the starting materials for the preparation of XeF;;+ salts. All three compounds are XeF5+ salt with polymeric anions. The indium compound is therrnally unstable, decomposing at 20 ° to XeF6 and InF3. u On the basis of vibrational spectra, the 4: 1 compounds can be formulated as Xe2Fll;' salts, while the 1 : 4 compounds are XeF5+ salt with polymeric anions. The 4: 3 and 3: 4 compounds are apparently intermediate compounds with both Xe2Fll+ and XeF5+ cations present. o The existence of an adduct XeF6' KrF2 was not confirmed 168 us ing the same experimental approach as the authors!" and also other synthetic routes. NOBLE-GAS COMBOUNDS 177

2.4.1. [XeFs+] and [Xe2Fll+] For the preparation of xenon(VI) fluorocomplexes many synthetic schemes have been used, which can be roughly represented by the following three groups: a) Direct method: The interaction of an excess of xenon hexafluoride with a binary fluoride in a melt of XeF6 (temperatures higher than 50°C.) Several complexes have been prepared by oxidizing a lower-valent fluorides in the presence of an excess of XeF6. In other cases the complexes have been prepared by mixing the component fluorides in a non reducing solvent. The stoichiometries of these complexes appear to be determined by the favoured coordination number of M (commonly six as in [MFn, [MF62-]) and the formation of either XeFs+ or Xe2FIl+. The latter is formed only when the anion is monomeric (MFxn-) and XeFs is available to form a p- bridge to XeFs+. b) Indirect method: The corresponding bin ary fluoride is sythetized in situ in a very reactive »molecular form«. In this case xenon hexafluoride is both oxidizer and complexing agent. As starting materials, hydrazinium salts or ammonium fluorocomplexesv"?" or oxides'!? have been used. Using this synthetic route a whole series of xenon(VI) fluorocomplexes have

been isolated (e. g. XeF6·FeF3, XeF6· GaF3, XeF6' InF3, XeF6· 2 AlF3, XeF6 . HfF4, XeF6 . ZrF4). c) Combined Method: In this method, KrF2 and XeF6 were used concurrently. This method is efficient in cases where the binary fluoride does not exist in the oxidation state in which the xenon(VI) fluorocomplexes are expected (e. g. xenon(VI) fluoronickelatesđ.Vj-") or in the cases where the bin ary fluoride exists in the oxidation state appropriate for the formation of xenon(VI) fluorocompound, but where this binary fluoride could not be obtained by using only XeF6 as oxidizer (e. g. xenon(VI) fluoroargen- ta te(III)129). Xenon hexafluoride forms complexes with monomeric and polymeric anions. In the case of monomeric anions we have the following set of complexes: Xe2Fll+MF6-, XeF5+ MF6-, XeFs+M2Fll-, (XezF1t)2 MF62-, (XeFs+h MF62-, (XeF;h MF63-. In the case of polymeric anions we have with tetravalent fluorides the following mole ratios: XeF6: MF4 4: 3, 3: 4, 1: 2, 1 : 4110,87and with MF3 1 : 1 and 1: 294,96. Among all structurally defined xenon(VI) fluorocomplexes, only in the cases of XeF6 . AgF3 and XeF6 . . AuF3 the anions are not octahedrally coordinated with fluorine. The Xe2Fll+ cation behaves vibrationally like two weakly coupled XeFs+ species. The F bridging of two XeFs groups in the Xe2Fll cation appears to be characterized by a »bridge stretch« at cca. 360 cm-lt05. For XeFs+ of C4v symmetry, nine fundamental vibrations were observed.166,167

2.4.2. [XeF7-] and [XeFi-] Complexes of XeF6 with fluoride ion donors, e. g. N02F, NoF, and alkali metal fluorides (with the exception of LiF) are listed in Table V. With alkali fluorides, complexes of the type MXeF7 and MzXeF8· have been reported. They

) 178 B. ZEMVA probably contain anions XeF7- and XeFs2-. The XeFs2- in the fully characterized compound (NO)z+ XeFs2- 134 has an approximately archimedian antiprism geometry, the stereo activity of the non-bonding xenon valence electron pair being only weakly evident. NF4+ XeF7- decomposes to NF3, F2 and XeF6 and evidence for (NF4)z XeFs as an intermediate has been given133•

TABLEV

Xenon(VI) Compounds with p- Donors

2 NaF . XeF657.125 N02F . XeF6131 2 KF . XeF6125 2 NOF . XeF6132 2 RbF . XeF6106.125 NFl XeF7-133 RbF . XeF61060125 (NF4+l2 XeFs2-133 2 CsF . XeF6106.125 CsF . XeF6106,125

3. KRYPTON CHEMISTRY The chemistry of krypton is limited to that of krypton difluoride and its derivatives. All efforts to confirm the synthesis of krypton tetrafluoride or higher fluorid es have failed. Furthermore, there is no firm evidence of the existence of oxides, oxyfluorides or chlorides.

3.1. Krypton Difluoride Krypton difluoride was first prepared by Turner and Pimentel+" who prepared it by u1traviolet photolysis of fluorine suspended in a solid mixture of argon and krypton at 20 DK. Because krypton diffluoride is thermodynamically unstable-"; all effective syntheses have to be carried out at lower temperatures (-196 DC). Methods used for the preparation of KrF2 in the laboratory include electric discharge of a gaseous mixture of krypton and fluorine at low temperature and pressure+". The method is claimed to yield about 250 mg KrFzlh. High frequency discharge in a mixture of krypton and CFzClps has also been claimed. Other methods involve irradiation of fluorine - krypton mixtures: with 1.5 MeV electrons-š"; with 10 MeV protonsv? OJ;' (with low efficiency) exposure to sunlight-v. A very efficient method for the preparation of KrF2 on a bigger laboratory scale has been developed (30-40 g/chargej-". This is based on the irradiation of a liquified mixture of fluorine and krypton with neal' UV light at -196 DC. The yield is about 1 g of KrFzlh. Properties: krypton difluoride is colourless in the solid and vapour phase. It decomposes at temperatures above -80 DC. The decomposition rate is very much dependent upon the pretreatment of the reaction vessel. If this is well pretreated with KrFz before starting, the decomposition of KrFz is much lower. Krypton difluoride has the lowest average bond energy of any known fluoride. The atomization energy of KrFz (tl Hat = 98 kJ mol-I) is smaller than the dissociation energy of elemental fluorine (tl Hdiss.= 153 kJ mol-I), thus indicating that krypton difluoride is closest in activity to atomic fluorine. NOBLE-GAS COMPOUNDS 179

TABLE VI

Some Physical Properties of KrF2 !J. H,ub/kJ mol-Ia 41 !J. Hfo(g)/kJ mol-Ib 60 ± 3 solid density/g cm'?" 3,301(calcd.) vapor pressure/mm Hga 10 ± 1 at 257.7K 29 ± 2 at 273.2K 73 ± 3 at 288.2K mean thermochemical bond energy (kJ mol:")" 49.0

a Ref. (143); b Ref. (136); c Ref.(144).

3.2. Krypton(II) Compounds So far only those krypton-fluorine compounds are known which exhibit a significant amount of ionic character and can be described in terms of (KrF)+ and (Kr2F3)+containing species. The (Kr2F3)+cation consists of two (KrF)+ un its joined by a fluorine bridge. All known krypton compounds are listed in Table VII. They are complexes of KrF2 with Lewis acid pentafluorides and have the stoichiometries 2 : 1, 1 : 1 and 1 : 2. In addition, there are a number of compounds in which additional KrF2 units appear to be weakly associated, probably with the cationic part of the compound. Recently, also the first adducts of KrF2 with tetravalent fluorides have been isolated: 2 KrF2 . MnF4 and KrF2 . MnF4153. The 1 : 1 compound is most probably (KrF)+ (lvInFst with a polymeric anion. Claims for compounds 2 KrF2 . TiF4 and 2 KrF2 . SnF, have not been supported by any characteriza tional evidence. Krypton complexes may be prepared directly by the reaction of KrF2 with the appropriate pentafluoride or using a solvent. In some cases the pentafluoride has been prepared in situ using KrF2 as the fluori.nating agent. So far no X-ray crystal structure of any KrF2 complex has been deterrnined and the characterization of the complexes has rested largely on vibrational spectroscopy. Raman spectroscopy has been especially valuable in the studies of these compounds, all of which are of low thermal stability. Observed Raman and infrared frequencies for Kr(II) compounds are best assigned on the basis of ionic formulations. (KrF)+ (M2Fll)-, (KrF)+ (MF6)- and (Kr2F3)+(MF6f. The appearance of Kr ... F and M ... F bridge-stretching and F-Kr ... F bending modes point to a strong interaction between cation and anion indicative of some covalent character to that bonding. For the (KrF)+ salts the stretching vibration is in the range from 626 cm"! (in the case of KrF+ Sb2Fll-) to 597 crrr" (in the case of KrF2• 2 NbFs). The spectra of the (Kr2F3)+cation have been correlated with those of the (Xe2F3)+species but additional peaks, which have no equivalents in the xenon spectra, suggest that the (Kr2F3)+is less symmetrical than its xenon analogue, with one long er and weaker Kr-F bond and one shorter and stronger Kr-F bond. An alternative interpretation is that this cation involves distorted KrF2 molecule fluorine bridged to a (KrF)+ cation. KrF2 is a very strong oxidizer. KrF+ salts are even more powerful electrophiles since the electron affinity of (KrF)+ is 12.5 ev' This makes it an effective source of F+, exploited effectively in the synthesis of BrF6+ and ClFt salts. •..... 00 O

TABLE VII K1'ypton(II) Compounds

Ti V Cr Mn Fe Co Ni 4145 KrF2' VF5150 2KrF ' MnF 4153 2KrF2' TiF 2 SO KrF2' MnF 4153 N t%J Zr Nb Mo Te Ru Rh Pd ~ (xKrF2' KrF)+' (Nb2F11r 151 KrF' .RUF6-154 KrF+'RhFo-154 KrF+'PdF6-154 ;<.. Kr?' Nb2Fll-152 KrF+'Ru2Fl1-154

Hf Ta W Re Os Ir Pt (xKrF2' KrF+)' (Ta2Fllf151 KrF+' PtFO-147.140 Kr2F3+'TaF6-151 Kr ?Pt2F 11-154 KrF"TaFo-152 KrF+'Ta2Fl1-151 Table VII to be continued .1

TabLe VII continued B C N O F Ne KrF+'BF4-145

Al Si P S Cl Ar

Cu Zn Ga Ge As Se Br Kr Z O (xKrF2' Kr2F3+)(AsF6r 146,147 KrF2' BrFS140 to Kr2F3+' AsF6-146,147 r- M KrF+' AsF6-146,147 6 ~ Ag Cd In Sn Sb Te I Xe ul 145 (xKrF2' Kr2F3)+' (SbF6-)146 115 o 2KrF2' SnF 4 KrF2' XeF6 O (xKrF2' KrF)'· (SbF6'V47.148 a;: 148 '"d Kr2F3+' SbF6- ,146 o KrF+' SbF6 -148,146 q KrF+' Sb2Fl1-48,102 Z ti ul Au Hg TI Pb Bi Po At Rn KrF+'AuF6-oo KrF+BiF6-72

-co 182 B. ZEMVA

4. RADON CHEMISTRY There are no stable isotopes of radon and the long est lived »natural« isotope, 222Rn, has a half-life of only 3.83 days. Experimental difficulties arise not only from radiation hazard, but also from the radiation decomposition of the reagents. The most detailed review of radon chemistry, to early 1981, was made by L. Stein!55. The low first ionization potencial of radon (10.7 eV) and its chemical similarity to xenon indicate that a number of other radon compounds apart from RnF2 should exist, such as RnF4' RnF6, RnCb, RnC14, RnO, Rn02, Rn03 and RnOF4. The existence of. RnF4 or RnF6 and Rn03 has been claimed.P" So far no complex compounds containing (RnF), (Rn2F3)+or (Rn)2+have been isolated and fully charazterized. Radon reacts spontaneously at room temperature with many solid compounds that contain oxidizing cations (e. g. XeF+ and O2+salts). By analogy with xenon, it is bellieved that radon also forms RnF+ salts.

5. CHEMISTRY OF LIGHTER NOBLE GASES The first ionization potentials of lighter noble gases Ar (15.76 eV), Ne (21.56 eV) and He (24.59 eV) show that these gases will remain truly inert gases. The only chance is probably for the formation of ArF+ cation in con- nection with a high ionization potential anion, which is also a very poor fluoride ion donor. A possible candidate is the Sb2Fll- ion but AuF6- could also serve the purpose. The 2 + oxidation state for krypton and xenon are more stabilized in the diatomic noble-gas fluoride cation (G-F)+. The energy of the bond in KrF+ is 1.63 eVG2as compared to XeF+ which is 2.03 eV21 and ArF+ which is 1.655 eV157.

6. CHALLENGE S OF THE NOBLE GAS CHEMISTRY IN THE FUTURE An extension of noble gas chemistry to other elements than Kr, Xe and Rn is possible only with argon in ArF+ salts. Further, new oxidation states and a greater range of ligands migth be used in radon chemistry. Otherwise, the extension of the chemical compounds of xenon and krypton in known oxidation states with other electronegative ligands might occur. The high proton affinities of the heavier noble gases158 (Xe;::: 6 eV; Kr > 4 eV) also raise the possibility of (XeH)+ and (KrH)+ salts being preparable. These are of potential value as protonating agents. The greatest extension of noble gas chemistry is expected in the exploitation of noble gas compounds as fluor i- nating and oxidizing agents.

7. POSSIBLE USES OF THE NOBLE GAS COMPOUNDS The greatest utility of the noble-gas compounds is derived from the weakness of their bonds and the relative inertness of the reduction product, the noble-gas atom. Their greatest potential value, therefore, lies in their employment as clean oxidizers. The lighter the noble-gas atom the greater is the oxidizing power of the compound. Because of the ready ionization of noble-gas fluorides to form cationc species (XeF+, KrF+ etc.), they are fast and efficient electrophilic oxidizers. NOBLE-GAS COMPOUNDS 183

The use of XeF2 as a fluorinating agent for the preparation of some ammonium fluorometalates with metals in higher oxidation state (e. g. (NR4)4UF!O)98suggests that it can have a valuable role as a potent electrophilic oxidizing agent in the generation of high oxidation state species. It has also been used as a source of F ligands in the substitution of F ligand into aromatic hydrocarbons-š" and as a route to oxidative reagents which do not contain a noble gas atom (e. g. source of S03 radicals)!". The interaction of XeF2 with water as a component in the inetraction has also been an effective route to the synthesis of perbromate'v". This could be further exploited as an effective [O] atom source: XeF2 + R20 ~ [XeO] + 2 RF. The most spectacular reagents so far have proven to be the GF+ salts (G = noble gas atom). They act as suppliers of F+ (synthesis of BrF6+ 161 etc.).

Reactions between radon and strong oxidizers, such as O2+ salts, can be used for scrubbing radon from air. This is of particular value in uranium mines.

AeknowLedgement. - The author is grateful to the Research Community of Slovenia for supporting this work.

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POVZETEK Pregled dela na sintezi binarnih fluoridov žlahtnih plinov B. Žemva Podan je kritični pregled dela na sintezi binarnih fluoridov žlahtnih plinov (posebno ksenona in kriptona) in njunih spojin v obdobju zadnjih 25 let. Peleg tega je podan tudi zgodovinski uvod v kemijo žlahtnih plinov, izziv za bodočnost in mož- nosti uporabe spojin ž1ahtnih plinov.