658 IN BRITAIN JULY \987 Twenty-five years of chemistry

John H. Holloway

This year we celebrate the 25th anniversary of the discovery of noble gas that can be made by direct synthesis are chemistry. The discovery toppled a belief that had grown for almost 70 the fluorides. Most compounds that have years-that no member of the family of 'inert gases' could form a been obtained consist of noble gas atoms thermodynamically stable . Now, a quarter of a bonded to the most electronegative century on, it is appropriate to look back at the origins, development and elements, and , but species bonded to chlorine, nitrogen and more noteworthy recent advances in the field. carbon have also been reported. In March 1962 Bartlett and Lohmann Chemistry in Britain. The second of these chemistry is the most extensive-Xe is reported the preparation of O2+[PtF6r.1 short papers, which was less than 250 known in oxidation states ranging from The production of O2+, by simply mixing words in length (Fig. 1), initiated what +2 to +8, while and species oxygen with an equimolar quantity of was to become a flood of activity in are known only in +2. at room tempera­ an entirely new area of chemistry. Stable compounds of , and ture, showed PtF6 to be 'an oxidizer of Before the end of August 1962, Howard are not known, but experimental unprecedented power',2 capable of oxi­ Claassen, Henry Selig and John Maim estimates of the bond strength in the dising molecular oxygen. Because the had prepared the first binary fluoride of a [ArFt cation12 suggest that continued first ionisation potential of molecular noble gas, XeF4,4 and by the end of the efforts to lock this into an adduct are oxygen (12.2 eV) is slightly more than that year, Rudolf Hoppe and his group had worthwhile. of xenon (12.13eV) Bartlett thought that announced the synthesis of XeF2,5 Jozef xenon gas might be similarly oxidised. Slivnik and his coworkers the preparation Xenon fluorides11 •13 Within a month he had confirmed his of XeF6,6 and Paul FieJds, Lawrence Stein Apart from XeF, which has been obtained suspicion and, in June he published a and Moshe Zirin the first evidence of the as an unstable free radical, there is no report of his observation of the spon­ existence of a radon fluorideJ By early evidence for the existence of fluorides in taneous and rapid oxidation of xenon by 1963* the existence of compounds of odd-numbered oxidation states. The the deep red PtF6 vapour to give a solid, krypton, xenon and radon was firmly fluorides, XeF2, XeF4 and XeF6 can be yellow-orange, quinquevalent platinum established. 11 obtained by reacting xenon directly with complex.3 Time has confirmed the validity of fluorine, but early reports of XeFa have The papers outlining the discoveries these early observations. Stable com­ not been confirmed. appeared in the Proceedings of the pounds are formed only with the heavier • The difluoride is prepared by heating Chemical Society,1.3 the predecessor of gases, and the only stable compounds an excess of xenon with fluorine or by irradiating mixtures of the gases with Fig. 1 . Bartlett's short paper announcing the discovery of Xe1PtF ef and carrying his ultraviolet light or sunlight. signature. Reprinted from Proceedings ofthe Chemical Society, June 1962, p 218. • The tetrafluoride is prepared by heat­ ing xenon and fluorine under pressure in a ratio of 1 :5 by volume. Xenon Hexaflaoroplatiate(v) Xe~.PtF.J­ • The hexafluoride is prepared by heat­ By NEIL BARTLETT ing xenon and fluorine under pressure in (DEPARTMENT OF CHEMISTRY, THE UNIVERSITY OF BRITISH COLUMBIA, a ratio of 1 :20. VANCOUVER 8, B.C., CANADA) All the xenon fluorides are colourless A RECENT Communicationl described the compound The composition of the evolved gas was established solids except the hexafluoride, which is dioXYlCllyl hexaftuoroplatinate(v), Ot.. PtF.-, which by mass-spectrometric analysis. yellow in the liquid and gaseous states. is formed when molecular Ol\;ygen is ol\;idised by Although inert-gas c1athrates have been described, The di- and tetrafluorides have low platinum hexaftuoridc vapour. Since the first i~nisa­ this compound is believed to be the first xenon tion potential of molecular ol\;y~n,2 I L! ev, is com­ charge-transfer compound which is stable at room volatilities, and linear and square planar parable with that of xenon,z 12·\3 ev, it appeared temperatures. Lallice-encrgy calculations for the structures in the solid and vapour states. that xenon might also be ol\;idised by the hexa­ xenon compound, by means of Kapustinskii's equa­ Of the 16 known , xenon fluoride. tion,3 give a value -- 110 kcal. mole-I, which is only Tensimetric: titration of xenon (AI RCO .. Reagent 10 kcal. Olole- I smaller than that calculated for the hexafluoride has the highest boiling Grade") with platinum hexaftuoride has proved the dioxygcnyl compound. These values indicate that if point. Th~s suggests that, unlike the existenc:e of a I: I compound, XePtF•. This is an the compounds are ionic the electron affinity of the other hexafluorides, this compound oran~yeUow solid, which is insoluble in carbon platinum hcxaftuoride must have a minimum value might be polymeric in the solid state. tetrachloride, and has a nqiiaible vapour pressure at of 170 kcal. mole-I. room temperature. It sublimes in a vacuum when However, uncertainty about its struc- heated and the sublimate, when treated with water The author thanks Dr. David Frost for mass . ture remains, mainly because it is readily vapour, rapidly hydrolyses, xenon and oxygen being spectrometric analyses and the National Research hydrolysed to the explosive Xe03 and evolved and hydrated platinum dioxide deposited: Council, Ottawa, and the Research Corporation for its strong fluorinating ability damages lXePtF. + 6H.O -+ lX. + O. + lPtO. + UHF financial support. (R,cei",d, May 4th, 1962.) most container and window materials, I Butlett and Lobmann, Proc. CM"'. 5«.. 1962, 115. • Field and Franldin, "Electron Impact Phenomena," Academic Prest, Inc., New York, 1957, pp. 114-116. making it difficult to study. Unlike other • Kapustinsltii, QIllUI. RI.... , 1956, .8, 284. hexafluorides it is not octahedral in the vapour but is probably a non-rigid (floppy), distorted octahedron. 14 When molten or

• Early in 1963 it emerged that three other groups.8-10 working independently. had all prepared XeF, at .. about the same time . 660 CHEMISTRY IN BRITAIN JULY 1987

solution and powerfully oxidising, con­ 2 ~) ~) verting for example 2Cr to C1 2, Cr(lIlkto Cr(VI) and even [Br03r to [Br04r.17 The tetra- and hexafluorides are progressively more powerful oxidisers and difficult to control. Tt"\eir instant hydrolysis by water, however, has provided interesting , fluorides and powerfully oxidising xenate and salts . • Hydrolysis reactions. The difluoride, XeF2, is stable in acid or neutral solution but decomposes immediately in base, liberating xenon, oxygen and HF. The tetrafluoride is instantly hydrolysed and, depending upon conditions, up to a third of the xenon may be retained in solution (rather than the half expected in a simple disproportionation, 2Xe(lv) - Xe(vl) + Xe(II)), and XeF2 may also be liberated. It has been suggested20 that the initial product may be XeOF2, which de­ composes to give XeF2 if the water supply Fig. 2. The structure of liquid XeFs. is limited. When larger amounts of water are present Xe02 and XeO may be 3 formed, which can then give Xe03' [111] The more vigorous reaction of XeF6 i with water can be controlled by passing dry nitrogen over crystalline XeF6 to sweep the vapour into water where the following reaction occurs: XeF6 + 3H20 - Xe03 + 6HF (e) Thus the hydrolysis of both XeF4 and XeF6 give the trigonal pyramidal Xe03 in solution. Aqueous solutions of Xe03 (known as 'xenic' acid) are quite stable but solid Xe03, which is easily obtained from the solutions, is a sensitive and violent explosive. • Reactions giving oxides, oxide fluorides and oxygen-containing salts. In aqueous solution Xe03 is a powerful, but kin­ etically slow, oxidising agent. The ad­ Fig. 3. The crystal structure of the cubic phase of XeFe (-80 0 C). One unit cell dition of alkali yields 'xenate' , contains 144 'XeFs' units, 24 tetramers and eight hexamers.1s [HXe04L and, although the salts can be isolated, solutions containing xenate in solution the mono mer is in equilibrium stemmed from reactions of the binary disproportionate readily to give the with tetrameric rings consisting of [XeFsr fluorides. 'perxenate' , [Xe06]4-, and xenon: and F ions, but at low temperatures the • As fluorinating agents. The difluoride 17 18 monomer disappears. The in is a versatile mild fluorinating agent. • 2[HXe04r + 20H- - [Xe06r- + Xe + O2 + . 2H 0 (t) the ([XeFsrF)4 tetramers take part in a It provides routes to some aromatic 2 complex scrambling mechanism in which compounds which are otherwise difficult The Xe(VIII) solutions are among the each of the four units is connected as a to make and reacts with carbon-carbon most powerful oxidising agents known. cogwheel to its neighbours (Fig. 2).'S In double and triple bonds, giving addition For example, the oxidation of Mn(lI) to products. The cleanliness of these reac­ the solid, although at least four forms [Mn04L unlike that by Xe03, is immedi­ exist, only the cubic structure-compris­ tions is of special significance with some ate and is accompanied by the evolution ing [XeFsr and F ions grouped as of the products, which are of medicinal of oxygen. tetramers and hexamers within the unit and biological importance. One example The addition of solid Ba2XeOS to cold cell-has been described in detail (Fig. is the ability of XeF to convert uracil to 2 concentrated H2S04 produces a second 3).'S The fact that both hexamers and 5-fluorouracil, one of the first antitumour oxide, Xe04, which is an unstable and tetramers occur suggests that the way in agents. explosive gas. Understandably, it has which the [XeFsr ion is bridged is not Xenon difluoride is also a common been studt.ed little, but infrared spec­ critical. Furthermore the bridging F ions inorganic oxidative fluorinating agent, troscopy and electron diffraction have are not close to the four-fold axis of the see equations (aHd).'7 shown it to be a tetrahedral molecule. [XeFsr groups. This suggests that in keep­ CH31 + XeF2 - CH31F2 + Xe (a) Of the known oxide fluorides, XeOF2, ing with the Gillespie and Nyholm valence .xeOF4, Xe02F2 and Xe03F2, the most shell electron-pair repulsion theory, HF stable is the liquid, XeOF . Colourless 2Sb + 3XeF2 - 2SbF3 + 3Xe (b) 4 which has provided the most accurate and volatile, with a square pyramidal and simple rationalisation of the shapes HF structure, it derives from the controlled SbF + XeF - SbFs + Xe (c) of noble gas compounds, a lone pair of 3 2 hydrolysis of XeFs: electrons may occupy this position. CFzCI.CFClz Re~CO),o + 3XeF2~ XeFs + H20 - XeOF4 + 2HF (g) Re(CO)sF.ReFs + 5CO + 3Xe (d) Reactions of the xenon fluorides This reaction, equation (g), readily goes With the exception of the reactions of Recently it was used in the preparation of further to produce Xe03, but the ex­ powerful oxidative fluorinating agents CrOF3.19 Soluble also in water, the pected intermediate, Xe02F2, has never such as PtFe, the chemistry of xenon has difluoride is stable in neutral or acid been isolated from the reaction. But,

t a 662 CHEMISTRY IN BRITAIN JULY 1987

Xe02F2 can be produced via the inter­ product of the reaction, which decom­ In accord with the enthalpy of ionis­ action of with the oxide poses to give XeF2 and carbon fluorides ation data, xenon tetrafluoride forms tetrafl uoride. at room temperature, is thought to be complexes with only the strongest fluor­ Xe( CF3h·24 ide ion acceptors-SbF and BiF5-but Xe03 + XeOF - 2Xe02F2 (h) 5 4 • Reactions with fluoride-ion acceptors xenon hexafluoride yields adducts with a The related and possibly even more and donors. 25 Reactions of fluoride-ion number of pentafluorides and also with hazardous reaction with Xe04 yields donors and acceptors with the binary some tetra- and trifluorides. Solutions of Xe03F2: fluorides and oxide fluorides of xenon XeFs in HF also have a higher conductivity give a variety of cations and anions than those of XeF and XeF4 because Xe04 + XeFs - Xe03F2 + XeOF4 (I) 2 (Table 1). More unusual entities such as both [XeF5f and [Xe2Fllf are readily • Substitution reactions. Reactions in­ [Xe2f and [Xe03Xr (X = F, Cl or Br) also formed in solution: volving the replacement of fluorine in occur. The representation of many of the xenon fluorides with other groups began complexes as salts, however, is an over­ ([XeF5fF-)4 + nHF ~ with studies of reactions of XeF2 with simplification, especially for the xenon 2[Xe2Fllf + [(HF)nFr (P) anhydrous oxygen-containing acids: 17 difluoride adducts. Also, some adducts of XeF -for example, those with XeOF 4 [Xe2Fllf+ nHF ~ XeF + HL -F-Xe-L + HF (J) 2 2 and the halogen pentafluorides-are not 2[XeF5f + [(HF)nFr (q) and XeF2 + 2HL - L-Xe-L + 2HF (k) donor-acceptor compounds but simple molecular adducts in which the com­ .In all cases, the linkage to xenon is These same two ions are found in ponent molecules preserve their mol­ via oxygen (L = OTeF , OSeF , OS02F, the adducts. The [XeF5f ion has the 5 5 ecular identities. OPOF , ON0 , OCI0 , OIF 0, OCOCH square pyramidal form found in XeFs and 2 2 3 4 3 The enthalpies of ionisation of the or OCOCF ). Only those species con­ [Xe2Fllf is best described in terms of an 3 binary fluorides: taining the highly electronegative -OTeF5 ionic model, [XeF5fF [XeF5f, but with some covalent character in terms of and -OSeF5 units are stable under XeFs(g) - [XeF5r (g) + F-(g) ordinary conditions. The -OS02F and tlH=9.24eV (m) ([XeF5fXeFs) to account for the non­ -OIF 0 derivatives decompose readily linearity of the Xe-F-Xe bond. 4 XeF 4(g) - [XeF r (g) + F(g) and the remainder are explosive. 3 A relatively poor fluoride ion donor, tlH= 9.66eV (n) The high electronegativity of the xenon oxide tetrafluoride, forms adducts with only the stronger Lewis acids. -OTeF5 group means that Xe(OTeF5h is XeF2(g) - [XeFr (g) + F-(g) Thus, with SbF , [XeOF nSb F r and similar to XeF2 and an interesting reac­ tlH= 9.45eV (0) 5 3 2 ll tion of Xe(OTeF h is that with halogeno­ [XeOF3nSbFsr are obtained. The re­ 5 imply that complex formation might be alkanes such as CF =CFCI, CF =CCI and lated dioxide difluoride, Xe02F2, yields 2 2 2 progressively easier in the sequence CF =CFH, which results in the addition of [Xe02Ff[Sb2Fl1r and [Xe02FnSbFsL 2 XeF4 XeF XeF . The chemical -OTeF5 groups across the double bond. 21 < 2 < s but these are less stable. Reaction of evidence supports this. Xenon difluoride The development of substitution reac­ FXeOTeF5 with AsF5 gives the complex reacts wjth AsF , SbF , BiF , and a tions involving Xe(IV) and Xe(vI) com­ 5 5 5 [XeOTeF5nAsFsr and reactions of this range of pentafluor­ pounds has required the use of the complex give rise to a number of as yet I ides, to give adducts that are often reagent B(OTeF h, an obvious relative of uncharacterised xenon cations. 5 written [XeFnM F L [XeFnMFsr and BF . Reactions involving this molecule 2 ll Fluoride ion acceptor behaviour is 3 Xe F have given the only Xe(IV) compound [ 2 3nMFsr but in which the cation limited to XeFs. With fluorides units are linked to the counter ion by to date, Xe(OTeF )4, while Xe(VI) gives it yields M2XeFS (M =Cs, Rb, K and Na) 5 fluorine bridges; the V-shaped [Xe2F3f Xe(OTeF )s, OXe(OTeF )4, 02Xe(OTeF h and MXeF7 (M = Cs, Rb), which lose XeFs 5 5 5 ion can be regarded as two [XeFf units and, under certain conditions in solution, when heated. The thermal stability of linked by a fluorine bridge. The most the related partly substituted derivatives. these adducts increases with increasing ionic adducts are those derived from Like many of the Xe(II) compounds, most molecular weight and the low decom­ the strongest Lewis acids. The related of these are also unstable. position temperature of the com­ derivatives of MOF4 (M = W or Mo), which Because noble gas chemistry is mostly plex, which is below 100 °C, provides a are weaker Lewis acids, have lesS ionic associated with Xe-F and Xe-O bond convenient way of separating XeFs from character. formation the key to the successful bond­ XeF2 and XeF 4 with which no reactions ing of xenon to other atoms to yield stable One of the most interesting reactions occur. Solid state structures have not compounds lies in producing suitable of [XeFf is its reduction with some been determined, but the related com­ highly electronegative groups. Such· a compounds to give the bright green plexes derived from NOF and N02F have group is -N(S02Fh which has provided paramagnetic cation [Xe2r.2S This cation been prepared and [NO]i[XeFs]2- has the route to the first Xe-N bond:22 can also be obtained by oxidising xenon been shown to contain a slightly distorted with salts. Furthermore, square antiprismatic [XeFs]2-. Valence Bartlett has shown that the first noble shell electron pair repulsion theory sug­ gas compound, Xe +[PtFsL has the gests that a ninth coordination position in variable composition, Xe(PtFs)x (x =1-2). which a lone pair can be accommodated F-Xe-N(S02Fh (I) Whether this compound contains Xe+, should be present in this anion. However, [Xe2f or [Xe2Ff or can be described in there is no evidence for it, and this The solid, once formed, is reasonably terms of [XeFnPtFsr and [XeFnPt2Fllr is implies a rare violation of the theory in stable. It contains a linear F-Xe-N still uncertain. noble ga'S chemistry. group and a planar configuration about the nitrogen.22.23 Its reaction with AsF5 yields the solid, [(FS0 hN-Xe-F-Xe­ Table 1 . The cationic and anionic derivatives of binary fluorides and oxide fluorides 2 25 N(S02FhnAsFsf, containing a V-shaped ofxenon. cationic relative of [Xe2F3f. Although, in general, the noble gases Fluorides Cationic derivatives Anionic derivatives form bonds with only the most electro­ XeF [XeFr; [Xe2F3r negative elements, fluorine and oxygen, 2 the success in making an Xe-N bond in XeF. [XeF3f compounds has increased speculation as XeFe [XeF6r; [Xe 2F"r [XeF7L [XeFe]2- to whether an Xe-C bond is possible. A XeOF. [XeOF3r [XeOF6r: [Xe30 3F13r plasma-induced reaction of XeF vapour 2 Xe02F2 [Xe02Fr [Xe02F3r with CFj radicals may have provided the XeOF [XeOF r answer. The volat,ile, waxy, white solid 2 3 664 HEM I TRY IN BRITAIN J LY 1987

used as fluorinating agents in industrial processes and radon chemistry being Scheme 1. The reaction of XeOF4 with CsF. applied in areas of radon monitoring and atmospheric cleansing. Had it not been for Bartlett's astute chemical awarenesS, however, it is possible that we might still CsF + XeOF (excess) o DC, vacuum) 4 - XeOF4 cSF.xeoF4 1 be talking of the 'inert gases'. 20 °C, vacuum - XeOF4 Dr John H. Ho 110 way is a senior lecturer in the department of chemistry, Univer­ sity of Leicester, Leicester LET 7RH.

References 1. N. Bartlett and D. H. Lohmann, Proc. Chem. Soc., 1962, 115. 2. N. Bartlett and F. O. Sladky in Comprehen­ Stable white solids containing the cations. By analogy with xenon chem­ sive , A. F. Trotman­ nM + [Xe0 3 F]~ anion consist of chains of istry, the reaction with [02nSbF6f is Dickenson (ed ), p 214. Oxford: Pergamon, pseudo-octahedral xenon atoms linked probably: 1973. by angular bridges. These are deposited 3. N. Bartlett, Proc. Chem. Soc., 1962,218. Rn(g) + 2([02nSbF f}(s) ­ 4. H. H. Claassen, H. Selig and J . G. Maim, when solutions of stoichiometric 6 J. Am. Chem. Soc., 1962, 84,3593. amounts of Xe03 and HF are concen­ [RnFnSb2F11ns) + 202(g) (t) 5. R. Hoppe et aI, Angew Chem., 1962, 74, trated. The haloxenates, CsXe03CI and 903. 13 25 6. J . Slivnik et aI, Croat. Chem. Acta, 1962, CsXe03Br have also been obtained. Krypton compounds . 34, 253. Complexes of XeOF4 with CsF, RbF, Krypton chemistry is more limited than 7. P. R. Fields, L. Stein and M . H. Zirin, KF and with NOF have been known for that of xenon. Although small amounts of J. Am. Chem. Soc., 1962,84,4164. some time. The former are obtained by the violet free radical, KrF, have been 8. J . G. Maim, I. Sheft and C. L. Chernick, treating the alkali fluorides with an ex­ observed in the y-irradiation of KrF2' J. Am. Chem. Soc., 1963, 85, 110. 9. F. B. Dudley, G. Gard and G. H. Cady, cess of XeOF4 and pumping to constant most of the chemistry derives from the Inorg. Chem., 1963,2,228. weight under appropriate conditions (see difluoride. No oxides or oxide fluorides 10. E. E. W ea ver, B. W einstock and C. P. Scheme 1). The structure of CsF'XeOF4 have been isolated and ea rly reports on Knop, J. Am. Chem. Soc., 1963, 85, 111 . contains an [XeOF f anion of distorted the preparation of KrF4 have not been 11 . H. H. Hyman (ed), Noble-gas compounds. 5 Chicago and London: University of octahedral geometry with a lone pair confirmed. Chicago Press, 1963. of electrons apparently occupying an The volatile, colourless, crystalline di­ 12. J. Berkowitz and W . A. Chupka, Chem. octahedral face. The CsF'3XeOF4 adduct fluoride is prepared by subjecting mix­ Phys. Lelt., 1970, 7,447. contains the [F(XeOF bf anion in which tures of krypton and fluorine to an electric 13. J . H. Holloway, Noble-gas chemistry. 4 London: M ethuen, 1968; H. -G. Horn in the XeOF4 units, of close to square discharge at low temperatures or by Gmelins Handbuck der Anorganischen pyramidal shape, are linked via three irradiating liquid krypton- fluorine mix­ Chemie, Erganzungswerk zur 8 Auflage Xe---F bridging bonds to a common tures with ultraviolet light.28 The linear Band 1, Ede lgasverbindungen, Verlag fluorine which is out of the plane of the molecular structure of KrF2 is similar to Chemie, W ei nhei m, 1970; N. Bartlett and F. O. Sladky, Comprehensive three xenons. Trigonally bonded fluorine that of XeF2 but, because of its much inorganic chemistry, A. F. Trotman­ is rare but, where it does occur (ie (XeF6)6 lower bond energy, it is a very powerful Dickenso n (ed ), p 213. Oxford: Pergamon, and NaSb3FlO), the arrangements about it fluorinating agent capable of oxidising 1973. are closely similar. iodine to IF , xenon to XeF , low-valent 14. L. S. Bartell, J. Chem. Phys., 1967, 46, 7 6 4530; R. M . Gavin, Jr, and L. S. Bartell, uranium, neptunium and plutonium J. Chem. Phys., 1968, 48,2460; H. Kim, Radon compounds7.13.25.27 compounds to the hexafluorides,29 and H. H. Claassen and E. Pearson, Inorg. Radon chemistry has assumed increas­ AgF2 to AgF3.30 KrF2 is also rapidly Chem., 1968,7,616. 15. K. Seppelt, Acc. Chem. Res., 1979, 12, ing importance since · it emerged that decomposed by water. 211 . radon can be oxidised and complexed by The adducts of KrF2 are analogous to 16. R. D. Burbank and G. R. Jones, J. Am. certain liquid fluorides and fluoride salts. those of XeF2 and are limited to cationic Chem. Soc., 1974,96, 43. Tt. ;" offers the possibility of removing derivatives formed with fluoride ion 17. K. Seppelt and D. Lentz, in Progress in Kr F inorganic chemistry, S. J . Lippard (ed), radon from mine atmospheres and of acceptors. Species such as [ 2 3nMF6r vol 29, p 167. New York: W iley, 1982. developing radon-monitoring devices.27 (M =As, Sb, Ta), [KrFnMF6f (M =As, Sb, 18. R. Filler, Isr. J. Chem., 1978, 17, 71 . The high radioactivity of radon has V, Ta, Pt and Au) and [KrFnM2Fllf 19. M . McHughes et aI, Inorg. Chem., 1986, meant that the chemistry of radon has (M =Sb, Nb and Ta) are known in the so lid 25,426. state and in solution and, more recently, 20. J . Huston, Inorg. Chem., 1982, 21 ,685. been explored by tracer methods. When 21 . C. J . Schack and K. O. Christe, J. Fluorine microcurie amounts of radon are heated compounds such as [KrFf[MOF5f (M = Chem., 1985, 27, 53. with fluorine a difluoride appears to be Cr, 31 Mo and W 32 ) have been prepared 22. R. D. LeBlond and D. D. DesMarteau, formed. The same species is also gen­ and characterised. The spectra of the J. Chem. Soc., Chem. Commun., 1974, 555. erated spontaneously when millicurie or ions are related to those of the xenon 23. J . F. Sawyer, G. J . Schrobilgen and S. J . larger amounts of radon are mixed with analogues but additional peaks in the Sutherland, Inorg. Chem., 1982, 21, gaseous or liquid fluorine in a small Raman for [Kr2F3f have been interpreted 4064. volume; the latter reaction is brought in terms of an unsymmetrical (rather 24. L. J . Turbini, R. E. Aikman and R. J . Lagow, J. Am. Chem. Soc., 1979, 101 , about by the intense a-radiation. Radon than symmetrical) cation. The oxidising 5833. dissolves in halogen fluorides as an in­ power of [KrFt and [Kr2F3f is stronger 25. H. Selig and J . H. Holloway, in Topics in volatile species. Electromigration studies than that of KrF2, capable of oxidising IF5 current chemistry, F. L. Boschke (ed), vol have shown that in pure solvents the to [IF6f, O to [02f, Xe to [XeF5f, BrF5 to 124, p 33. Berlin: Springer- Verlag, 1984. 2 26. L. Stein et aI, J. Chem. Soc., Chem. radon is present as a cation and ionis­ [BrF6f and Au to [AuF6f and [AuF5]. Commun., 1978,502. ations involving the following equilibria Noble gas chemistry is now well estab­ 27. L. Stein, Radiochim. Acta, 1983,32,163. have been suggested: lished. Still inextricably linked with 28. J . Slivnik et aI, J. Fluorine Chem., 1975, fluorine chemistry, it has provided a 5,273. RnF2 ~ [RnFf + F- (r) 29. L. B. Asprey, P. G. Eller and K. A. Kinkead, means to accomplish novel fluorinations, [RnFf ~ Rn 2+ + F- (s) Inorg. Chem., 1986,25,670. of introducing us to new classes of 30. R. Bougon et aI, Inorg. Chem., 1984, 23, The existence of the [Rn- Ff cation compounds and new high oxidation state 3667. seems certain in view of the spontaneous species. In the next few years it promises 31 . J. H. Hollow ay and G. J . Schrobilgen, Inorg. Chem., 1981,20,3363. reaction of radon with a variety of solids to take us into the realms of applied 32. K. O. Christie, W . W . Wilson and R. A. such as [02nSbF6f containing oxidising chemistry, with noble gas fluorides being Bougon,lnorg. Chem., 1986,25,2163.