The Chemistry of Xenon(IV) - ACS Publications - American

He was elected a Fellow of the Royal Society of Canada in 1999. Professor Schrobilgen has served on the Executive Committee of the Division of Fluorin...
100 downloads 0 Views 2MB Size
Download PDF
Review pubs.acs.org/CR

The Chemistry of Xenon(IV) Jamie Haner and Gary J. Schrobilgen* Department of Chemistry, McMaster University, Hamilton, Ontario L8S 4M1, Canada 5.4. Xenon Dioxide, XeO2 5.4.1. Synthesis and Raman Spectra of Xe16/18O2 5.4.2. Computational Studies of the Solid-State Structure of XeO2 5.4.3. Implications of XeO2 on Earth’s Missing Xenon 6. Polyatomic Ligand Derivatives of Xenon(IV) 6.1. Xenon(IV) Derivatives of the F5TeO-Group 6.1.1. Syntheses of F4−nXe(OTeF5)n (n = 1−4) and Characterization by 19F and 129Xe NMR Spectroscopy 6.1.2. Characterization of Xe(OTeF5)4 by Raman Spectroscopy and X-ray Diffraction 6.1.3. 19F and 129Xe NMR Studies of [FnXe(OTeF5)3−n]+ (n = 0−2) 19 6.2. F and 129Xe NMR Studies of F3XeOIOF4 6.3. Possible Existence of Xe(OCOCF3)4 6.4. Salts of the [C6F5XeF2]+ Cation 6.4.1. Synthesis of [C6F5XeF2][BF4] 6.4.2. Multi-NMR Studies of [C6F5XeF2]+ and C6F5IF2 6.4.3. Structures of [C6F5XeF2][BF4], [C6F5XeF2][BF4]·1.5CH3CN, and [C6F5XeF2][BF4]·2HF 6.4.4. Bonding in the [C6F5XeF2]+ Cation 6.4.5. Raman Spectra of [C6F5XeF2][BF4] and [C6F5XeF2][BF4]·2HF 7. Conclusions and Outlook Author Information Corresponding Author Notes Biographies Acknowledgments Dedication References

CONTENTS 1. Introduction 2. Xenon Tetrafluoride, XeF4 2.1. Synthesis of XeF4 2.2. Structural Characterization and Physical Properties of XeF4 2.3. Thermochemical Properties of XeF4 2.4. Mö ssbauer Spectra of XeF4 and XeCl4 2.5. Molecular Addition Compounds of XeF4 2.5.1. XeF4·XeF2 2.5.2. ([XeF5][CrF5])4·XeF4 3. Fluoride Ion Donor Properties of XeF4 3.1. [Mg(XeF2)(XeF4)][AsF6]2 3.2. Salts of the [XeF3]+ Cation 3.2.1. Syntheses of [XeF3]+ Salts and Solution Studies 3.2.2. Crystal Structures of [XeF3]+ Salts 3.2.3. Raman Spectra of [XeF3]+ Salts 3.2.4. Theoretical Studies of the [XeF3]+ Cation 4. Fluoride Ion Acceptor Properties of XeF4; The [XeF5]− Anion 4.1. Syntheses of [XeF5]− Salts 4.2. X-ray Crystal Structure of [N(CH3)4][XeF5] 4.3. 19F and 129Xe NMR Spectra of [XeF5]− 4.4. Vibrational Spectrum of [XeF5]− 4.5. Bonding in [XeF5]− 5. The Oxide Fluorides and Oxide of Xenon(IV) 5.1. Syntheses and Properties of XeOF 2 , F2OXeNCCH3, and XeOF2·nHF 5.1.1. 19F, 17O, and 129Xe NMR Spectra of F2OXeNCCH3 5.1.2. Raman Spectra of XeOF2, F2OXeNCCH3, and XeOF2·nHF 5.1.3. X-ray Crystal Structure of F2OXeNCCH3 5.2. The [XeOF3]− Anion 5.2.1. Raman Spectrum and Computational Studies of [XeOF3]− 5.3. Xenon(IV) and Mixed Oxidation State Xenon(II)/Xenon(IV) Hydroxy Fluoride Cations 5.3.1. Syntheses of [H(OXeF 2) n ][AsF6 ] and [FXeII(OXeIVF2)n][AsF6] (n = 1, 2) 5.3.2. Structures of [HOXeF 2 ] + , [HOXe(F) 2 OXeF 2 ] + , [FXe I I OXe I V F 2 ] + , and [FXeIIOXeIV(F)2OXeIVF2]+ © XXXX American Chemical Society

A B B E F G G G H H H I J K L M N N O O P P Q Q

Z AA AB AB AC AC

AC AD AE AE AF AF AF AG AH AI AJ AJ AK AK AK AK AL AL AL

1. INTRODUCTION Although many xenon compounds in the II and VI oxidation states have been synthesized, the chemistry of Xe(IV) and its compounds is less developed.1,2 Several factors have contributed to the relative scarcity of Xe(IV) compounds: their strong oxidant properties, the endothermic natures of the oxide and oxide fluoride derivatives, and their tendencies to undergo redox eliminations and disproportionations.3−6 Prior to 2000, the number of structurally characterized Xe(IV) species were limited to [XeF3]+,7−9 [XeF5]−,10 XeF4,11−13

Q R U U U V V

Special Issue: 2015 Fluorine Chemistry Received: August 5, 2014

X A

DOI: 10.1021/cr500427p Chem. Rev. XXXX, XXX, XXX−XXX

Chemical Reviews

Review

Xe(OTeF 5 ) 4 , 1 2 , 1 4 , 1 5 Xe(OTeF 5 ) 4 − n F n (n = 0−3), 1 2 [C6F5XeF2]+,16,17 F3XeOIOF4,18 [FnXe(OTeF5)3−n]+ (n = 0− 2),19 and to preliminary reports of XeOF220−22 and [XeOF3]−.22 Xenon tetrafluoride is the first noble-gas fluoride to be synthesized23,24 and structurally characterized (see section 2.2) and is the synthetic precursor for all Xe(IV) compounds; for this reason, this Review initially focuses on the synthetic methodologies that led to XeF4 and its structural characterization. Much of the early work relating to the synthesis of XeF4, its characterization, and properties has been surveyed in a general review of noble-gas chemistry by Neil Bartlett, the discoverer of noble-gas reactivity, and F. O. Sladky,25 and covers the period 1962−1970. Several other more specialized reviews have been published on noble-gas chemistry,26−43 some of which include the chemistry and properties of Xe(IV) species. The most recent comprehensive review of noble-gas chemistry spans 2000− 2012.44 The present Review provides a comprehensive survey of developments pertaining to the syntheses, properties, structures, and bonding of Xe(IV) compounds from 1962 to 2014.

phase infrared spectrum, the authors proposed a square-planar or tetrahedral geometry. The room-temperature stable colorless crystals of XeF4 were sublimed into a silica spectrophotometer cell and became a standard exhibit used to convince local and visiting chemists that xenon compounds are real (Figure 1).24

2. XENON TETRAFLUORIDE, XeF4 2.1. Synthesis of XeF4

The first noble-gas compound was discovered in 1962 by Neil Bartlett.23 The compound, then formulated as “Xe+PtF6−”, was formed in the reaction of xenon gas with PtF6 vapor. This discovery prompted researchers at Argonne National Laboratory to repeat this experiment. They also demonstrated that RuF6 reacts with xenon to form “Xe+RuF6−”. Because RuF5 is a stable compound, it was postulated that RuF6 and PtF6 could serve as F• sources (eq 1) in the oxidative fluorination of xenon. Thus, the reactivity of xenon with elemental fluorine at high temperatures and pressures was investigated at an early stage in the history of noble-gas chemistry.24 These experiments soon led to the discovery of the first binary fluoride of a noble gas and the first Xe(IV) compound, XeF4.45,46

MF6 → MF5 + F•

Figure 1. Crystals of XeF4 in a 22 mm diameter silica spectrophotometer cell. Reprinted with permission from ref 24. Copyright 1963 University of Chicago. This photograph also appears on the cover of Science, 1962, 138, 69.

Although XeF4 was the first binary fluoride of xenon to be discovered, it is generally regarded to be the most difficult to synthesize in high purity among the series of binary xenon fluorides, XeF2, XeF4, and XeF6, when high-temperature, highpressure synthetic methods are applied. The synthesis of highpurity XeF4 under thermal conditions involves the heating of one part Xe (0.361 g; 0.00275 mol) with five parts F2 (0.524 g, 0.0138 mol) in a 130 mL welded nickel reactor at 400 °C for 1 h followed by rapidly cooling the reactor in a water bath and removal of the excess fluorine at −196 °C (Figure 2).48 If reaction conditions are not stringently controlled, product mixtures may result from equilibria 3−5.

(1)

In the original experiment, a 9.6:1.06 molar ratio of F2:Xe was condensed into a nickel reaction vessel.45 The mixture was heated to 400 °C for 1 h and then cooled to −196 °C, followed by removal of unreacted F2 under dynamic vacuum. The vessel was then warmed to −100 °C to transfer any xenon that may not have reacted into a U-trap held at −196 °C. However, when researchers opened the valve of the reaction vessel at −100 °C, no xenon pressure was detected. Although the intention was to react xenon with fluorine, the experimentalists were nonetheless surprised by the results. That afternoon, they wrote in their notebook, “Go home somewhat puzzled. Where is the Xe?”24 Subsequent quantitative experiments using the material remaining in the reaction vessel established the relative amounts of fluorine and xenon that were consumed, thus providing the formulation, XeF4, and the reaction stoichiometry (eq 2).

Xe + 2F2 → XeF4

Xe + F2 ⇌ XeF2

(3)

XeF2 + F2 ⇌ XeF4

(4)

XeF4 + F2 ⇌ XeF6

(5)

Thus, the choices of initial Xe/F2 composition, reaction temperature, and reaction pressure are critical to the purities of binary xenon fluorides synthesized by this method (Figure 3).28 Gas-phase equilibrium constants have been determined for eqs 3−5 (Table 1) from Xe and F2 mixtures at various temperatures and pressures.50 The data in Table 1 and Figure 3 show that at low temperatures and high F2 pressures, the formation of XeF6 is favored, whereas at low F2 pressures and high temperatures, the formation of XeF2 is favored. Accordingly, it is not possible to prepare high-purity XeF4 by the thermal method unless it is further purified. The relatively low volatility of XeF4 allows for the more volatile XeF2 and XeF6 impurities to be separated by

(2)

The molecular formula of XeF4 was also established by reaction of XeF4 with H2 at 400 °C in a nickel reaction vessel.45 The resulting reaction products, Xe and HF, were analyzed gravimetrically and by fluoride analysis, respectively. Xenon tetrafluoride was reported to have a melting point of ca. 114 °C26 and a vapor pressure of 3 Torr at room temperature. Large colorless crystals readily formed under static vacuum in a glass vessel within several hours. On the basis of a preliminary gasB

DOI: 10.1021/cr500427p Chem. Rev. XXXX, XXX, XXX−XXX

Chemical Reviews

Review

A number of alternative procedures for the synthesis of XeF4 have been reported. A flow method was reported to yield XeF4 by passage of a Xe/F2 mixture through a 4-ft. × 1/2-in. o.d. length of nickel tubing packed with nickel sheet that was heated to a dull red color.54 Xenon tetrafluoride was collected in 30−50% yield from the effluent gas in a trap held at −72 °C. A closed circulation system was also reported to form XeF4 in 97% yield when a 3:1 mixture of F2:Xe was circulated through a column packed with pieces of nickel sheet that were heated to 560 °C. The XeF4 was collected in a trap at 0 °C.55 In flow preparations, the purities of the XeF4 were not provided. It was also shown that XeF4 is formed when Xe/F2 mixtures are circulated through an electric glow discharge at −78 °C.56 A report of the synthesis of KrF457,58 by circulation of Kr/F2 (1:2) mixtures at −214 to −212 K through an electric glow discharge appeared shortly after the discovery of XeF4. Independent attempts to reproduce the synthesis of KrF4 by this method confirmed that a krypton fluoride was indeed formed;59 however, it was shown to be KrF2. Krypton(II) is the highest known oxidation state of krypton,58 and calculations verify the instabilities of higher oxidation state krypton fluorides toward dissociation to KrF2 and F2.60 Irradiation of a mixture of xenon and fluorine at ambient temperatures with γ-rays from a 60Co source was shown to form XeF4 along with a significant amount of XeF2.61,62 When the reactor was cooled to ca. −35 °C, only XeF2 formed, which was ascribed to the condensation of XeF2 in cooler regions of the reactor such that its vapor pressure and further gas-phase reaction to form XeF4 were suppressed. In a related study, crystalline XeF4 exposed to 60Co γ-radiation at 77 K turned blue and was shown by ESR spectroscopy to give rise to the XeF• radical.63,64 The formation of XeF4 by UV irradiation of mixtures of F2 and Xe gases in a water-cooled reactor with a Hg lamp (365−366 nm) has been thoroughly investigated.29,65,66 Studies of the photochemical reaction of Xe and F2 showed that the reaction is accelerated by 4-fold when unpurified F2 gas containing ca. 1 mol % HF was used.66 A 3-fold increase in reaction rate was achieved in the presence of solid NiF2. The presence of O2 (14 mol %) increased the rate of the photochemical reaction by approximately 2-fold, which is

Figure 2. Apparatus for the preparation and purification of xenon tetrafluoride. (A) 30 000-psi valve, Monel body, Teflon packing. (B) 30 000-psi Monel tee. (C) 30 000-psi Monel cross. (D) Welded nickel U-tube. (E) Welded Monel Bourdon gage (0−1000 mm). (F) 130 mL welded nickel reaction vessel (1 7/8-in. o.d. × 1/32-in. wall × 3-in. long). (G) 1500 mL welded nickel storage and measuring vessel (4-in. o.d. × 1/8-in. wall × 9 1/2-in. long). (H) 85 mL welded nickel storage and measuring vessel (1 1/8-in. o.d. × 1/16-in. wall × 6 1/4-in. long). (J) Brass valve: Hoke type A413. (K) Soda-line trap. (L) Monel valve: Hoke type 413. Reprinted with permission from ref 48. Copyright 1966 John Wiley & Sons.

fractional sublimation. Although XeF2 is more difficult to remove because its vapor pressure is similar to that of XeF4, repeated removal of equilibrated vapor phase fractions under dynamic vacuum has proven successful for the purification of XeF4 samples containing ca. 1% XeF2.51 Xenon hexafluoride impurity can also be removed by complexation with excess NaF (Na2XeF8),52 followed by sublimation of the less Lewis acidic XeF4 from the mixture. Differences in fluoride ion donor strengths can also be exploited to purify crude XeF4 mixtures (section 3.2). The equilibrium constants for the formation of XeF4 from its elements have also been determined from thermodynamic data51 and are in reasonably good agreement with the experimental values (Table 2).50 Because of the experimental difficulties encountered in measuring the equilibrium constants of the Xe/F2 system, the calculated equilibrium constants are likely to be more accurate. The temperature dependency of the reaction of Xe and F2 was also demonstrated by reacting 10:1 molar mixtures of F2:Xe at 120, 150, and 200 °C to form primarily XeF2, XeF4, and XeF6, respectively.53

Figure 3. (a) Equilibrium pressures of xenon fluorides as a function of temperature. Initial conditions: 125 mmol of Xe, 275 mmol of F2 per 1000 mL. (b) Equilibrium pressures of xenon fluorides as a function of temperature. Initial conditions: 125 mmol of Xe, 1225 mmol of F2 per 1000 mL. Reprinted with permission from ref 49. Copyright 1967 Academic. C

DOI: 10.1021/cr500427p Chem. Rev. XXXX, XXX, XXX−XXX

Chemical Reviews

Review

Table 1. Experimental Equilibrium Constantsa for the Xe/F2 System50 temp (K) 298.15

523.15

[1.23 × 10 ] [1.37 × 1011] [8.2 × 105] 13

K1(XeF2) K2(XeF4) K3(XeF6)

573.15

[8.80 × 10 ] 1.43 × 103 0.944

[1.02 × 10 ] 1.55 × 102 0.211

4

4

623.15

673.15

773.15

[1670] 27.2 0.0558

[360] 4.86 0.0182

29.8 0.50 0.0033

a

K1 = (XeF2)/(Xe)(F2); K2 = (XeF4)/(XeF2)(F2); K3 = (XeF6)/(XeF4)(F2). Square brackets denote calculated values. Reproduced with permission from ref 50. Copyright 1966 American Chemical Society.

Table 2. Equilibrium Constanta for the Formation of XeF4 from Heat Capacity Measurements51 temp (K)

ΔGf°/T (J K−1 mol−1)

K4 (atm−2), calcd51

289.5 400 500 600 700 800 900 1000

−440.59 ± 3.1 −264.81 ± 2.3 −162.39 ± 1.9 −94.47 ± 1.6 −46.23 ± 1.4 −10.26 ± 1.2 17.57 ± 1.1 39.70 ± 1.0

1.03 × 1023 6.79 × 1013 3.04 × 108 8.61 × 104 2.60 × 102 3.43 0.121 8.44 × 10−3

K4 (atm−2), exptl50

1.05 × 109 2.03 × 105 4.66 × 102 5.01

a

K4 = (XeF4)/(Xe)(F2)2. Reproduced with permission from ref 51. Copyright 1972 American Institute of Physics.

hν

(6)

O2 + F• → O2 F•

(7)

O2 F• + F2 → O2 F2 + F•

(8)

O2 F2 → O2 F• + F•

(9)

•

Xe + F → XeF

(11)

XeF• + F2 → XeF2 + F•

(12)

(14)

XeF3• + F2 → XeF4 + F•

(15)

XeF3• + F• → XeF4

(16)

XeF4 + F• → XeF5•

(17)

XeF5• + F2 → XeF6 + F•

(18)

XeF5• + F• → XeF6

(19)

XeF4 + F2 HooooI XeF6 14 ° C

(20)

350 ° C

XeF4 + F2 ⎯⎯⎯⎯⎯→ XeF6 NaF

(21)

The potent low-temperature fluorinator, O2F2, was shown to be capable of oxidizing xenon to XeF4 (eq 16)67 when a 10-fold excess of O2F2 was combined with Xe at −196 °C. The reaction mixture was allowed to slowly warm from −155 to −120 °C over a 2−4 h period. The reaction mixture was subsequently warmed to −78 °C and held at this temperature overnight to promote decomposition of the remaining O2F2 to O2 and F2. In addition to O2 and F2, a mixture of XeF2 and XeF4 was initially obtained. A second treatment of the XeF2/XeF4 mixture with O2F2 resulted in a quantitative yield of XeF4. Xenon hexafluoride does not form, even in the presence of large excesses of O2F2, indicating that O2F2 is not potent enough to oxidatively fluorinate Xe(IV) to Xe(VI). The oxidation of Xe to Xe(IV) by O2F2 is in disagreement with an earlier report that only XeF2 was formed upon reaction of Xe with O2F2 (eq 22).73 Xe + 2O2 F2 → XeF4 + 2O2

(10) •

XeF2 + F• → XeF3•

hν

In contrast, the reaction rate for the high-pressure, thermal synthesis of XeF4 was shown to be unaffected by the presence of oxygen when reactions of 1:100 and 4.6:1 molar ratios of O2:F2 at 573.15 K were compared.50 Room-temperature UV irradiation of XeF2 in the presence of partially soluble F2 gas in aHF has also been used to prepare XeF4.68 The nature of F2 activation by HF or NiF2 is not fully understood, but similar reports such as the nonirradiative room-temperature synthesis of XeF2 in aHF support F2 activation by HF.69 Catalysis by NiF2 has also been reported in the thermal reaction of fluorine and xenon mixtures.70 In all of the aforementioned XeF4 syntheses, homolytic cleavage of F2 into F• radicals is likely the first step (eq 10). Fluorine radicals combine with Xe to form XeF• radicals (eq 11) followed by their reaction with F2 or F• atoms to form XeF2 (eqs 12, 13). Subsequent reaction of XeF2 with F• atoms (eq 14) followed by reaction with F2 or F• atoms (eqs 15, 16) would yield XeF4. Under high temperature and pressure conditions,49 XeF4 can be further fluorinated to XeF6 (eqs 17−19). The formation of XeF2 is reported to be ca. 20 times faster than the formation of XeF4.69 F2 → 2F•

(13)

In the photochemical synthesis of XeF4 from Xe and F2, only very small amounts of XeF6 are produced.71 The photochemical equilibrium shown in eq 20 was examined by UV photolysis of XeF6 at 14 °C in a copper reactor outfitted with sapphire windows. Xenon hexafluoride was shown to rapidly dissociate to XeF4 and F2, with only a small equilibrium concentration of XeF6 remaining. The equilibrium was verified by the addition of AsF5 to the reactor containing XeF4, F2, and traces of XeF6. Ultraviolet irradiation resulted in complete F2 consumption with conversion of XeF6 and XeF4 to [XeF5][AsF6]. Photolytic conversion of XeF6 to XeF4 was suppressed when excess F2 was added to the photochemical reactor.71 A thermal synthesis of XeF4 has also been developed on the basis of the decomposition of XeF6 in the presence of excess NaF at 350 °C (eq 21).29,72

attributed to an increased concentration of fluorine radicals produced by radical chain reactions 6−9. This explanation is supported by a later report on the reactivity of O2F2 with Xe (see below).67

F2 → 2F•

XeF• + F• → XeF2

(22)

Xenon tetrafluoride has also been synthesized by thermal decomposition of Xe(II) fluoride salts. Pyrolysis of [XeF][PtF6] or XePd2F10 at 150 °C yielded XeF4 and XePt2F10 or Pd2F6, D

DOI: 10.1021/cr500427p Chem. Rev. XXXX, XXX, XXX−XXX

Chemical Reviews

Review

respectively (eqs 23 and 24).74 In these cases, Xe(II) is oxidized to Xe(IV), and the Pt(V) and Pd(IV) anions are reduced to Pt(IV) in XePt2F10 and Pd(II)/Pd(IV) in Pd2F6, respectively. The isostructural salts, XePt2F10 and XePd2F10, have been shown by Raman spectroscopy to contain [XeF]+ units and most likely the polymeric [M2F9]nn− anions.74 2[XeF][PtV F6] → XePtIV 2F10 + XeF4

(23)

XePd 2 IV F10 → Pd 2 II/IV F6 + XeF4

(24)

study, but solution values range from 331.7 to 325 Hz.86 The shielding and 1J(129Xe−19F) couplings of XeF2, XeF4, and XeF6 were calculated87 with nonrelativistic Hartree−Fock (HF) and relativistic Dirac−Hartree−Fock (DHF) methods.87 The shielding constants calculated for XeF4 by the DHF method (σ(Xe) = 1605.4 ppm, σ(F) = 275.7 ppm) gave the best agreement with the experimental values86 (σ(Xe) = 1471 ppm, σ(F) = 210.5, 208.7 ppm), but the calculated shielding constant of fluorine was much further from the experimental values than that of xenon. The calculated 1J(129Xe−19F) coupling (−4058.1 Hz) is in very good agreement with the experimental values (4060 and 3849 Hz).86 The X-ray crystal structure of XeF4 (P21/n) revealed a molecular solid in which the XeF4 molecules are square-planar within ±3σ (D4h point symmetry) with Xe−F bond lengths of 1.92(3) Å and F−Xe−F bond angles of 86(3)°.75 Significantly improved structural parameters were obtained from a singlecrystal neutron diffraction study (Figure 4), which yielded Xe−F bond lengths of 1.951(2) and 1.954(2) Å and F−Xe−F bond angles of 90.0(1)°.13

2.2. Structural Characterization and Physical Properties of XeF4

Shortly after the initial report of XeF4,45 the proposed squareplanar structure of XeF4 was unambiguously established by the gamut of physical methods that were then available including single-crystal X-ray75 and neutron diffraction13 and infrared,11 Raman,11 and NMR spectroscopies.76−78 The mass spectrum of XeF4 was measured and a minor peak corresponding to the [XeF4]+ parent cation was detected, as well as major peaks for [XeF3]+, [XeF2]+, [XeF]+, and [Xe]+ in increasing intensities.79,80 The magnetic susceptibility of XeF4, determined by Gouy’s method at room temperature, is χM = −50.6 × 10−6 g−1 mol−1 and becomes temperature dependent below 240 K.81 Early on, several groups reported room- and low-temperature 19 F and 129Xe (I = 1/2, 26.44% natural abundance) NMR parameters for XeF4 in aHF that included 19F and 129Xe NMR chemical shifts and the 1J(129Xe−19F) coupling constant.76,82−84 The 19F chemical shift of solid XeF4 was also measured.84,85 The NMR parameters for XeF4 have since been remeasured in various solvents and referenced to their IUPAC standards (Table 3). The 19 F and 129Xe chemical shifts and the 1J(129Xe−19F) coupling are discussed in a comprehensive review on multi-NMR spectroscopy of noble-gas compounds.2

Figure 4. Neutron diffraction crystal structure of XeF4 (P21/n). Thermal ellipsoids are shown at the 50% probability level. Redrawn with permission from ref 13. Copyright 1963, American Association for the Advancement of Science.

Table 3. 19F and 129Xe NMR Parameters for XeF4 solvent

T (°C)

δ(129Xe), ppm

δ(19F), ppm

CCl3F CCl3F BrF5 CH3CN CH3CN gas gas

24 24 25 −42 24 94 120

202.9 166.1 253 335.3 316.9 163 165

−15.66

−20.1 −18.7 −17.8 −16

1

J(129Xe−19F)

refa

3817 3801 3823 3913 3895 4060 3849

12 88 89 90 10 86 86

The gas-phase Raman and infrared spectra of XeF4 are comprised of six fundamental vibrational bands and were assigned under D4h symmetry.11 The bands correspond to nine vibrational modes belonging to the irreducible vibrational representations, Γvib = A1g (R) + B1g (R) + B2g (R) + A2u (IR) + B2u (ia) + 2Eu (IR). Descriptions of the normal modes of vibration my be found in ref 91. The XeF4 molecule is centrosymmetric and therefore mutual exclusion applies, with no overlap of infrared- and Raman-active bands. Originally, the B2u fundamental was erroneously assigned to B1u symmetry.11 This error was subsequently propagated through nearly every publication pertaining to the vibrational spectrum of XeF492−95 as well as the vibrational spectrum of the isoelectronic [IF4]− anion, which was assigned by direct comparison with the vibrational spectrum and assignments of XeF4.96 The alleged ν4(B1u) fundamental at 221 cm−1 was determined from a very weak band at 442 cm−1 in the Raman spectrum of solid XeF4, which was attributed to the fundamental’s first overtone. In a later study, an extremely weak band at 433 cm−1 in the gas-phase Raman spectrum was assigned to the overtone of a fundamental at 216 cm−1 (ν4(B1u)). In both experimental spectra, the signalto-noise levels of the presumed overtones are extremely low, and therefore these and several other reported assignments (see Table 4 footnotes) should be regarded as tentative.

Samples were referenced externally at 30 °C with respect to the neat liquid references: XeOF4 (129Xe) and CCl3F (19F). A positive chemical shift denotes a resonance occurring to high frequency of the reference compound. a

The 19F and 129Xe chemical shifts and the 1J(129Xe−19F) coupling constants of gaseous XeF2, XeF4, and XeF6 were also determined to obtain their NMR parameters in the absence of solvent−solute interactions and bulk susceptibility effects.86 Of the series of gaseous xenon fluorides, XeF2 (δ(129Xe) = −2074 ppm, δ(19F) = −183 ppm), XeF4 (δ(129Xe) = 163 to 165 ppm, δ(19F) = −17.8 to −16 ppm), and XeF6 (δ(129Xe) = −35 ppm, δ(19F) = 126.4 ppm), XeF4 has the most deshielded 129Xe nucleus but a δ(19F) chemical shift that is intermediate with respect to those of XeF2 and XeF6. The 1J(129Xe−19F) coupling of XeF4 (4060 Hz) is less than that of XeF2 (5627 Hz). No 1 129 J( Xe−19F) coupling was observed for gaseous XeF6 in this E

DOI: 10.1021/cr500427p Chem. Rev. XXXX, XXX, XXX−XXX

Chemical Reviews

Review

stretching force constants of XeF2 (2.694 mdyn Å−1), XeF4 (2.829 mdyn Å−1), and XeOF4 (2.964 mdyn Å−1) increase with increasing xenon oxidation state (electronegativity). Stretching force constants were also calculated using general valence force fields for the series [IF4]− (2.221 mdyn Å−1), [BrF4]− (2.227 mdyn Å−1), [ClF4]− (2.13 mdyn Å−1), and XeF4 (3.055 mdyn Å−1). The higher Xe−F stretching force constant of XeF4 relative to those of the isovalent [XF4]− (X = Cl, Br, I) anions reflects the greater electronegativity of Xe(IV) and covalent character of the Xe(IV)−F bond.96

Table 4. Infrared and Raman Spectra of XeF4a,b

2.3. Thermochemical Properties of XeF4

Unlike the oxides of Xe(IV) and Xe(VI), XeF4 is not prone to disproportionation. The disproportionations of XeF4 into XeF2 and XeF6 (eq 25, ΔH° = 5 kcal mol−1, ΔS° ≈ 0 kcal K−1 mol−1) or into Xe and XeF6 (eq 26, ΔH° = 9 kcal mol−1, ΔS° ≈ −9 kcal K−1 mol−1) are unfavorable under standard conditions.25

Frequencies are given in cm−1. bValues in parentheses denote relative Raman intensities. The abbreviations denote a shoulder (sh), broad (br), not observed (n.o.), very strong (vs), and medium (m). cFrom reference 92. Combination bands at 1093(br) (ν1/ν4 + ν1/ν4), 1050(br) (ν1/ν4 + ν2/ν1), 1011(br) (ν2/ν5 + ν2+ν5), 599(0.3) [2 × 299], 588(0.2) [2 × 294], 448(0.2), 325(b), and 283(br) (ν1 + 60/74) as well as lattice modes at 101(6), 80(sh), 74(14), and 60(10) were observed in the Raman spectrum. dFrom reference 11. Combination bands at 1136, 1105, and 302 were reported for the gas-phase infrared spectrum of XeF4. A band at 123 cm−1 was initially assigned to the ν7(Eu) mode but was later shown to be an impurity.50,97 A very weak overtone at 442 (vw) [2ν5] was reported for the gas-phase Raman spectrum which was attributed to a fundamental at 221 cm−1. The Eu band is split into bands at 591 and 581 cm−1. eFrom ref 94. Very weak overtones at 433(vw) [2ν5] and 322(vw) [2ν7] were reported for the gas-phase Raman spectrum which were attributed to fundamentals at 216 and 161 cm−1, respectively. fFrom reference 95. Calculated values at the CCSD(T)/aug-cc-pVDZ level. gThe gas-phase assignments, B2g(ν4), A2u(ν3), B1g(ν2), and B2u(ν5) correspond to values that were initially assigned as B 2g (ν 5 ), A 2u (ν 2 ), B 1g (ν 3 ), and B 1u (ν 4 ), respectively.11 hFrom reference 93; isotopically pure 136XeF4. The Eu band is split into two bands at ca. 589 and 579 cm−1 in the gas-phase infrared spectrum which arise from rotational fine structure and are attributed to the ΔJ = +1 and −1 lines. a

2XeF4(g) → XeF2(g) + XeF6(g)

(25)

3XeF4(g) → Xe(g) + 2XeF6(g)

(26)

39

In a recent review, calculations of the energies for disproportionation of XeF4 were presented for equations 27− 29 (B3LYP; SDD pseudopotential and concomitant 2ζ basis set for Xe, and the 6-311++G** basis set for F). The disproportionation of XeF4 to [XeF3]+ and [XeF5]− has a barrier of 161 kcal mol−1 (eq 27), while the loss of fluorine from XeF4 has a barrier of 16 kcal mol −1 (eq 28). The redox disproportionation of XeF4 into XeF2 and XeF6 has a barrier of only ca. 5 kcal mol−1 (eq 29). 2XeF4 → [Xe IV F3][Xe IV F5]

(27)

XeF4 → Xe IIF2 + F2

(28)

2XeF4 → Xe IIF2 + Xe VIF6

(29)

Calculations of the unit cell of XeF4 were performed at 200 GPa.39 The compression of crystalline XeF4 showed the appearance of eight new Xe---F contacts (2.297 Å), and the compression of the Xe−F bonds from 1.914 and 1.900 to 1.894 Å. It appears that the Xe−F bonds in XeF4 are extremely incompressable, which was attributed to the repulsive effect of the xenon and fluorine lone pairs. It was suggested that reactions 27−29 may be favorable at such high pressures. The atomization energy at 0 K and heats of formation at 0 and 298 K were determined for XeF4 using high-level CCSD(T) quantum-chemical calculations.95 In the series of binary fluorides of xenon, XeF2, XeF4, and XeF6, the heats of formation increase with increasing oxidation state (Table 5). The calculated total dissociation energies for XeF4 (−42.5 ± 1.0 kcal mol−1, 0 K; −43.5 ± 1.0 kcal mol−1, 298 K) are less negative than the experimental values (−50.2 kcal mol−1, equilibrium 0 K;50 −57.7 ± 2 kcal mol−1, photoionization (PI)98) derived from the

High-level calculations (CCSD(T)) of the vibrational frequencies of XeF4 are generally in good agreement with the experimental values.60 The experimental vibrational frequencies and assignments obtained for XeF4 and calculated frequencies are given in Table 4. Band splittings observed in the solid-state Raman spectra have been attributed to vibrational coupling (factor-group splitting) between XeF4 molecules within the crystallographic unit cell.92 An infrared spectroscopic study of 136Xe-enriched XeF4 gave an effective resolution of 0.002−0.004 cm−1 and allowed for a rotational analysis of the ν2 (A2u) band and a very precise determination of the gas-phase Xe−F bond length (1.93487(3) Å).93 The stretching force constants for XeF4 were determined from its gas-phase infrared spectrum. Urey−Bradley types of interactions between nonbonded atoms and between atoms and lone pairs were also taken into account. As expected, the Xe−F

Table 5. Calculated and Experimental Heats of Formation of the Binary Xenon Fluorides and [XeF5]−a species

theory (0 K)a

theory (298 K)a

exptl equilibrium (0 K)b

exptl PI (0 K)c

XeF2 XeF4 [XeF5]− XeF6

−23.3 ± 0.3 −42.5 ± 1.0 −160.6 ± 1.0 −55.9 ± 2.0

−23.9 ± 0.3 −43.5 ± 1.0 −162.9 ± 1.0 −58.3 ± 2.0

−25.3 −50.2

−28.0 ± 0.5 −57.2 ± 2

−68.1

−(90−3+8)

From ref 95; CCSD(T), values are in kcal mol−1. bFrom ref 50. cFrom ref 98. Reproduced with permission from ref 95. Copyright 2005 American Chemical Society. a

F

DOI: 10.1021/cr500427p Chem. Rev. XXXX, XXX, XXX−XXX

Chemical Reviews

Review

experimental heats of formation, which lead to calculated heats of formation that are less negative than the experimental values. The calculated total bond dissociation energy for XeF4 is 7.7 kcal mol−1 lower than the experimental equilibrium value and 15.1 kcal mol−1 lower than the photoionization value.98 The experimental heats of formation for XeF4 and XeF6 are too large and are likely in error because it is difficult to obtain accurate experimental data on such reactive species. The atomization energies of XeF2 (63.65 kcal mol−1), XeF4 (125.08 kcal mol−1), and XeF6 (179.21 kcal mol−1) •increase with increasing number of fluorine atoms. The calculated dissociation energies at 0 K are given in eqs 30−35. These calculations show that loss of F2 from XeF2 is slightly less favorable than the loss of F2 from XeF4, and the loss of F2 from XeF4 is significantly more favorable than the loss of F2 from XeF6. As the number of fluorine atoms around Xe increases, the consequent steric crowding favors F2 elimination relative to the lower fluorides. XeF2 → Xe + F2

ΔH = 23.3 kcal mol−1 (30)

XeF4 → XeF2 + F2

ΔH = 19.2 kcal mol−1 (31)

XeF4 → Xe + 2F2

ΔH = 42.5 kcal mol−1 (32)

XeF6 → XeF4 + F2

ΔH = 13.4 kcal mol−1 (33)

Figure 5. Mössbauer spectra of XeCl4 and XeF4 referenced to natural abundance xenon hydroquinone clathrate. Reproduced with permission from ref 103. Copyright 1964 American Institute of Physics.

ray from metastable 129mXe results in ground state 129Xe and XeCl4. Xenon tetrachloride is stable for at least ca. 10−9 s, the lifetime of 129mXe. The doublet quadrupole splitting observed in the 129Xe Mössbauer spectrum that results from allowed transitions between the excited 129mXe (I = 3/2) state and the 129 Xe (I = 1/2) ground state and a nonspherical charge distribution around the xenon atom is consistent with the formation of planar XeCl4. The quadrupole splitting of XeCl4 (25.60 ± 0.18 mm s−1) is considerably less than that of XeF4 (41.74 ± 0.29 mm s−1)105 and can be attributed to the greater p electron imbalance due to the greater electronegativity of fluorine.106 Thus, the p electron imbalance is greater in XeF4 than in XeCl4 as reflected in their respective calculated valence shell electron configurations, 5s25p4.1 and 5s25p2.9. There was no appreciable yield of XeCl2 (quadrupole splitting, 28.2 ± 0.1 mm s−1)73 or Xe0, which does not display a quadrupole splitting. 129

XeF6 → XeF2 + 2F2 ΔH = 32.6 kcal mol−1 (34) XeF6 → Xe + 3F2

ΔH = 55.9 kcal mol−1 (35)

Significant variations in the standard entropy, S°, determinations for XeF4 occur in the early literature.99−101 A survey of the standard entropy determinations has accounted for the early discrepancies.51 One reason for variations in the standard entropy values was that a small sample (1.8727 g) of unspecified purity was initially used to determine the heat capacity of XeF4.101 The heat capacity was later redetermined from 20 to 300 K with a large (65.948 g), very pure sample of XeF4.51 The vapor pressure equation determined for XeF4100 (valid between 275 and 390.25 K, eq 36) was used along with thermodynamic data derived from the heat capacity to arrive at a S° value of 323.17 ± 2.0 J K−1 mol−1, which is in good agreement with the value determined from the fundamental vibrational frequencies94 of 323.98 ± 0.4 J K−1 mol−1. Agreement of the entropy derived from low-temperature heat capacity measurements and the entropy derived from vibrational data provides some validation of the assignments of the ν5 (formerly ν4) and ν7 fundamentals of XeF4, which were never directly observed but were derived from very weak bands that were attributed to overtones (see Table 4). log Pmm = −

2.5. Molecular Addition Compounds of XeF4

2.5.1. XeF4·XeF2. The molecular addition compound, XeF4· XeF2, was discovered in the course of developing a synthetic protocol for XeF4.107 This compound formed when the direct high-temperature fluorination of Xe by F2 to XeF4 was incomplete, yielding mixtures of XeF2 and XeF4.107 The compound was subsequently prepared by fusion of an equimolar mixture of XeF2 and XeF4 in a Kel-F tube.92 The compound was initially believed to be a high-density phase of XeF4 based on preliminary X-ray crystallographic data and its infrared spectrum.107 The fully refined X-ray crystal structure; however, revealed the coexistence of XeF4 and XeF2 in a 1:1 molar ratio with XeF2 and XeF4 occupying separate and alternating planes, giving the layered structure depicted in Figure 6.108 The geometrical parameters of XeF4 (Xe−F, 1.95(1) and 1.94(1) Å; F−Xe−F, 89.0(4)°) and XeF2 (Xe−F, 2.01(1)Å) do not differ substantially from those of solid XeF4 (1.951(2) and 1.954(2) Å; F−Xe−F, 90.0(1)°)13 and XeF2 (1.999(4) Å, F−Xe−F, 180.0°).109 Moreover, the shortest F2Xe---FXeF3 (3.28−3.42 Å), FXeF--XeF4 (3.36 Å), and F3XeF---XeF4 (3.37 Å) contact distances are only somewhat less than the sum of the fluorine and xenon van der Waals radii (3.63 Å).110 Consequently, XeF4·XeF2 is best described as a molecular addition compound. The Raman spectrum of solid XeF4·XeF2 has been compared to those of XeF4 and XeF2.92 Bands at 548 and 219 cm−1 were

3226.21 − 0.43434 log T + 12.301738 T (36)

2.4. Mössbauer Spectra of XeF4 and XeCl4

An attempt to synthesize the endothermic compound, XeCl4, by reaction of XeF4 with BCl3 was reported (eq 37).102 No XeCl4 was detected, and only Xe and Cl2 gases were formed (eq 38). 3XeF4 + 4BCl3 → 3XeCl4 + 4BF3

(37)

3XeF4 + 4BCl3 → 3Xe + 2Cl 2 + 4BF3

(38)

Evidence for the existence and structure of XeCl4 has been obtained by Mössbauer emission spectroscopy (Figure 5).103,104 This study utilized the β− decay of 129I in K129ICl4·H2O to form the excited nuclear state, 129mXe, at 4.2 K. Emission of a 40 keV γG

DOI: 10.1021/cr500427p Chem. Rev. XXXX, XXX, XXX−XXX

Chemical Reviews

Review

Figure 6. Packing in the X-ray crystal structure of XeF4·XeF2 (P21/c). Reproduced with permission from ref 108. Copyright 1963 American Chemical Society.

readily assigned to XeF4, but assignment of three intense lines at 511, 508, and 503 cm−1 could not be assigned with certainty. Because the bands corresponding to XeF4 (548 and 219 cm−1) were not factor-group split as they are in solid XeF4 (see Table 4), it was assumed that only one of the three unassigned lines corresponds to ν4 of XeF4 which was tentatively assigned to 511 cm−1. The remaining two bands (503 and 508 cm−1) were tentatively assigned to a factor-group split XeF2 mode. Although the Xe−F bond lengths of XeF4 and XeF2 in XeF4·XeF2 show little change relative to their pure phases, the Xe−F stretching frequencies are shifted to somewhat higher frequencies relative to those of solid XeF4 (551, ν4(Bg); 543, ν1(Ag); 506, ν5(Bg); 503, ν2(Ag); 238, ν6(Bg); and 216, ν3(Ag); space group P21/n (C2h))92 and XeF2 (496 cm−1, ν1(A1g), space group I4/mmm (D4h)).111 2.5.2. ([XeF5][CrF5])4·XeF4. The reaction of CrF2 with excess XeF6 yielded ([XeF5][CrF5])4·XeF4, which was characterized by single-crystal X-ray diffraction and vibrational spectroscopy.112 A mixture of CrF2 and excess XeF6 was heated at 120 °C for 10 d followed by removal of the reduction product, XeF4, and excess XeF6 under dynamic vacuum. Heating ([XeF5][CrF5])4·XeF4 at 50 °C for 70 h led to further removal of XeF6 and XeF4 to yield XeF6·2CrF4. Single crystals of ([XeF5][CrF5])4·XeF4 were obtained by recrystallization from aHF.112 The [CrF5]− anions form infinite parallel chains of distorted CrF6-octahedra that are separated by rows of alternating XeF4 molecules and [XeF5]+ cations and are linked to each other through long (F5Cr)−F--Xe(F4)---F−(CrF5) bridge interactions (3.126(7) Å) (Figure 7). The xenon valence electron lone pairs lie above and below the XeF4-plane, preventing the interacting fluorine ligands of the [CrF5]− anions from approaching XeF4 along the axial lone pair axis. Instead, the approach occurs between the lone pairs and the equatorial fluorine ligands of XeF4, thereby minimizing the repulsions between the interacting fluorine ligands and the xenon electron lone pairs of XeF4. The two crystallographically distinct [XeF5]+ cations also interact with the anions so that one xenon is eight-coordinated and the other is nine-coordinated. The Xe−F bond lengths of XeF4 (1.90(2) and 1.97(2) Å) and F−Xe−F bond angles (91.6(7)° and 84.5(5)°) deviate only slightly from the geometric parameters of XeF4 (see above). Four vibrational bands for the XeF4 molecule in [XeF5][CrF5])4·XeF4 were reported.112,113 Three frequencies (554, 545, and 504 cm−1) in the Raman spectrum are very similar to those of XeF4 (ν1, 543 and ν2, 503 cm−1). It was suggested that the bands observed at

Figure 7. Coordination environment of XeF4 in ([XeF5][CrF5])4·XeF4 (Pbca). Reproduced with permission from ref 112. Copyright 1992 Elsevier.

545 and 554 cm−1 likely arise from a single factor-group split vibrational mode. One broad band (563 cm−1) in the infrared spectrum was reported to be similar to that of XeF4 (ν6, 561 cm−1), but the ν6 fundamental of XeF4 in the complex occurs at 584.0 cm−1.93 The similarity of the structural and vibrational parameters of XeF4 in ([XeF5][CrF5])4·XeF4 to those of XeF4 indicates that ([XeF5][CrF5])4·XeF4 is best described as a molecular addition compound. The magnetic behavior of ([XeF5][CrF5])4·XeF4 was also investigated114 to deduce the magnetic characteristics of Cr(IV). The study revealed that antiferromagnetic superexchange interactions are present within the [CrF5]− chains. The zerofield cooled and field cooled magnetic measurements obeyed the Curie−Weiss law in the temperature range 25−300 K, where C = 4.11 emu K mol−1, and θp = −15 K. No minimum was observed for the reciprocal magnetic susceptibility at temperatures as low as 2 K. The vertex-sharing CrF6 octahedra lead to weak interchain exchange interactions. The negative value of θp is characteristic of antiferromagnetic interactions that arise from vertex-sharing octahedra.

3. FLUORIDE ION DONOR PROPERTIES OF XeF4 3.1. [Mg(XeF2)(XeF4)][AsF6]2

Xenon tetrafluoride is a weak fluoride-ion acceptor as exemplified by the formation of stable [XeF5]− salts (see section 4).10 Xenon tetrafluoride is also fluoro-basic and functions as a ligand. A significant number of adducts have been prepared in which XeF238,44 and KrF2115 molecules are F-coordinated to metal cation centers. The syntheses of XeF4 coordination compounds are more challenging than those of XeF2 due to the lower fluorobasicity of XeF4. The only metal cation complex of XeF4 that is presently known has been obtained from the reactions of a 1:1 or a higher molar ratio of XeF4:XeF2 with Mg(AsF6)2 in aHF, which exclusively yielded [Mg(XeF2)(XeF4)][AsF6]2.68 Reactions involving excess XeF4 yielded [Mg(XeF2)(XeF4)][AsF6]2 and unreacted XeF4, which was removed under dynamic vacuum at room temperature H

DOI: 10.1021/cr500427p Chem. Rev. XXXX, XXX, XXX−XXX

Chemical Reviews

Review

(see below). When 2:1 or higher molar ratios of XeF2:[Mg]2+ were used, the only product isolated was [Mg(XeF2)2][AsF6]2, which is a further affirmation of the greater fluoro-basicity of XeF2 relative to that of XeF4.68 The complex salt, [Mg(XeF2)(XeF4)][AsF6]2, is currently the only example known in which XeF4 functions as a ligand and, excluding the molecular addition compound XeF2·XeF4, is a rare example of a Xe(II)/Xe(IV) mixed oxidation state compound (see section 2.5.1). The coordination environment around [Mg]2+ is pseudooctahedral with four equatorial fluorine atoms belonging to four crystallographically distinct [AsF6]− anions and two axial fluorine atoms belonging to one XeF2 and one XeF4 molecule (Figure 8).

Figure 9. Packing diagram for [Mg(XeF2)(XeF4)][AsF6]2 viewed along the b-axis of the unit cell showing the alteration of the positions of the XeF2 and XeF4 molecules. Reproduced with permission from ref 68. Copyright 2009 Wiley−VCH.

those of uncoordinated XeF2 (1.999(4) Å)77 and those of [XeF]+ in [XeF][SbF6].109 The bridging (Xe---Fb, 2.059(7) Å) and terminal (Xe−Ft, 1.936(7) Å) bonds of the coordinated XeF2 are elongated and contracted when compared to the Xe−F bonds in uncoordinated XeF2,109 whereas they are contracted and elongated when compared to the Xe---Fb and Xe−Ft bonds in [XeF][SbF6] (2.278(2) and 1.885(2) Å, respectively). As observed in coordinated XeF4, the Xe−Ft bond is shorter than the Xe---Fb bond. The Raman spectrum of a single crystal of [Mg(XeF2)(XeF4)][AsF6]2 showed two strong bands at 552 and 605 cm−1 and a shoulder at 596 cm−1, which were attributed to Xe−F stretching bands of coordinated XeF4. The Xe−F stretching frequencies are intermediate with respect to those of free XeF4 (Table 4) and [XeF3][SbF6] (see section 3.2), consistent with XeF4 coordination, and polarization of the bridging Xe−F bond. A strong band at 575 cm−1 was assigned to the stretching mode of the terminal Xe−F bond of coordinated XeF2, and a band at 460 cm−1 was tentatively attributed to a coupled vibrational mode of XeF2 and XeF4, but was not explicitly assigned. In the same study, the reaction of a “higher molar excess” of XeF4 with [Mg][AsF6]2 in aHF gave a Raman spectrum reminiscent of [Mg(XeF2)(XeF4)][AsF6]2. However, the band associated with XeF2 at 575 cm−1 was weak, suggesting that an XeF2 adduct also formed as a minor species. The observation of two intense Raman bands at 549 and 622 cm−1, with a shoulder at 590 cm−1, suggested the presence of a compound containing coordinated XeF4 that was structurally similar to [Mg(XeF2)(XeF4)][AsF6]2. This compound was tentatively formulated as “Mg[AsF6]2 + nXeF4” and was assumed to contain at least two coordinated XeF4 molecules to achieve six-coordination for the central [Mg]2+ cation. Attempts to crystallize this compound from aHF over a 4-week period only resulted in the isolation of crystalline [Mg(XeF2)(XeF4)][AsF6]2. The presence of small amounts of XeF2 and the apparent reduction of Xe(IV) to Xe(II) could not be accounted for, but may be related to the long crystallization period.

Figure 8. Coordination environment of [Mg]2+ in the crystal structure of [Mg(XeF2)(XeF4)][AsF6]2 (P21/n). Reproduced with permission from ref 68. Copyright 2009 Wiley−VCH.

The XeF2 and XeF4 groups are trans to one another with very short [Mg]2+---F contacts (1.935(7) and 1.956(7) Å, respectively). The greater fluoride ion donor strength of XeF2 is not clearly reflected in the bridging Xe−F bond length differences, which are equal within ±3σ. The [Mg]2+ cations and [AsF6]− anions form chains in which the trans-coordinated XeF2 and XeF4 molecules alternate (Figure 9).68 A long contact (3.103(8) Å) between the terminal fluorine atom of XeF2 in one layer and the xenon atom of XeF4 in an adjacent layer presumably contributes to the stability of the lattice. It is noteworthy that the Xe−F bonds of the coordinated XeF4 molecule are intermediate with respect to those of uncoordinated XeF413 and the [XeF3]+ cation in β-[XeF3][SbF6].116 The bridging Xe---Fb bond length is significantly longer (2.083(6) Å) than the Xe−F bonds of free XeF4 (1.951(2) and 1.954(2) Å), but significantly shorter than the bridging Xe---Fb bond in β-[XeF3][SbF6] (2.485(1) Å). The Xe−F bonds trans (Xe−Ft , 1.871(7) Å) and cis (Xe−Fc , 1.916(7) and 1.906(7) Å) to the Xe---Fb bridge bond are contracted relative to the Xe−F bond of free XeF4. The Xe−Ft and Xe−Fc bond lengths are equal, within ±3σ, to the Xe−Ft and Xe−Fc bonds in β-[XeF3][SbF6] (trans, 1.839(2) Å; cis, 1.894(2) and 1.901(2) Å). As observed in β-[XeF3][SbF6], the Xe−Ft bond is shorter than the Xe−Fc bonds in [Mg(XeF2)(XeF4)][AsF6]2. Following the same pattern, the Fc−Xe−Fc angle of coordinated XeF4 (172.4(3)°) is also intermediate with respect to those of free XeF4 (180.0(1)°)13 and β-[XeF3][SbF6] (159.45(8)°).116 Furthermore, the Xe−F bonds of the coordinated XeF2 molecule are intermediate with respect to

3.2. Salts of the [XeF3]+ Cation

Although a large number of XeF2 and XeF6 complexes with strong fluoride ion acceptor molecules were reported at an early stage in the development of noble-gas chemistry,117,118 no XeF4 I

DOI: 10.1021/cr500427p Chem. Rev. XXXX, XXX, XXX−XXX

Chemical Reviews

Review

complex was identified. An early report of an unstable XeF4− SbF5 complex, which was believed to have the composition XeF4· 2SbF5, was not substantiated by analytical, spectroscopic, or other methods.119 Another early report claiming XeF4 interacts with SbF5 or TaF5 to form XeF2 adducts is erroneous.120 The formation of the latter adducts was most likely due to the use of XeF4 that was significantly contaminated with XeF2. Bartlett and co-workers121 reported that XeF4 does not form stable adducts with either AsF5 or IrF5 in BrF5 solution or with RuF5.122 Xenon tetrafluoride thus came to be regarded as a very weak fluoride ion donor and led to the ranking XeF6 > XeF2 ≫ XeF4 for the fluoride ion donor abilities of the binary xenon fluorides. 3.2.1. Syntheses of [XeF3]+ Salts and Solution Studies. The earliest indication of the possible existence of an XeF4 adduct with a fluoride ion acceptor was obtained when XeF4 or XeF2/XeF4 mixtures were dissolved in SbF5. Dissolution was accompanied by gas evolution, to give green solutions. Upon removal of excess SbF5 by pumping at 25 °C, an easily decomposed white solid, which could not be isolated or identified, and the then known yellow XeF2·2SbF5 complex were obtained. The 19F NMR spectra of these solutions were reported to contain peaks arising from XeF2, which were later shown to be due to [XeF]+ and its 129Xe satellites,123 and other unidentified peaks arising from fluorine on xenon. The gas that was evolved in these early experiments was likely Xe and/or O2, which could result from XeII/XeIV reduction if the SbF5 is contaminated with HF/H2O ([H3O]+). The green solution color is now known to arise from the [Xe2]+ cation, which results from the reaction of Xe with [XeF]+ in SbF5.124−126 A subsequent 19F NMR study of SbF5 solutions of XeF4 and XeF2/XeF4 mixtures prepared from pure XeF2 and XeF4 and freshly distilled SbF5 yielded yellow solutions and revealed the presence of the [XeF3]+ cation (Figure 10). The ionization of XeF2 to [XeF][SbnF5n+1] in SbF5 increases the ionicity and decreases the viscosity of the solvent mixture so that XeF4 dissolves at 25 °C, whereas XeF4 must be heated to 50 °C in pure SbF5 to effect dissolution. The decreased viscosity also increases spectral resolution. The 19F NMR spectrum consisted of a second-order doublet and triplet (AB2) spin−spin coupling pattern (where A = Fax and B = Feq) flanked by natural abundance 129 Xe satellites, where δ(19Feq) = 23.0 ppm, δ(19Fax) = 38.7 ppm, 2 19 J( Feq−19Fax) = 174 Hz, 1J(129Xe−19Feq) = 2440 Hz, and 1 129 J( Xe−19Fax) = 2620 Hz.7,127 The integrated 19F resonance intensities are consistent with the expected T-shaped (disphenoidal) C2v geometry of the cation predicted by the VSEPR model of molecular geometry.128 A 129Xe NMR study89 reported the 129Xe NMR spectrum of [XeF3]+ in SbF5 (25 °C). The 129Xe spectrum consisted of an overlapping doublet (1J(129Xe−19Feq) = 2384 Hz) of triplets (1J(129Xe−19Fax) = 2609 Hz) centered at 595 ppm with respect to external liquid XeOF4 (Figure 11). The observed T-shaped geometry results from a trigonal bipyramidal, AX3E2 VSEPR arrangement of five electron pairs in the valence shell of xenon where a fluorine atom and the valence electron lone pairs of xenon occupy the equatorial plane and two fluorine atoms occupy the axial positions of the trigonal bipyramid. When solutions of XeF4 in SbF5 solution were pumped to constant weight, a yellow compound was obtained having a mass that was consistent with eq 39 and the formulation, [XeF3][Sb2F11]. This salt formulation and the geometry of [XeF3]+ were confirmed by comparison of the Raman spectrum of [XeF3][Sb2F11] with those of the isostructural ClF3, BrF3, and IF3 molecules (see section 3.2.3).129

Figure 10. 19F NMR spectrum (56.4 MHz, 26 °C) of the [XeF3]+ cation (0.20 m XeF4 and 0.50 m XeF2 in SbF5 solution): (A) axial fluorines and (a) 129Xe satellites; (B) equatorial fluorine and (b) 129Xe satellites. Modified with permission from ref 127. Copyright 1974 American Chemical Society.

Figure 11. 19F NMR spectrum (56.4 MHz, 26 °C) of the [XeF3]+ cation (0.20 m XeF4 and 0.50 m XeF2 in SbF5 solution): (A) axial fluorines and (a) 129Xe satellites; (B) equatorial fluorine and (b) 129Xe satellites. Modified with permission from ref 89. Copyright 1978 American Chemical Society. SbF5, 50 ° C

XeF4 + 2SbF5 ⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯→ [XeF3][Sb2 F11]

(39)

Since the discovery of the first [XeF3]+ salt, a number of other [XeF3]+ salts have been synthesized by the interaction of the strong group 15 Lewis acid fluoride ion acceptors, SbF5, BiF5, and AsF5, with XeF4. Two phases of [XeF3][SbF6] have been identified by Raman spectroscopy (see below), a high-temperature α-phase and a room-temperature β-phase. The pale yellowgreen α-[XeF3][SbF6] salt was synthesized by fusing [XeF3][Sb2F11] with an excess amount of XeF4 at 80 °C (eq 40).129 80 ° C

XeF4 + [XeF3][Sb2 F11] ⎯⎯⎯⎯→ 2α‐[XeF3][SbF6]

(40)

The room-temperature β-phase has been prepared by reaction of SbF5 and a stoichiometric excess of XeF4 in aHF. Crystallization from aHF at room temperature followed by removal of the solvent and excess XeF4 only yielded pale yellowgreen β-[XeF3][SbF6] (eq 41).8 The fluorine-bridged [Xe2F7]+ cation did not form when the aforementioned reactions were carried out in the presence of excess stoichiometric amounts of XeF4.9 This contrasts with the known Xe(II) and Xe(VI) analogues, [Xe2F3]+130,131 and [Xe2F11]+,132−134 whose salts have been structurally characterized. J

DOI: 10.1021/cr500427p Chem. Rev. XXXX, XXX, XXX−XXX

Chemical Reviews

Review

aHF, RT

XeF4 + SbF5 ⎯⎯⎯⎯⎯⎯⎯⎯→ β‐[XeF3][SbF6]

(41)

135

136

The [XeF3][BiF6] and [XeF3][Bi2F11] salts have been prepared by reaction of stoichiometric amounts of XeF4 and BiF5 in aHF (eqs 42 and 43). In the case of [XeF3][BiF6], the reaction was carried out at 40 °C to give a partially soluble solid, contrasting with the [XeF3][Bi2F11] salt, which was soluble in aHF at room temperature. Both salts are pale yellow to yellow in color. aHF, 40 ° C

XeF4 + BiF5 ⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯→ [XeF3][BiF6]

Figure 13. X-ray crystal structures of [XeF3][BiF6] (P1̅). Redrawn with permission from ref 135. Copyright 1977 Royal Society of Chemistry.

(42)

aHF, RT

XeF4 + 2BiF5 ⎯⎯⎯⎯⎯⎯⎯→ [XeF3][Bi 2F11]

(43)

lone pair domains. The Fax−Xe−Fax angle subtended by the axial fluorine atoms at xenon deviates significantly from linearity and is bent away from the more sterically demanding electron lone pair domains of xenon toward the equatorial fluorine atom (Figure 14). In all three structures, the [XeF3]+ cation strongly interacts with a fluorine ligand of the anion by means of a fluorine bridge (Fb; Table 6). The [XeF3]+ cation in β-[XeF3][SbF6] is coordinated to the fluorine atoms of two [SbF6]− anions with Xe---Fb distances of 2.484(2) and 2.646(2) Å. In the salts, [XeF3][Sb2F11] and [XeF3][BiF6], the [XeF3]+ cations are only strongly coordinated to a single Xe---Fb bond (2.25(2) [BiF6]−, 2.490(4) [Sb2F11]− Å). The fluorine bridges in all three salts are significantly less than the sum of the fluorine and xenon van der Waals radii (3.63 Å)110 and are nearly coplanar with the fluorine atoms of the [XeF3]+ cation. These structural features indicate a significant degree of covalent character in the Xe---Fb bridge bond. The AX2E2 VSEPR arrangement at Fb causes the Xe---Fb− Pn angles to be bent ([SbF6]− 138.91(8)°, [BiF6]− 150.3°, [Sb2F11]− 169.0(2)°). The [BiF6]− salt is distinguished from the [SbF6]− and [Sb2F11]− salts by its shorter Xe---Fb bond and elongation of its Bi−Fb bond relative to the terminal Bi−F bonds of the ion pair. The tendency for the “incoming” Fb atom lone pair to avoid the Xe lone pairs is evident from the confinement of their Fb approaches to the XeF3-plane and away from the lone pairs in the trigonal plane defined by Feq, Xe, and the lone pairs of Xe. In the [SbF6]− and [Sb2F11]− salts, the two Fb---Xe−Fax angles are very similar, that is, 73.5(4)° and 125.6(4)° ([SbF6]− salt) and 72.7(3)° and 125.3(3)° ([Sb2F11]− salt). In these cases, the asymmetries of the Fb---Xe---Fax contact angles are consistent with Fb atom approaches that occur between Fax and the trigonal plane to avoid the two electron lone pairs that occupy that plane. The corresponding Fax−Xe−Fax angles are 159.39(8)° for the [SbF6]− salt and 161.3(2)° for the [Sb2F11]− salt. In contrast, [XeF3][BiF6] displays nearly equal Fb---Xe−Fax angles of 89.7(9)° and 97.1(10)° and a wide Fax−Xe−Fax angle of 173.2(1)°. The angle difference arises from the stronger Xe--Fb interaction, with the incoming Fb lone pair approaching the midpoint of the edge defined by the two Xe lone pairs of the trigonal bipyramid, which forces the lone pairs into an arrangement that more closely resembles square-planar XeF4 and a AX4E2 VSEPR arrangement. In this case, the xenon coordination is between a trigonal bipyramidal AX3E2 and a square-planar AX4E2 arrangement. The structural differences are consistent with a fluoride ion affinity for BiF5 (466.9 kJ mol−1) that is significantly less than those of SbF5 (495.0 kJ mol−1) or Sb2F10 (526.8 kJ mol−1).144,145 3.2.2.2. [H5F4][SbF6]·2[XeF3·HF][Sb2F11] and [XeF3·HF][Sb2F11]. In an attempt to obtain a soluble [HOXeF2]+ salt suitable for crystal growth (see section 5.3), the synthesis of the

In an attempt to synthesize [XeF3][As2F11] from XeF4 and excess liquid AsF5 at −100 °C, only [XeF3][AsF6] was obtained (eq 44) with no evidence for [XeF3][As2F11].129 The [XeF3][AsF6] salt was stable in the presence of excess AsF5 but dissociated to XeF4 and AsF5 when AsF5 was pumped off at −78 °C. Estimates of the standard free energies for the formation of solid [XeF3][AsF6] and [XeF3][As2F11] from AsF5 liquid and XeF4 solid using volume based thermodynamics (VBT) suggest instability or marginal stability for [XeF3][AsF6] (ΔG° = 20 ± 26 kJ mol−1) and [XeF3][As2F11] (ΔG° = 14 ± 15 kJ mol−1) when the calculated errors are taken into account.109 The known room temperature-stable salts, [XeF3][SbF6] (ΔG° = −30 ± 54 kJ mol−1) and [XeF3][Sb2F11] (ΔG° = −54 ± 15 kJ mol−1), display more negative standard free energies. AsF5, −100 ° C

XeF4 + AsF5 ⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯→ [XeF3][AsF6 ]

(44)

+

3.2.2. Crystal Structures of [XeF3] Salts. 3.2.2.1. [XeF3][PnF6] (Pn = Sb, Bi) and [XeF3][Sb2F11]. The original crystal structures of β-[XeF3][SbF6]8 and [XeF3][Sb2F11]9 have been redetermined at −173 °C (Figure 12),109,116 and represent the

Figure 12. X-ray crystal structures of β-[XeF3][SbF6] (P21/n) and [XeF3][Sb2F11] (P1̅). Thermal ellipsoids are drawn at the 50% probability level. Redrawn with permission from refs 116 and 109. Copyright 2013 and 2010 American Chemical Society.

highest precision structures for these [XeF3]+ salts. Consequently, these structures, along with the original lower precision structure of [XeF3][BiF6]102 (Figure 13), are referred to in the ensuing discussion. A comparison of the geometrical parameters for the known [XeF3]+ salts, ClF3, BrF3, and IF3, is provided in Table 6. The crystal structures can be approximately described as consisting of [XeF3]+ cations and [PnF6]−/[Sb2F11]− anions. The [XeF3]+ cations have the planar T-shaped geometry observed in solution by 19F NMR spectroscopy and in the solid-state Raman spectra of these salts (see above). As predicted by the VSEPR rules of molecular geometry,128 the axial Xe−Fax bonds are longer than the equatorial Xe−Feq bond due to the greater repulsive interactions between the axial bond pair domains and their neighboring equatorial bond pair and electron K

DOI: 10.1021/cr500427p Chem. Rev. XXXX, XXX, XXX−XXX

Chemical Reviews

Review

Table 6. Comparison of the Structures of [XeF3]+, [XeF3]+·HF, IF3, BrF3, and ClF3 bond length/anglea X−Feq (Å) X−Fax (Å) X---Fb (Å) Fax−X−Fax (deg) Fax−X−Feq (deg) Feq−X---Fb (deg) Xe---Fb−Pn (deg) ref a

β-[XeF3][SbF6]

[XeF3][BiF6]

[XeF3][Sb2F11]

[XeF3·HF][Sb2F11]b

IF3

1.840(2) 1.893(2) 1.901(2) 2.484(2) 159.39(8) 79.72(8) 79.74(7) 152.74(7) 138.91(8) 116

1.81(2) 1.92(3) 1.94(3) 2.25(2) 173.2(1) 84.2(1.1) 89.1(1.1) 173.8(1.1) 150.3 135

1.832(4) 1.883(4) 1.908(4) 2.490(4) 161.3(2) 80.3(2) 81.2(2) 154.5(2) 169.0(2) 109, 137

1.838(2) 1.880(2) 1.890(2) 2.462(2) 161.12(6) 81.0(1) 81.4(1) 160.2(1)

1.872(4) 1.983(3)

1.721 1.810

1.598 1.698

160.3(2) 80.2(1)

86.2

87.5

139, 140

141−143

116

BrF3

ClF3

160.3(2) 138

X = Xe, I, Br, Cl. b[XeF3·HF][Sb2F11] in the salt [XeF3·HF][Sb2F11]·[H5F4][SbF6].85

longer than that of the parent salt, [XeF3][Sb2F11] (Xe---Fb, 2.490(4) Å). The Xe−Feq bond lengths of [H5F4][SbF6]·2[XeF3·HF][Sb2F11] (1.838(2) Å) and [XeF3][SbF6] (1.840(2) Å) are equal, but significantly shorter than the Xe−F bonds of XeF4 (1.951(2), 1.954(2) Å).13 The Xe−Fax bond lengths (1.880(2) and 1.890(2) Å) are also very similar to those of the β-[XeF3][SbF6] salt (1.893(2) and 1.901(2) Å), and are also significantly shorter than in XeF4. In contrast with other [XeF3]+ salts, which fluorine bridge to their respective anions,8,9,109,116,135 the [XeF3]+ cation of [H5F4][SbF6]·2[XeF3·HF][Sb2F11] has a short contact (2.462(2) Å) to an HF molecule that lies in the [XeF3]+ plane, so that [XeF3·HF]+ is better formulated as an HF Lewis acid− base adduct, [F3Xe---FH]+, rather than as a protonated XeF4 molecule, [F3XeF---H]+. This Xe---F contact is similar to the short secondary contact that occurs between the cation and anion of β-[XeF3][SbF6] (2.485(1) Å) and [XeF3][Sb2F11] (2.490(4) Å). The Feq−Xe−Fax (81.4(1)°, 81.0(1)°), Fax−Xe− Fax (162.3(1)°), and Fax−Xe---FF (78.8(1)°, 118.77(9)°) bond angles of XeF3·HF are also very similar to those of [XeF3][SbF6] (Feq−Xe−Fax, 79.72(8)°, 79.74(7)°; Fax−Xe−Fax, 159.45(8)°; Fax−Xe---FSb, 73.06(6)°, 127.47(6)°). The positionally averaged Xe−Fax/Xe−Feq and Xe−Fax/Xe−FH bond lengths of [XeF3· HF][Sb2F11] are 1.865(1) and 2.186(1) Å, respectively, and are in good agreement with the averages of the corresponding bond lengths in [H5F4][SbF6]·2[XeF3·HF][Sb2F11], 1.859(3) and 2.176(3) Å. 3.2.3. Raman Spectra of [XeF3]+ Salts. The vibrational spectra of [SbF6]−, [Sb2F11]−, [BiF6]−, [Bi2F11]−, and [AsF6]− salts of [XeF3]+ have been studied in detail by Raman spectroscopy.129,136,137 As inferred from the X-ray crystal structures of β-[XeF3][SbF6], [XeF3][Sb2F11], and [XeF3][BiF6], these salts are comprised of ion pairs that result from partial abstraction of a fluoride ion by [XeF3]+ from the counteranion. Consequently, the anion geometries are distorted from their gas-phase geometries, which is also evident in their Raman spectra. The original papers cited herein should be consulted for full discussions of the anion spectra. The six vibrational modes of the cation were initially assigned by comparison with the vibrational spectra of the isovalent ClF3, BrF3, and IF3 molecules.129 The vibrational frequencies of [XeF3]+ were assigned under C2v symmetry (Table 7) and correspond to the vibrational representation Γvib = 3A1 + 2B1 + B2, where the σv(xz) is the molecular plane and all modes are Raman and infrared active.91 Table 8 provides a comparison between the vibrational frequencies and their assignments for the [Sb2F11]− salt of [XeF3]+, and those of ClF3, BrF3, and IF3. The

Figure 14. A depiction of the Fb---Xe fluorine bridge contacts in β[XeF3][SbF6]. Reproduced with permission from ref 8. Copyright 1974 American Chemical Society.

[SbF6]− salt was attempted in the superacid medium, aHF/ SbF5.116 In addition to [XeF3][SbF6], two complex [XeF3]+ salts, [H5F 4][SbF 6]·2[XeF3·HF][Sb2F11] (P1̅) and [XeF 3·HF][Sb2F11] (C2/c), were isolated from HF/SbF5 solutions of XeOF2·nHF (Figure 15). Among the interesting features of these salts are the HF acidium ion, [H5F4]+, which had not been structurally characterized before, and [XeF3·HF]+.

Figure 15. XeF3+ cations in the X-ray crystal structures of (a) [H5F4][SbF6]·2[XeF3·HF][Sb2F11] (P1̅), and (b) [XeF3][SbF6] (P21/ c). Thermal ellipsoids are shown at the 50% probability level. Redrawn with permission from ref 116. Copyright 2013 American Chemical Society.

There is a 2-fold positional disorder between an Fax atom and the F atom of coordinated HF in [XeF3·HF][Sb2F11]. The secondary coordination spheres of xenon are very similar in [XeF3·HF][Sb2F11] and [H5F4][SbF6]·2[XeF3·HF][Sb2F11], with long Xe---F contacts ranging from 2.935(1) to 2.967(1) Å and from 2.847(3) to 3.088(3) Å, respectively. There is an additional contact (Xe---F, 3.088(3) Å) in the ordered crystal structure of [H5F4][SbF6]·2[XeF3·HF][Sb2F11] that occurs between [XeF3]+ and [SbF6]− of [H5F4][SbF6]. The Xe---F cation−anion contacts found in [XeF3·HF]+ are significantly L

DOI: 10.1021/cr500427p Chem. Rev. XXXX, XXX, XXX−XXX

Chemical Reviews

Review

Table 7. Raman Spectra and Assignments for the [XeF3]+ Cation of [α-XeF3][SbF6], [β-XeF3][SbF6], [XeF3][AsF6], [XeF3][BiF6], [XeF3][Sb2F11], and [XeF3][Bi2F11]a,b α-[XeF3][SbF6]

β-[XeF3][SbF6]

[XeF3][BiF6]

[XeF3][AsF6]

[XeF3][Sb2F11]

[XeF3][Bi2F11]

643(100)

663(100) 643(56 576(94) 564(94) 212(3) 199(2) 612(25) 604(21) 335(2) 318(2)

645(100)

643(85)

645(39)

648(100)

ν1(A1)

ν(XeFeq)

557(83)

571(100)

583(100)

574(98)

ν2(A1)

νsym(XeF2ax)

ν3(A1)

δsym(FaxXeFax)

ν4(B1)

νasym(XeF2ax)

ν5(B1)

δasym(FaxXeFax)

ν6(B2)

δ(XeF3)o.o.p. ν(Xe---F)d ref

573(88) 205(2) 609(9) 305, sh

232(1) 129

257(1) 129

198(7) 609(10)

253(3) 377(1) 136

209(6) 608(58) 607(38) 316(18)

619, sh

611(36) 606(35)

305, sh

129

129

251(9) 444(9) 136

[XeF3]+ (C2v) assgntsc

Frequencies are given in cm−1. bValues in parentheses denote relative Raman intensities. The abbreviations denote a shoulder (sh), out of plane (o.o.p.), equatorial (eq), axial (ax), symmetric (sym), and asymmetric (asym). cThe corresponding reference should be consulted for the anion frequencies and their assignments. dAssignments of the ν(Xe---F) stretching bands are tentative. a

Table 8. Infrared and Raman Spectra of [XeF3]+, ClF3, BrF3, and IF3a,b

Frequencies are given in cm−1. bValues in parentheses denote relative Raman intensities. The abbreviations denote polarized band (p), depolarized band (dp), very strong (vs), strong (s), medium (m), weak (w), very weak (vw), and shoulder (sh). c[XeF3][Sb2F11] from ref 129. dFrom ref 146. e From ref 148. fFrom ref 147. Reproduced with permission from ref 129. Copyright 1976 American Chemical Society. a

calculated [XeF3]+ frequencies are in good agreement with the experimental values and their assignments (see section 3.2.4). Each [XeF3]+ vibrational band of β-[XeF3][SbF6] is split into a doublet due to vibrational coupling within the unit cell (factorgroup splitting), whereas those of α-[XeF3][SbF6] are not split (Table 7). Modes associated with the Xe---F bridges have been assigned for several salts in Table 7; however, their assignments are tentative in view of their broadness and low intensities. On the basis of their crystal structures, [XeF3][BiF6]135 is more strongly ion-paired than β-[XeF 3 ][SbF 6 ] 116 or [XeF 3 ][Sb2F11];109 however, only the νsym(XeF2ax) and δsym(Fax−Xe− Fax) modes of [XeF3][BiF6] differ significantly from those of other salts listed in Table 7. 3.2.4. Theoretical Studies of the [XeF3]+ Cation. The gasphase [XeF3]+ cation has been studied by quantum-chemical methods (PBE1PBE/aug-cc-pVTZ, B3LYP/aug-cc-pVTZ, and MP2/aug-cc-pVTZ) (Table 9).116 The calculated geometry of [XeF3]+ is in accordance with the experimentally observed T-

shaped (C2v) geometry. All computational methods provide very good agreement with experiment for the Xe−Fax bond lengths, but tend to overestimate the Fax−Xe−Feq bond angles. The Xe− Feq bond lengths were very well reproduced using the PBE1PBE and MP2 methods, whereas the B3LYP and CCSD(T) methods overestimated these bond lengths.16,149 The B3LYP method reproduced the experimental frequencies better than the PBE1PBE and MP2 methods. All three methods overestimated the ν4(B1) vibration to be higher than ν1(A1), which is opposite to what is observed experimentally. An oxidant strength scale has been developed with the use of local density functional theory (LDFT) calculations.150 The scale ranks the oxidizing strengths of oxidative fluorinators and is based on the relative F+ detachment (FPD) energies of 36 oxidative fluorinators, which decrease with increasing oxidizer strength. The FPD energy is defined as [AF]+ → F+ + A. The [XeF3]+ cation ranks near the middle of the scale (152.4 kcal mol−1), but well below [KrF]+ (115.9 kcal mol−1), the strongest M

DOI: 10.1021/cr500427p Chem. Rev. XXXX, XXX, XXX−XXX

Chemical Reviews

Review

Table 9. Calculateda and Experimentalb Raman Spectra and Structural Parameters of [XeF3]+ calcd bond lengths (Å)

B3LYPc

PBE1PBEc

MP2d,e

CCSD(T)d

Xe−Fax

1.906

1.884

1.883

1.891

Xe−Feq bond angles (deg) Feq−Xe−Fax

1.867

1.845

1.838

1.853

Feq−Xe−Fax assgnts (C2v)

84.2

83.7

168.4

167.3

ν1(A1) ν2(A1) ν3(A1) ν4(B1) ν5(B1) ν6(B2)

83.3

83.0

166.6 166.0 vibrational frequencies (cm−1)c,d,g,h

644(16)[14] 598(29)[