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The Oxide Fluoride Chemistry of Bromine, Selenium and Sulphur Thesis submitted for the degree o f Doctor o f Philosophy at the University o f Leicester by LEE JO H N W O O T T O N Department Faculty University 1997 o of o UMI Number: U095961 All rights reserved INFORMATION TO ALL USERS The quality of this reproduction is dependent upon the quality of the copy submitted. In the unlikely event that the author did not send a complete manuscript and there are missing pages, these will be noted. Also, if material had to be removed, a note will indicate the deletion. Dissertation Publishing UMI U095961 Published by ProQuest LLC 2013. Copyright in the Dissertation held by the Author. Microform Edition © ProQuest LLC. All rights reserved. This work is protected against unauthorized copying under Title 17, United States Code. ProQuest LLC 789 East Eisenhower Parkway P.O. Box 1346 Ann Arbor, Ml 48106-1346 STATEMENT The experimental work described in this thesis has been carried out by the author in the Department of Chemistry at the University of Leicester between October 1993 and September 1996. The work has not been submitted, and is not presently being submitted, for any other degree at this or any other university. Department of Chemistry University of Leicester University Road Leicester U.K. L E I7 R H i Abstract The Oxide Fluoride Chemistry of Bromine, Selenium and Sulphur Lee J. Wootton The transition metal carbonyls [Re2 (CO)10], [Mn 2 (C O )10] and [Ru(CO)3 (PPh3)2], and elemental iodine have been reacted with Xe(OSeF5)2. The products have been fully characterised by a combination of mass spectrometry, infrared spectroscopy, 1 9 F, 13C and 3 1 P { 1 H} (where appropriate) NMR spectroscopies. Further characterisation of the novel compounds Xe(OSeF5)2, [Re(CO)5 (OSeF5)] and [Mn(CO) 5 (OSeF5)] by EXAFS spectroscopy is reported. An extensive review of the halogen oxide fluorides has been carried out and attempts were made to synthesise a range of fluorides and oxide fluorides of bromine. The bromine fluorides [BrF 2 ][AsF6], [BrF 4 ][Sb 2 F u ] , K[BrF4] and Cs[BrF6] were successfully characterised using EXAFS spectroscopy. The compound Cs[BrOF4] has been synthesised and the application of EXAFS spectroscopy has yielded internal bond parameters. The area of fluorosulphate chemistry has been reviewed and reactions have been carried out between the superacid H S 0 3F and a range o f transition metal carbonyl complexes and Ti, Hf and Zr derivatives. The complexes produced were characterised using mass spectrometry, infrared spectroscopy and 1 H, 1 3 C, 1 3 C{ 1 H} and 19F NMR spectroscopies. The protonation of the carbonyl clusters [Ir4 (CO)12], [Os3 (CO)12] and [Ru 3 (C O )12] by H S 0 3F was investigated. The systems were found to be the same as those previously observed for the superacid AHF. Contents Statement i Abstract ii Contents iii List of Tables ix List of Figures xii Acknowledgements xv Abbreviations xvi Chapter One Introduction 1.1 General Introduction 1 1.2 Characterisation 3 1.2.1 EXAFS spectroscopy 1.3 4 Summary 7 References 9 Chapter Two Oxidation Reactions using Xenon Bis(seflate) 2.1 Introduction 10 2.2 Preparative Routes to Compounds Containing the 14 Seflate Group 2.3 Stability of Seflate Compounds 16 2.4 Electronegativity of the Seflate Anion 17 2.5 Spectroscopic Characterisation of Seflate Compounds 21 iii 2.5.1 Fluorine-19 NMR spectroscopy 2 1 2.5.2 Vibrational spectroscopy 24 2.5.3 Mass spectrometry 26 2.5.4 X-ray crystallography and EXAFS spectroscopy 27 2 .6 Covalent Bonding 28 2.7 Ionic Bonding 29 2 .8 Xenon Bis(seflate) 31 2.9 Preparation and Properties of Xenon Bis(seflate) 33 2 .1 0 The Reaction Between [Re2 (CO)10] and Xe(OSeF 5 ) 2 39 2 .1 1 The Reaction Between [Mn2 (CO)10] and Xe(OSeF 5 ) 2 46 2 .1 2 The Reaction Between [Ru(CO)3 (PPh3)2] and Xe(OSeF 5 ) 2 53 2.13 The Reaction Between I2 and Xe(OSeF 5 ) 2 57 2.14 Discussion 6 6 References 70 Chapter Three Bromine Oxide Fluoride Chemistry 3.1 Introduction 75 3.2 Structures of the Oxide Fluorides 75 3.3 The Halogenyl Fluorides, X 0 2F 78 3.3.1 Chlory 1 fluoride 78 3.3.2 Bromyl fluoride 78 3.3.3 Iodyl fluoride 79 3.4 3.5 The Halogen Oxide Trifluorides, XOF 3 80 3.4.1 Chlorine oxide trifluoride 80 3.4.2 Bromine oxide trifluoride 81 3.4.3 Iodine oxide trifluoride 82 The Perhalogenyl Fluorides, X 0 3F 83 3.5.1 3.6 3.7 3.8 Perchloryl fluoride 83 3.5.2 Perbromyl fluoride 84 3.5.3 84 Periodyl fluoride The Halogen Dioxide Trifluorides, X 0 2 F 3 85 3.6.1 Chlorine dioxide trifluoride 85 3.6.2 Bromine dioxide trifluoride 8 6 3.6.3 Iodine dioxide trifluoride 8 6 The Halogen Oxide Pentafluorides, XOF 5 8 8 3.7.1 Chlorine oxide pentafluoride 8 8 3.7.2 Bromine oxide pentafluoride 8 8 3.7.3 Iodine oxide pentafluoride 8 8 The Halogen Oxide Fluorides 89 3.8.1 89 Chlorosyl fluoride 3.9 Summary 89 3.10 The Unusual Nature of Bromine (VII) 90 3.11 Area of Study 93 3.12 EXAFS Spectroscopic Study of the Bromine Fluorides 95 3.12.1 Discussion 96 3.13 The Synthesis and EXAFS Characterisation of Cs[BrOF4] 104 3.14 The Synthesis of K [B r04] 109 3.15 The Synthesis of B r0 3F 110 3.16 The Attempted Synthesis of BrOF 3 114 3.17 The Attempted Synthesis of B r0 2F 116 3.18 Conclusion 117 References 119 v Chapter Four Displacement and Oxidation Reactions using Fluorosulphonic Acid 4.1 Introduction 125 4.2 Properties 125 4.3 Synthetic Routes to Metal Fluorosulphate Complexes 129 4.3.1 129 Syntheses involving S2 0 6 F 2 or S 2 0 6 F 2 -H S 0 3F 4.3.1.1 Limitations of the S 2 0 6 F 2 -H S 0 3F system 131 4.3.2 Displacement reactions 134 4.3.3 Syntheses involving B rS 0 3F 136 4.3.4 Insertion reactions 136 4.3.5 Oxidising reactions involving H S 0 3F 137 4.4 Decomposition of Fluorosulphates 137 4.5 Spectroscopic Characterisation of Fluorosulphate 139 Compounds 4.5.1 4.6 Vibrational spectroscopy 139 4.5.2 X-ray crystallography 145 4.5.3 Fluorine-19 NMR spectroscopy 146 4.5.4 Mossbauer spectroscopy 146 4.5.5 Magnetic studies and electronic spectroscopy 146 Single Crystal X-ray Analysis of Fluorosulphate 147 Compounds 4.7 Recent Developments in Fluorosulphate Chemistry 153 4.7.1 153 Cationic carbonyl metal species 4.7.2 Superacids 157 4.8 Area of Study 160 4.9 The Reaction of [Ir4 (CO)12], [Ru 3 (CO)12] and 161 [Os 3 (CO)12] with H S 0 3F vi 4.9.1 Summary 164 4.10 The Reaction Between [Fe2 (CO)10] and H S 0 3F 166 4.11 168 The Reaction Between Re or Mn Carbonyl Derivatives and H S 0 3F 4.12 The Reaction Between [Cp2 MX2] (M = Ti, Zr or 172 H f and X = Me or Cl) and H S 0 3F 4.13 The Reaction Between [W(CO)6] and H S 0 3F 177 4.14 The Reaction Between [Mo(CO)6] and H S 0 3F 178 4.15 The Reaction Between [Co2 (CO)8] or [Cr(CO)6] 179 and HSO 3 F 4.16 Summary 180 References 181 Chapter Five Experimental 5.1 5.2 5.3 Handling of Materials 186 5.1.1 Metal vacuum line 186 5.1.2 Inert atmosphere dry box 186 Reaction Vessels 188 5.2.1 Metal reactors 188 5.2.2 Glass apparatus 188 5.2.3 Fluoroplastic apparatus 190 Analytical T echniques 192 5.3.1 192 Nuclear magnetic resonance spectroscopy 5.3.2 Infrared spectroscopy 192 5.3.3 192 Mass spectrometry 5.3.4 EXAFS spectroscopy 5.4 194 Solvents 195 vii 5.5 5.4.1 Anhydrous hydrogen fluoride 195 5.4.2 Dichloromethane 195 5.4.3 Acetonitrile 196 5.4.4 Fluorosulphonic acid 196 Preparation of Fluorides, Oxide Fluorides, Seflate and 196 Fluorosulphate Species 5.6 5.5.1 Preparation of XeF 2 196 5.5.2 Preparation of Xe(OSeF 5 ) 2 197 5.5.3 Preparation of K [B r04] 198 5.5.4 Reactions involving Xe(OSeF 5 ) 2 200 5.5.5 Preparation of BrF 3 201 5.5.6 Preparation of BrF 5 201 5.5.7 Preparation of K[BrF4], KtBrFg] andC s[B rF6] 201 5.5.8 Preparation of [BrF 2 ][AsF6] and [BrF 4 ][Sb 2 F n ] 202 5.5.9 Preparation of Cs[BrOF4] 203 5.5.10 Preparation of B r0 3F 203 5.5.11 Reactions involving H S 0 3F 204 5.5.12 Attempted synthesis of BrOF 3 205 5.5.13 Attempted synthesis of B r0 2F 206 5.5.14 Attempted synthesis of K [B r0 2 F2] and K[BrOF4] 206 Sources of Chemicals and Methods of Purification References 208 211 viii List of Tables 1.1 The oxides, fluorides and oxide fluorides of xenon 2 2.1 Seflate derivatives of the main group elements 13 2.2 Seflate derivatives of the transition metals 14 2.3 Proton-1 NMR chemical shifts for CH3X and CH 2 X 2, 17 X = halogen or seflate 2.4 Aab values for seflate compounds 23 2.5 Solvent effects on the value of R for [Ti(OTeF5)4] 24 2.6 The dependance of v(Se-O) on covalent or ionic character 25 2.7 Vibrational modes of the seflate group 26 2.8 Bond angles for Xe(OSeF 5 ) 2 33 2.9 EXAFS and crystal data for Xe(OSeF 5 ) 2 36 2.10 EXAFS data for [Re(CO)5 (OSeF5)] 44 2.11 EXAFS data for [Mn(CO)5 (OSeF5)] 51 2.12 Fluorine-19 NMR spectral data for the products of 59 the reaction between I 2 and five molar equivalents of Xe(OSeF 5 ) 2 2.13 Fluorine-19 NMR spectral data for the products of 65 the reaction between I 2 and three molar equivalents of Xe(OSeF 5 ) 2 2.14 A comparison of the v(CO), v(Se-O) and v(Te-O) 6 6 values for various carbonyl derivatives 2.15 The comparative chemistry of XeL2, L = fluoride, 6 8 seflate or teflate 3.1 Structures of the known and possible oxide fluoride compounds of bromine (V) and bromine (VII) ix 76 3.2 Standard electrode potentials (in acid solution) between highest oxidation states of non metals 3.3 EXAFS and X-ray crystal data for K[BrF4] and Cs[BrF6], [BrF 2 ][AsF6] and [BrF 4 ][Sb 2 F n ] 3.4 EXAFS data for Cs[BrOF4] 4.1 Physical properties of H 2 S 0 4, H S 0 3 F, S 0 2 F2, CF 3 SO 3 H and HF 4.2 Reaction times and temperatures involved in the formation of fluorosulphate derivatives 4.3 Infrared vibrational data and assignments for K [S 0 3 F] 4.4 Infrared vibrational data and assignments for [C o(S 0 3 F)2], [Fe(S0 3 F)2] and [N i(S0 3 F)2] 4.5 The Raman vibrational data and assignments for K [B r(S 0 3 F)4] and K [I(S 0 3 F)4] 4.6 Infrared vibrational data for the -S 0 3F group in [Fe(S0 3 F)3], [Sn(S0 3 F) 2 Me2], [S n(S 0 3 F) 2 Cl2], K [B r(S 0 3 F)4] and K [S 0 3 F] 4.7 Comparison and assignment of the infrared vibrational data for K [S 0 3 F] and [R e(S0 3 F )(C 0)5] 4.8 Bond lengths and angles for C s[S 0 3 F] 4.9 Bond lengths and angles for C s[A u(S0 3 F)4] and C s[S b(S0 3 F)6] 4.10 Infrared spectroscopic data for K [S 0 3 F] and [F e(S 0 3 F)2] 4.11 Infrared vibrational data for [Re(CO)5 (S 0 3 F)] 4.12 Infrared spectroscopic data for [Cp2 T i(S 0 3 F)2] and K [B r(S 0 3 F)4] 4.13 Proton-1 NMR chemical shifts for [Cp2 TiX2] 176 (X = -S 0 3 F, -OTeF5, -F and -Cl) 4.14 Fluorine-19 NMR chemical shifts for various covalent monodentate fluorosulphate complexes xi 177 List of Figures 2.1 Teflate derivatives 12 2.2 Resonance canonical forms of the seflate anion 25 2.3 The gas phase structure of F 5 SeOSeF 5 28 2.4 The X-ray crystal structure of Xe(OSeF 5 ) 2 32 2.5 Fuorine-19 NM R spectrum of Xe(OSeF 5 ) 2 35 2.6 Background-subtracted EXAFS and the Fourier 37 transform spectra for Xe(OSeF 5 ) 2 2.7 Fluorine-19 NM R spectrum for the product of the 41 reaction between [Re2 (CO)10] and Xe(OSeF 5 ) 2 2.8 Carbon-13 NM R spectrum for the product of the 42 reaction between [Re2 (CO)10] and Xe(OSeF 5 ) 2 2.9 Electron-impact and accurate mass spectrum for 43 [Re(CO) 5 (OSeF5)] 2.10 Background-subtracted EXAFS and the Fourier 45 transform spectra for [Re(CO) 5 (OSeF5)] 2.11 Fluorine-19 NM R spectrum for the product of the 49 reaction between [Mn 2 (CO)10] and Xe(OSeF 5 ) 2 2.12 C arbon-13 NM R spectrum for the product of the 50 reaction between [Mn 2 (CO)10] and Xe(OSeF 5 ) 2 2.13 B ackground-subtracted EXAFS and the Fourier 52 transform spectra for [Mn(CO)5 (OSeF5)] 2.14 Fluorine-19 NM R spectrum for the products of the 55 reaction between [Ru(CO)3 (PPh3)2] and Xe(OSeF 5 ) 2 2.15 Fluorine-19 NM R spectrum for the products of the reaction between [Ru(CO)3 (PPh3)2] and Xe(OSeF 5 ) 2 xii 56 2.16 Phosphorus-31 NMR spectrum for the products of the 56 reaction between [Ru(CO)3 (PPh3)2] and Xe(OSeF 5 ) 2 2.17 Fluorine-19 NMR spectrum for the products of the 60 reaction between I 2 and five molar equivalents of Xe(OSeF 5 ) 2 2.18 Fluorine-19 NM R spectrum for the products of the 61 reaction between I 2 and five molar equivalents of Xe(OSeF 5 ) 2 2.19 Fluorine-19 NM R spectrum for the products o f the 61 reaction between I2 and five molar eqivalents of Xe(OSeF 5 ) 2 2.20 Fluorine-19 NM R spectrum for the products o f the 64 reaction between I 2 and three molar eqivalents of Xe(OSeF 5 ) 2 3.1 The isomeric forms of I 0 2 F 3 87 3.2 Background-subtracted EXAFS and the Fourier 98 transform spectra for K[BrF4] 3.3 Background-subtracted EXAFS and the Fourier 99 transform spectra for Cs[BrF6] 3.4 Background-subtracted EXAFS and the Fourier 100 transform spectra for [BrF 2 ][AsF6] 3.5 Background-subtracted EXAFS and the Fourier 101 transform spectra for [BrF 4 ][Sb 2 F n ] 3.6 Proposed reaction scheme for the reaction between 105 bromine pentafluoride and the alkali metal nitrates 3.7 Background-subtracted EXAFS and the Fourier 108 transform spectra for Cs[BrOF4] 3.8 Fluorine-19 NMR spectrum of B r0 3F in BrF 5 xiii 112 4.1 The protonation of acetamide 126 4.2 Variety of fluorosulphate derivatives 128 4.3 The bonding modes of the fluorosulphate ligand 140 4.4 M olecular structures of [A u(S0 3 F)4]‘ and [Sb(S 0 3 F)6]' 149 4.5 M olecular structure of [A u(S 0 3 F)3] 151 4.6 Crystal structure of m er-[Ir(S 0 3 F) 3 (CO)3] 156 4.7 Proposed structure o f [Ir4 (CO) 1 2 H2]2+ 163 4.8 C arbon-13 and 13C {1 H } NMR spectra of [Ir4 (CO) j2] 164 in H S 0 3F 5.1 M etal vacuum line 187 5.2 M etal reactor 189 5.3 Glass appartatus 189 5.4 Apparatus for the transfer of volatile reagents under 191 static vacuum 5.5 NM R samples fitted inside a 5 mm o.d. precision NM R tube 193 5.6 FEP cell used for the collection of EXAFS data 195 xiv Acknowledgements Firstly, I would like to say thank you to my family. Their help and support over the years has meant a lot to me, and without them, none of this would have been possible. I would also like to take this opportunity to thank my supervisor Dr Eric Hope and Professor John Holloway for their help and guidance throughout my time in the Fluorine Group at Leicester. A very big thank you goes out to Anne Crane. I must also acknowledge D r’s G. Griffith and G. Eaton for their help in recording NM R and mass spectra respectively. Finally, I would also like to say thank you to hundreds of people, but I can’t. So, to all those people whom have touched my life, cheers, you’ve made me what I am. To all the members of the Fluorine Group, past and present, particularly Dr P. Bhattacharyya, and to the rest of the chemistry department in general, thanks. And lastly, but not least, there are several people in particular whom I am grateful too; Dr Lee Peck for being a sound mate! Danny and Lindsey, for numerous relaxing evenings when the work got to much. My family, because they are definitely worth mentioning twice. Raj and Russ for m emorable times. Adam and all the guests of 3 Greenhill Road, quality. Last, but not least, my girlfriend Orla Mary Teresa McLoughlin, thanks for being their through thick and thin, and a big thanks for still loving me even though I ’ve messed up on more than one occasion. xv Abbreviations AHF : anhydrous hydrogen fluoride Cp : cyclopentadienyl (rj5 -C 5 H5) 8 : NM R chemical shift EXAFS : Extended X-ray Absorption Fine Structure FEP : tetrafluoroethylene / perfluoropropylene copolymer Hz : Hertz I R : Infrared K e l-F : poly(chlorotrifluoroethylene) Me : Methyl NM R : nuclear magnetic resonance O. D. : outer diameter, I. D . : internal diameter ppm : parts per million t-Bu : tertiary butyl v : stretching frequency ( d ) : doublet ( d d ) : doublet o f doublets ( t ) : triplet ( q ) : quintet c m '1 : wavenumbers w : weak m : medium s : strong v : very s h : shoulder b r : broad UV : ultra-violet Avi/ 2 : full width half height xvi CHAPTER ONE Introduction 1.1. General Introduction. The ability of oxygen and fluorine atoms to stabilise high and unusual oxidation states is a direct consequence of their very high electronegativities and their very low reduction potentials. The concept of electronegativity was first introduced by Pauling[1] and is defined as the ability of an atom to attract electron density towards itself in a molecule. In order to stabilise high oxidation states, strong covalent bonds are needed to restore electron density to the central atom. The strength of a bond, and its nature, depends on the relative electronegativities of the atoms in the bond. As the difference in the electronegativities of the two atoms in a bond gets smaller, so the nature of the bond shifts from ionic towards covalent. Fluorine and oxygen have high electronegativities, 3.98 and 3.44 respectively. Electronegativity varies with size, nuclear charge and, more importantly, oxidation state of an element. As one descends a group, the electronegativity of the neutral element decreases. However, as the oxidation state of an element increases, shielding of the nuclear charge becomes less effective and the element becomes less polarisiable or more electronegative. As a consequence, a large majority of the highest oxidation state compounds of the elements of Groups 16, 17 and 18 are oxides, fluorides and oxide fluorides. In general, to stabilise these highest oxidation states it is usually necessary to replace fluorine with oxygen atoms, e.g. Table 1.1,[2] which highlights that XeF 8 is unknown, whereas, X e 0 3 F 2 and X e 0 4 have been isolated. It also appears that, especially with respect to the transition metals, the substitution of fluorine for oxygen atoms tends to destabilise lower oxidation states, e.g. the lowest oxidation states of the oxide fluorides of Cr and Mn are V and VII respectively / 33 whereas, those of the lowest binary fluorides are II and III (CrOF 3 and M n 0 3 F, CrF 2 and M nF3). The electronegativities of the halides decrease down the group (cf. F (3.98) > Cl (3.16) > Br (2.96) > I (2.66)), so that fluorine, a first row element, 1 has an electronegativity considerably higher than that of chlorine, bromine or iodine. These changes in electronegativity are evident in the halide chemistry of Group 16: SF6, SeF 6 and TeF 6 are all stable molecular covalent species,1[4] whereas, the highest oxidation state chlorides are SC12, SeCl4 and TeCl4. This trend is even more marked for the bromides and iodides, and is a result of the fact that the heavier halides become progressively more easy to oxidise and are therefore less able to stabilise high oxidation states. Table 1.1. The oxides, fluorides and oxide fluorides of xenon. Oxidation Fluoride Oxide Oxide fluoride state II - XeF 2 - IV - XeF 4 XeOF 2 VI X e03 XeF 6 XeOF 4 X e 0 2 F2 VIII X e04 - X e 0 3 F2 As will be looked at in Section 2.4, polyatomic ligands such as -OSeF 5 are known to stabilise high and unusual oxidation states.[5] It appears that the accumulation of five fluorines around the selenium atom produces an extreme electron deficiency at selenium, which extends as far as the oxygen atom. In this instance, and because fluorine is normally able to restore electron density via n bonding, the electronegativity of the oxygen atom may exceed that of fluorine. 2 1.2. Characterisation. High oxidation state fluorides and oxide fluorides are normally very moisture-sensitive, highly corrosive and volatile; properties which are not conducive to obtaining reliable structural data. Nevertheless, electron diffraction and microwave spectroscopy have been successfully used to structurally characterise a range of gaseous non-metal fluorides and oxide fluorides, e.g. Br 0 3 F [6] (pseudo tetrahedral: d(Br-C>) = 1.582(1) 1.708(3) A, A, d(Br-F) = ZO B rO = 114.9(3)° and ZO B rF = 103.3(3)°) using electron diffraction, and SeOF2[7] (trigonal: </(Se-0 ) = 1.576(3) A, d(Se-F) = 1.729(1) A, Z FS eF = 92.22(10)° and ZO SeF = 104.82(1)°) using microwave spectroscopy. X-ray crystallography is the definitive method for structurally characterising solids, however, the properties of these materials tends to result in them not meeting the prerequisite for good quality crystals required by this technique. Nevertheless, there has been success in the structure determination of [NMe 4 ]+[IOF6] '[8] (pseudo pentagonal-bipyramidal: <7(1-0) = 1.775(6) d(I-F) mean = 1.854(9) A and A). Fluorine-19 NMR and infrared spectroscopies are powerful techniques when dealing with these types of molecules. Chapters Two and Four take a detailed look at these techniques which, when applied to seflate, -OSeF5, and fluorosulphate, -S 0 3 F, derivatives, offer the principal means of characterisation. EXAFS spectroscopy (Section 1.2.1) can provide intem uclear distances for very unstable materials. Although such data is only one dimensional, when combined with the above spectroscopic information, it can produce an essentially complete local structural characterisation. This approach has proven very successful, and is capable of distinguishing between M = 0 and M -F , 191 e.g. M n 0 3F (rf(Mn-O) = 1.59(2) C r0 2 F 2 W(Cr-O) = 1.55(2) A and </(Cr-F) = 3 A and d(Mn-F) = 1.72(2) 1.71(2) A). A) and 1.2.1. EXAFS spectroscopy. The development of synchrotron radiation sources[10] such as those at the Daresbury Synchrotron Radiation Laboratory (CCLRC) has provided experimenters with X-ray sources several orders of magnitude brighter than those previously obtained from conventional X-ray tubes. The level of understanding of extended X-ray absorption fine structure (EXAFS) spectroscopy has advanced such that reliable structural information can be extracted from X-ray absorption spectra . [ U 1 Additionally, the application of EXAFS spectroscopy does not require compounds to be crystalline and can provide structural information on powders, unstable materials, solutions and at different temperatures. A typical X-ray absorption spectrum exhibits decreasing absorption as the photon energy is increased. Superimposed on this smooth background is a sequence of steeply rising discontinuities in the absorption at energies characteristic of each element in the sample. These abrupt increases in absorption occur whenever the incident photon has sufficient energy to promote a core electron to unoccupied valence levels or to the continuum. The edges are labelled according to the core electron being promoted, K edge arises from Is excitation, L edges arise from 2s or 2p excitation and so on. With the discovery of absorption edges came the observation that the absorption near the edge and beyond does not vary smoothly, rather, there is a wealth of fine structural information which is characteristic of the chemical environment of the X-ray absorbing atom. A typical absorption edge consists of a series of approximately Lorentzian lines superimposed on a steeply rising absorption step. W ithin about 25 eV of the absorption edge most of the structure is due to bound state transitions. However, additional structural information is observed over several hundred electron volts past the edge. This long range oscillation, EXAFS, is considered to result from interference between the atom and the photoelectron 4 wave propagating from the X-ray absorbing atom and the wave backscattered by neighbouring atoms. The absorption process may be viewed as a oneelectron transition from a highly localised core orbital to a delocalised continuum state, which is sensitive to the immediate environment of the absorbing atom. Analysis of the positions and relative intensities of the absorption edge features can reveal details about the metal site symmetry, its oxidation state and the nature of the surrounding ligands. More importantly here, interpretation of the phase, amplitude and frequency of the EXAFS oscillations can provide information about the type, number and distances of atoms in the vicinity of the absorber. In an absorption experiment, ionisation gas detectors are mounted in front and behind the sample, and the relative absorbance is obtained by taking the log of the ratio of the currents in each detector. The absorbance of the particular element of interest is superimposed on both the spectrometer baseline and the background absorption (due to cell windows, solvent, air and other elements present in the sample). The phenomenon known as EXAFS, %, is simply the relative modulation of the absorption coefficient p, of a particular atom compared with the smooth background absorption coefficient j l l s, normalised by the absorption coefficient p 0 that would be observed for the free atom (Eqn. 1.1). X = (p - Ps)/M« Ecln - L 1 - EXAFS results from interference between the out-going photoelectron wave from the absorbing atom, and the back scattered waves of the surrounding atoms. Theoretical determination of EXAFS rests on the ability to calculate the relative phases and amplitudes of the out-going and the back-scattered photoelectron waves. In order to interpret an EXAFS spectrum, it is necessary to subtract the background. This is performed using the program EX,[12] 5 developed at The University of Leicester. The Fourier transform of the EXAFS from k (k = photoelectron wave vector) space to R (distance) space provides information about radial distribution. The Fourier transform of a data set from which the background has not been correctly subtracted usually results in large o peaks below 1 A. Curve fitting analysis is carried out to derive a parameterised function that will model the observed EXAFS, and then iterate the structure dependent parameters in this theoretical EXAFS spectrum until the fit with the experimentally observed EXAFS is optimised. This is achieved using the program EXCURV92.[13] The final values of the optimised EXAFS should yield structural information about the compound. A number of variable parameters exist in the program and these include AFAC, which is the proportion of electrons which perform the EXAFS type scatter, and VPI, which takes into account inelastic losses and the core hole lifetime. These values should be comparable for similar types of species. EXAFS spectroscopy only gives one dimensional information except when measured for single crystals. However, the sensitivity of the technique is very high and this characteristic has made it of unique value in the study of metal-containing biological systems. Overall, EXAFS spectroscopy is an invaluable technique especially with reference to the work under taken in this thesis. Although not as accurate as other structural techniques, distances can be obtained with accuracies of up to ± 0 .0 1 A, which is excellent considering that, for the compounds studied in the present work, such information may not be obtainable by other techniques. 6 1.3. Summary. The work undertaken in this thesis is concerned with some of the oxide fluoride chemistry of sulphur, selenium and bromine. The simple oxide fluorides of sulphur and selenium (SOF2, S 0 2 F 2 and SOF4, and SeOF2, S e 0 2 F 2 and SeOF4) are well known.[3] In addition, there are a whole host of complex oxide fluorides which, in the case of sulphur, fall into two categories: those which contain the -S 0 3F group as a structural unit e.g. the series of polysulphuryl difluorides S 2 0 5 F 2 - S7 O 2 0 F2, and those whose structural group is -SF 5 e.g. SF 5 OF, (SF5)20 and (SF 5 O) 2. For selenium, a series of complex oxide fluorides is known e.g. F 5 SeOF, (SeF5)20 and (FSeO)2, however, the chemistry is not as diverse as that observed for sulphur. None of the simple oxide fluorides of tellurium have been isolated, although, a number of complex oxide fluorides are known,[3] e.g. (F 5 Te)20 and (FTeO)2. This chemistry reflects the increased size of the tellurium atom which leads to an increased coordination number, hexavalent tellurium usually attaining a coordination number of six. The work undertaken here was designed to attempt to expand the number of derivatives and exploit new synthetic pathways to complexes of the S (IV) (-SO 3 F) and Se (IV) (-OSeF5) fluoroanions. In contrast, the oxide fluoride chemistry of bromine is not so extensive,[6] and the aim was to attempt to establish pathways to new bromine oxide fluorides, the properties of which it was hoped would lead to new areas of coordination and reaction chemistry. Chapters Two and Four describe the synthesis of novel low-valent metal derivatives containing the high-valent ligands -SO 3 F and -OSeF5. The two ligands have been described as “pseudo fluorides”, and indeed, the high valent complexes [Sb(S 0 3 F)6]' and [I(OSeF5)5] have few analogues besides their respective fluoride derivatives, [SbF6]' and IF5. However, as will be shown, in the area of low valent transition metal derivatives the properties of the -S 0 3F and -OSeF 5 ligands are quite different from that of F', making the term pseudo fluoride inappropriate. Chapters Two and Four begin with reviews of the areas 7 of interest, and cover the history, synthetic approaches, limitations and a detailed look at the respective spectroscopic techniques needed to characterise these type of species. Chapter Three is directed towards the isolation and characterisation of the bromine oxide fluorides. These compounds are of fundamental importance as textbook examples of rare, unusual and discrete molecular geometries. The introduction consists of a review of halogen oxide fluoride chemistry and serves to highlight the corresponding dearth of bromine oxide fluorides relative to the respective chlorine and iodine analogues. 8 References Chapter One [1] L. Pauling, The Nature o f the Chemical Bond, Ithaca, New York, 3rd edn., 1960, ch. 3, 64-107. [2] F. A. Cotton and G. Wilkinson, Advanced Inorganic Chemistry, 5th edn., 1988, ch. 15. [3] J. H. Holloway and D. Laycock, Adv. Inorg. Chem. Radiochem., 1983, 2 7,157 and references cited therein. [4] N. C. Norman, Periodicity and the p-Block Elements, ed. J. Evans, Oxford, Oxford, 1st edn., 1994, ch. 5, 57-71. [5] K. Seppelt, Angew. Chem., Int. Ed. Engl., 1982, 21, 877. [6 ] R. J. Gillespie and P. H. Spekkens, Isr. J. Chem., 1978,17,11 and references cited therein. [7] I. C. Bowater, R. D. Brown and F. R. Burden, J. Mol. Spectrosc., 1968, 28,461. [8 ] A. Mahjoub and K. Seppelt, J. Chem. Soc., Chem. Commun., 1991, 840. [9] W. Levason, J. S. Ogden, A. K. Saad, N. A. Young, A. K. Brisdon, P. J. Holliman, J. H. Holloway and E. G. Hope, J. Fluorine Chem., 1991, 53, 43. [10] E. A. V. Ebsworth, D. W. H. Rankin and S. Cradock, Structural M ethods in Inorganic Chemistry, Blackwell, Oxford, 2nd edn., 1991, ch. 1, p. 15. [11] E. A. V. Ebsworth, D. W. H. Rankin and S. Cradock, Structural M ethods in Inorganic Chemistry, Blackwell, Oxford, 2nd edn., 1991, ch. 8 , 366-371 and references cited therein. [12] EX, A. K. Brisdon, University of Leicester, 1992. [13] EXCURV92, SERC Daresbury Laboratory Program, N. Binstead, J. W. Campbell and S. J. Gurman, 1992. 9 CHAPTER TWO Oxidation Reactions using Xenon 5is(seflate) 2.1. Introduction. The serendipitous discovery of pentafluoroorthoselenic (VI) acid, HOSeF5, colloquially known as seflic acid, was made by S e p p e lt^ in 1972. This followed the similarly unexpected synthesis of pentafluoroorthotelluric (VI) acid, HOTeF5, (teflic acid) by E ngelbrecht^ and Sladky in 1964. Engelbrecht et al. intended to synthesise T e 0 2 F 2 by combining B a[T e04] and H S 0 3F at 160°C, in a reaction analogous to that used in the preparation of S e 0 2 F2. Tellurium dioxide difluoride, T e 0 2 F2, still unknown today, was not isolated, but instead, HOTeF 5 was formed. This was later explained in terms of the tendency of hexavalent tellurium to achieve a co-ordination number of 6 . Generally, compounds of the third row of non-metals resemble the second row o f non-metals; whereas an increase in coordination number is observed on going to the fourth row, e.g. [SO 4 ]2', [S e04]2" cf. [H4 T e 0 6]2". Seppelt was attempting to synthesis SeOF 4 when he discovered HOSeF5. The methods he employed were the same as those used to synthesis the analogous SOF4. The fluorination of SeOF 2 using a range of halogen fluorides or elemental fluorine produced SeF 4 and SeF 6 respectively. However, the fluorination of SeOF 2 in the presence of HF afforded HOSeF 5 (Eqn. 2.1). Selenium oxide tetrafluoride has since been prepared by the vacuum pyrolysis of Na[OSeF 5 ] J 3 1 It is five coordinate and is the only known example of hexavalent selenium. As befits this unusual coordination number, it is unstable above -100°C, forming highly viscous polymeric products. SeOF 2 + F 2 + HF HOSeF 5 Eqn. 2.1. The sulphur analogue of seflate and teflate can be p r e p a r e d , b u t below -60°C it is kinetically stable to reducing the coordination about the sulphur. 10 Above this temperature it readily releases HF, and consequently has found no synthetic use to date. Both HOTeF 5 and HOSeF 5 are strong Bronsted Lowry acids, and studies have shown the acid dissociation constant of HOTeF 5 to be between that of concentrated hydrochloric and nitric acids with a pK& = 9 .2 .^ Further work by Engelbrecht et al. demonstrated that HOTeF 5 could be readily converted to salts , [ 6 , 7 1 opening the door to a whole new area of coordination chemistry; and teflate derivatives are now known for most elements (Figure 2.1). The ability of the teflate ligand, [TeF 5 0 ] ', to stabilise both high and low oxidation state compounds is unsurpassed by any other polyatomic ligand, and the extremes of its covalent and ionic bonding are evident in the compounds Xe(OTeF 5 ) 6 and [Mn(CO) 5 (OTeF5)]. The wide variety of teflate chemistry is a consequence of the high stability of the group and its ease of introduction into a complex. Several synthetic approaches e x is t^ and include the use of, i) chlorides, fluorides or methyl compounds, which undergo displacement reactions with teflic acid to produce the corresponding teflate derivative, and HC1, HF or CH 4 respectively ii) boron tris(teflate), B(OTeF5)3, which undergoes metathesis to generate BF 3 iii) xenon bis(teflate), Xe(OTeF5)2, which is an extremely strong oxidising reagent, and provides a route into teflate derivatives which does not rely on the replacement of a group as the driving force iv) silver teflate, [Ag(OTeF5)]2, and mercury bis(teflate), Hg(OTeF5)2, which undergo displacement reactions with chloride derivatives. (However, these latter reagents have found only limited use owing to the difficulties involved in their preparation, and the separation of the products.) In marked contrast, the high oxidation potential of the Se (VI) centre in seflic acid has limited its use as a reagent (Tables 2.1 and 2.2); the only compounds which are compatible with seflic acid being fluorides and oxides. This can be partially overcome by the use of [Hg(OSeF5)2], itself prepared from HgF 2 and HOSeF5, which reacts gently with chlorides in inert solvents. In view of the dearth of metal-seflate species and the similarity between the iso- 11 Figure 2.1. Teflate derivatives. Low Valent Anionic [Re(CO)5(OTeF5)]t8l Cs[OTeF5][u ] [Mn(CO)5(OTeF5)[9’10l [Sb(OTeF5)6]‘ti2] Teflate High Valent Homoleptic B(OTeF5)3[13’14] C(OTeF5)4[15>16>i7] Organometallic [Cp2Ti(OTeF5)2][8] [CpFe(OTeF5)2]t8] Xe(OTeFsy i8’19l U(OTeF5)6t2°l Table 2.1. Seflate derivatives of the main group elements. Group 1 Group 15 Group 17 Li[OSeF 5 ] [ 1 0 >2 1 ’2 2 1 [N 0 2 ][0S eF 5] [23) F(OSeF5)[26] Na[OSeF5] [ia21'22] [POF 2 ][OSeF5] [24) Cl(OSeF 5 ) [ 2 6 -2 7 1 K[OSeF 5 ] [ 1 0 ’2 1 ’2 2 1 As(OSeF5)3[25] Br(OSeF5)[26] Rb[OSeF 5 ] [ 2 1 ' 2 3 1 Cs[OSeF5] [21'23] Br(OSeF5)3[26] Group 16 Rb[Br(OSeF 5 ) 4 ] [ 2 3 -2 6 1 (OSeF5)2[26,27] F,I(OSeF 5 )5 J 17'28] ( x = 1 -5 ) Group 14 F 2 Se(OSeF5)2[34l C(OSeF5)4[29] F S 0 2 - 0 - S 0 2 -0S eF 5[31] Group 18 CH^(OSeF5)4. ; 29! SF 5 -OSeF5[35l36] Xe(OSeF5)2[27,37] o C 5 F 5 -OSeF5[30] 0= S e(0S eF 5 ) 2 [ 2 3 1 FXe-OSeF5[37’38] FCO-OSeF5[30] F 5 Se-0-SeF5[35] F5 SeO-Xe-OTeF5I27'38] ( x = 0 -4 ) CF 3 CO-OSeF5[31] (CH 3 )3 Si-OSeF5t32>33] Table 2.2. Seflate derivatives of the transition metals. [F(4-i)Ti( ° s eF5) J 39,40 i Reference ii Compound [0 = V (0 S eF 5)3] 39,40 [ 0 2 Cr(OSeF5)2] 39,40 [Hg(OSeF5)2] 10,21,27 structural -OSeF 5 and -OTeF 5 groups, relevant data on metal-teflate compounds are discussed where appropriate, in this chapter. 2.2. Preparative Routes to Compounds Containing the Seflate Group. Three main synthetic routes have been employed for the introduction of the seflate group into a compound:- i) Acid displacement reactions with seflic acid, HOSeF5. As described above, the high oxidation potential of hexavalent selenium restricts the usefulness of seflic acid, and thus only fluorides and oxides are compatible with this reagent (Eqn. 2.2[37]). XeF 2 + 2 HOSeF 5 -------- ► Xe(OSeF 5 ) 2 + 2 HF Eqn. 2 .2 . The compounds Br(OSeF<;)^2^ and [Hg(OSeF5)2] F1,21’27J have also been prepared from their respective precursors, BrF 3 and HgF2, by this method. 14 ii) The difficulty associated with using HOSeF 5 can be avoided when using [Hg(OSeF5)2], this reacts gently with chlorides (Eqn. 2.3). C H ^ C l^ + Vi Hg(OSeF 5 ) 2 --------► CHx(OSeF5)4.x + V4HgCl2 Eqn. 2.3. (x = 0 - 3 ) By this route [CrO 2 (OSeF5)2],[39'40] [VO(OSeF5)3],[39'40] C(OSeF5)4[29] and As(OSeF5)3^29' have been prepared from C r0 2 Cl2, VOCl3, CC14 and AsC13 respectively. iii) The use of Xe(OSeF 5 ) 2 as a clean reagent for the introduction of the seflate group into molecules has not been exploited at all. The only compounds produced and studied are thermal and photochemical decomposition products^35] (Eqn.'s 2.4, 2.5 and 2.6). UV F 5 SeOXeOSeF 5 — — ► 2 OSeF 5 -------- ► F 5 SeOOSeF 5 ” X Eqn. 2.4. G F 5 SeOOSeF 5 A » F,SeO X eO SeF, ---- ► 130 C F 5 SeOSeF 5 + SeF 4 + SeF 6 + 0 Xe + F 5 SeOSeF 5 2 Eqn. 2.5. Eqn. 2.6. Other reagents have found isolated uses in the preparation of seflate species and these include F 2 OPOSeF5^17,28^ and C10SeF5^28^ (Eqn.’s 2.7 and 2 .8 respectively). IF 5 + F 2 OPOSeF 5 -------- ► Fxl(OSeF5)5.x + POF 3 Eqn. 2.7. (x = 0 - 5 ) IC13 + C!OSeF 5 ► I(OSeF 5 ) 3 + Cl2 Eqn. 2.8. 2.3. Stability of Seflate Compounds. In general, seflate compounds must be stored in an inert atmosphere dry box, and reactions must be performed in dry prepassivated vessels using rigorously dried solvents. Seflate compounds may undergo various decomposition reactions. Loss of the whole ligand, with the formation of the resulting peroxide,[26] is common where the bond to the central atom is weak (Eqn. 2.9). Xe(OSeF 5 ) 2 Xe + (F 5 SeO ) 2 Eqn. 2.9. The elimination of selenium oxide tetrafluoridej39,40^ 0= S eF 4, from a coordinated seflate group is encountered when the -OSeF 5 group is attached to very electropositive, or coordinatively unsaturated central atoms (Eqn. 2.10). Ti(OSeF 5 ) 4 -> FJCTi(OSeF5)4_;c + 0=S eF 4 Eqn. 2.10. (x = 0 - 4 ) The elimination of oxygen is rare for all chalcogen pentafluorooxo species and the only examples known are shown in Equations 2.11,^27^ and 2 . 12.[31] 2 F 5 SeOOSeF 5 2 F 5 S 0 -0 -0 S F 5 -> -> F 5 SeOSeF 5 + 0 2 F 5 SOOSF 5 + 0 2 16 Eqn. 2.11. Eqn. 2.12. 2.4. Electronegativity of the Seflate Anion. The seflate group resembles fluorine in its ability to stabilise unusual oxidation states. For example, the compounds I(OSeF5)5, Xe(OSeF 5 ) 2 and Br(OSeF5)3, have few analogues other than their related fluorides. The high stability of xenon compounds, such as Xe(OTeF5)6, raised the question of the group's electronegativity. However, electronegativity has no simple definition and becomes even more ambiguous when applied to a group. In efforts to determine the group electronegativity of seflate relative to that of fluorine, a number of investigations have been carried out but with contrasting results, for example: - i) The NM R chemical shifts of the methyl protons of CH 3 X, X = I, Br, Cl, F and OSeF5^29J are presented in Table 2.3. By extrapolation o f a plot of the electronegativities of the halogens vs. 8 ( 1 H), it was concluded that the seflate group has an electronegativity of ~ 4.1 on the Pauling scale which is higher than that for fluorine (3.98). A similar result was also obtained for the di substituted methyl complex, CH 2 (OSeF5)2. Table 2.3. Proton-1 NMR chemical shiftsa for CH3X and CH 2 X 2, X = halogen or seflate. CH 3 (OSeF5) 4.50 CH 2 (OSeF 5 ) 2 6.30 c h 3f 4.26 c h 2f 2 5.54 c h 3c i 3.05 c h 2 c i2 5.33 CH3Br 2 .6 8 CH 2 Br 2 4.94 c h 3i 2.19 c h 2 i2 3.90 a ppm relative to TMS. 17 ii) The reaction of IF 5 with F 2 OPOSeF 5 leads to substitution of the fluorine atoms on the iodine by seflate^18,28^ (Eqn. 2.13). IF 5 + F 2 OPOSeF 5 -> FxI(OSeF5)5.^ + POF 3 Eqn. 2.13. (x = 0 - 5 ) Valence shell electron pair repulsion theory (VSEPR) predicts that in a square-based pyramid the axial position is always occupied by the least electronegative ligand. When the above reaction was monitored using 19F NMR spectroscopy, it was noted that the only reaction when an axial seflate ligand was observed was when complete substitution occurred, i.e. the formation of I(OSeF5)5. The behaviour of the seflate ligand cannot be explained kinetically: longer reaction times and heating do not alter the results. Hence, in accord with VSEPR theory, the order of the substitution indicates that the seflate group possesses an electronegativity higher than that of fluorine. Iodine pentateflate, I(OTeF5)5, cannot be prepared as outlined above, but may be synthesised via a different route (Eqn.’s 2.14 and 2.15). IF 3 + B(OTeF 5 ) 3 -> I(OTeF 5 ) 3 + BF 3 I(OTeF 5 ) 3 + Xe(OTeF 5 ) 2 -> I(OTeF 5 ) 5 + Xe Eqn. 2.14. Eqn. 2.15. iii) The ligand properties of -OSeF 5 and -OTeF 5 groups in the pseudo-trigonalbipyramidal molecules F 2 Se(OSeF5)2, F 2 Se(OTeF 5 ) 2 and F 2 Te(OTeF 5 ) 2 have also been investigated^34^ by 77Se and 125Te NMR spectroscopy. All three compounds possess axial -OSeF 5 or -OTeF 5 ligands, with the fluorine ligands in the equatorial plane. VSEPR theory states that the axial position of a trigonal-bipyramidal molecule is occupied by the more electronegative ligand; therefore, this indicates that both seflate and teflate ligands possess a higher electronegativity than that of fluorine. 18 iv) A correlation of the 31P NMR chemical shifts and P = 0 stretching frequencies for POF 2 -X, X = Cl, F or OSeF 5 has indicated[28] that seflate and fluorine ligands have approximately equal electronegativities. The evidence presented for seflate complexes closely matches that obtained for the corresponding teflate systems. However, the teflate anion has been more extensively studied. Schrobilgen et a l synthesised a series of teflate compounds of Te, I and Xe and their fluorine analogues, and studied them using 125Te and 129Xe NMR and 127I and 129Xe Mossbauer spectroscopies J 41^ The NM R chemical shifts and Mossbauer quadrupole splittings of the central Te, I and Xe atoms were used to assess the relative electronegativities of fluorine and teflate ligands. In 129Xe and 125Te NMR experiments the chemical shift range is exceedingly large: ~ 7500 ppm^42^ and 3000 ppm ^ respectively. Consequently, the chemical shifts are very sensitive to changes in electron density at the xenon or tellurium nucleus. A comparison of the 129Xe and 125Te NMR chemical shifts for the above series of compounds, revealed that the fluoride species were significantly more deshielded than their teflate analogues, implying that fluoride has a greater electronegativity than that of teflate. In the M ossbauer spectroscopic studies, isomer-shift differences between fluorine and teflate containing complexes were investigated. However, results proved inconclusive as the differences were within experimental error. The Mossbauer quadrupole splittings recorded for the central xenon, iodine and tellurium atoms on the other hand, established that fluorine was more electronegative than the teflate group: the latter being given a value of 3.87 (Pauling’s scale), compared with that of 3.98 for fluorine. O f the studies carried out, those performed by Schrobilgen appear to be the most definitive, and indicate 19 that fluorine possesses a greater electronegativity than the teflate group. Although a similar study has not been carried out on the seflate group, a similar result may be anticipated. In the case of the trigonal bipyramidal species, the axial and equatorial regions of space are clearly geometrically different and there appears to be no exception to the VSEPR rules. In an effort to explain why fluorine occupies the axial position Schrobilgen s u g g e s te d ^ that the fluorine atoms, because of their high electronegativity, have less electron density close to the central atom than other ligands. In the case of the seflate ligand, electron density on the oxygen atom may be significantly diminished by the interaction of the non bonding electron pairs with the ligand group -SeF5. This pn-dn interaction is important mainly in systems of very high oxidation states. Hence, although fluorine is a more electronegative element than oxygen, fluorine needs more space for its non-bonding pairs of electrons. The square-based pyramidal geometry of I(OSeF 5 ) 5 may be regarded as a pseudo octahedron. In such a case, the differences between the equatorial and axial positions become more subtle, making predictions more difficult. Furthermore, there are a few examples in main group, transition metal and actinide chemistry were the axial position of a pseudo octahedron is occupied by the more electronegative ligand. For the iodine dioxide tetrafluoride anion, [IO 2 F 4 ]', both cis and trans isomers are known to e x i s t , ^ despite the fact that oxo ligands normally prefer to adopt a pseudo axial position: doubly bonded oxygens exhibiting steric characteristics^45^ similar to that of a non-bonded pair of electrons. W hilst the precise electronegativity of the seflate group remains unknown, it is undoubtedly high. This is a consequence of the inductive effect of the five fluorine atoms bound to selenium, which is augmented by some pn-d n back bonding between the oxygen and the selenium, the final result being an electronegativity of a similar magnitude to that of fluorine. 20 2.5. Spectroscopic Characterisation of Seflate Compounds. Although seflate derivatives are extremely air- and moisture-sensitive, techniques such as infrared, Raman and multinuclear NMR spectroscopies are convenient methods for characterisation. Seflate derivatives show characteristic spectral fin g e rp rin ts ,^ which are sensitive to the type of bonding and oxidation state of the element to which the seflate is coordinated. The strength of the selenium-oxygen bond is related to the type of bond which exists between the oxygen atom and the rest of the molecule. The degree of ionicity of the seflate ligand can be demonstrated by measuring v(Se-O) in the infrared spectra and 8 19 FAX in the 19F NMR spectra of compounds containing the ligand. Free [OSeF5]' would have the strongest interaction between oxygen and selenium, which would result in a short Se-O bond, a high v(Se-O), and a high frequency 5 1 9 F ax due to deshielding of the axial fluorine. 2.5.1. Fluorine-19 NMR spectroscopy. O f the routine analytical techniques 19F NMR spectroscopy is the most important: the 19F nucleus is a 100% spin Vi with a wide chemical shift range. The seflate group contains two different fluorine environments, the four equatorial fluorine atoms Fe and the unique axial fluorine Fa, giving an AX 4 spin s y s t e m A first order AX 4 pattern is observed for the majority of ionic species, invariably the A part of the spectrum being at higher frequency than the X portion. As the interaction between the seflate and the group to which it is bound increases, that is to say becomes more covalent, so A and X become closer, leading to second-order AB 4 spectra. This is the case for seflic a c i d ^ where 5Fa 75.9 and 8 Fe 66.1 ppm. For F 5 SeO-OSeF5,[27] the A and B 4 parts are nearly coincident, 8 Fa 55.2 and 5Fe 54.4, and the spectrum appears to be a single resonance. If the covalency increases still further, the A part of the 21 spectrum moves to lower frequency than the B part. This is observed for CF 3 CO-OSeF 5 ,^31J 5Fa 61.2 and 5Fe 73.5 ppm. The appearance of the spectrum depends on the ratio, R , [49,50] Qf th e p a. Fe coupling constant, 7(FaFe), to the chemical shift difference, 5(FaFe) (Eqn. 2.16). j? = 7(FaFe) 8(FaFe) Eqn. 2.16. 7(FaFe) = Coupling constant (Hz) between Fa and Fe. 5(FaFe ) = Difference in the chemical shift of 5Fa and 5Fe (Hz). The coupling constants /( F aFe) for seflate compounds are typically 215240 Hz, whilst 5(FaFe) can vary over a large range. This is a direct result of the axial fluorine being sensitive to the nature of bonding of the oxygen to the central atom; the chemical shift of the equatorial fluorines (Fe) changing little for different compounds. Therefore, the parameter R can be used as an indicator of the nature of the bonding. For a spectrum in which the second order nature limits the information available, computer simulation programs ^5 ^ can be used to calculate chemical shifts: /( F aFe) varies little and so R can be found by matching the simulated spectrum with the actual spectrum. The value of R is dependent upon the NMR spectrometer operating frequency. For instance, a seflate species could have a coupling constant, 7(FaFe), of 220 Hz and a difference in chemical shifts, 8 (FaFe), of 8 ppm. Hence, on a machine operating at 300 MHz, R = 0.097. However, at 400 MHz, R = 0.073. 22 It is 8 (FaFe) which provides the information about the bonding which is present in a complex. We define here a new parameter, Aa b , to assess the relative ionicity or covalency of a seflate species (Eqn. 2.17). Aab = S(Fe) - S(Fa) Eqn. 2.17. Table 2.4 shows the calculated Aab values for a range of seflate species. W ithin Table 2.4, Aab becomes more positive from top to bottom and this indicates increasing covalent character. Solvent effects have been shown to profoundly affect the 19F NMR parameters in teflate compounds This is demonstrated in Table 2.5, where solvent effects on [Ti(OTeF5)4] considerably alter the observed chemical shifts. Thus, a degree of caution should be exercised when using Aab for different solvent systems. However, Table 2.4 demonstrates that Aab may be used to infer the nature of the bonding present within a molecule. Table 2.4. Aab values for seflate compounds. Compound Solvent Ref. -36.8 c h 3c n 23 229 -18 Neat 26 69.4 234 -1 1 .1 c fc i3 27,37 54.2 52.1 230 -2 . 1 Neat 26 F 5 SeOOSeF 5 55.2 54.1 230 -0 . 8 Neat 27,26 F 5 SeOCOCF 3 61.2 73.5 211 12.3 Neat 31 F 5 SeOSeF 5 62.7 76.0 226 13.3 Neat 35 F 5 S e 0 S 0 2 0 S 0 2F 57.1 78.1 216 Neat 31 SFa 5Fe J ( FaFe) A ab ppm ppm Hz ppm [N 0 2 ][OSeF5] 108.9 72.1 224 IOSeF 5 92.1 74.0 Xe(OSeF 5 ) 2 80.5 FOSeF 5 23 2 1 Table 2.5. Solvent effects on the value of R for [Ti(OTeF5)4] . Solvent R §Fa SFe J (FaFe) ppm ppm Hz (CH3)2SO -15.0 -33.0 171 0.16 (CH 3 OCH 2 ) 2 -42.5 -47.3 174 0.65 c h 3c n -42.3 -48.8 194 0.54 -40.3 -48.9 187 0.38 Genetron 113 -47.3 -45.2 187 -1.58 CC14 -49.5 -43.3 183 -0.52 c h 3n o 2 2.5.2. Vibrational spectroscopy. The seflate group, [OSeF5]", possesses C4v symmetry for which the following vibrational representation is obtained; r V ib = 4 A | + 2 Bj + B 2 + 4 E All of these modes are Raman active but only the Ay and E vibrations are infrared active. The highest frequency observed is assigned to the Se-O stretch which is in accordance with the partial double bond character^46^ of this type of bond (Figure 2.2). 24 Figure 2.2. Resonance canonical forms of the seflate anion. F \ /F Se, .S e , o- O The Se-O distance and stretching frequency vary in a characteristic and understandable manner. This variation depends on the nature of the element to which the seflate oxygen is bonded, or ion paired, as well as the strength of the interaction. The extremes of covalent and ionic bonding are evident in the molecules [N 0 2 ][0 S e F 5] and F 5 SeOSeF5, which have values of v(Se-O) of 918 and 760 cm - 1 respectively (c f Table 2.6). Table 2.6. The dependence of v(Se-O) on covalent or ionic character. Compound v(Se-O) cm ' 1 [N 0 2 ][0 S eF 5] 918 Ref. 23~~ Xe(OSeF 5 ) 2 787 27,37 F 5 SeOOSeF 5 765 27,36 F 5 SeOSeF 5 760 35 The vibrational modes and assignments expected for the seflate anion are presented in Table 2.7. Mayer and Sladky assigned these modes by comparison of the spectral data for Cs[OTeF5]^11,53^ with those for the isoelectronic C4v species, IOF5. Due to the differences in mass and effective charge of the central atom, most modes are observed at lower frequency when going from [OTeF5]' to IOF5. A similar shift would be expected when going from [OSeF5]‘ to [OTeF5]'. 25 Table 2.7. Vibrational modes of the seflate group. Assignment in C4v Description of vibration point group v (Se - 0 ) V i( A j) v 2 ( A i) ^ sym ( ^ 6 v3 ( A i ) - Feq) v (Se - Fax) v4 ( A j) 8 sym (out-of-plane SeF4) v5 ( B ! ) v sym (out-of-phase SeF4) V6 ( B ! ) 8 a sy m v7 ( B 2 ) 8 (out-of-plane SeF4) sym (in-plane-SeF4) v8 (E ) ^ a sy m v9 ( E ) v io( E ) Vn(E) (ScF4) 8 (F - Se - F4) 8 (O - Se - F4) S asym (in-plane SeF4) 2.5.3. M ass spectrometry. Mass spectrometry can be a particularly useful and informative technique. Selenium has six isotopes which, when coupled with the isotopic distribution of the other elements, can lead to complicated but characteristic patterns. Using computer programs it is possible to simulate the expected isotopic distribution, and these can be used to verify the composition of the species in question. A survey of the literature indicates that the parent ion is rarely observed for seflate-containing species. Loss of an entire group usually yields an intense fragment. The elimination of 0=S eF4, leaving one fluorine behind, is also common. This was what was found for Br(OSeF 5 ) 3 ; [ 2 6 1 no parent ion was 26 observed, but the loss of a seflate group produced [Br(OSeF5)2]+, m/z 461 ( 6 %), and the subsequent loss of 0= S eF 4 produced [FBr(OSeF5)]+, m/z 289 (8%). Of the ionisation techniques available, electron impact has been the most useful to date. This technique requires the sample to be slightly volatile. It is then ionised by an interaction with a beam of electrons to produce a radical cation, [M']+. The drawbacks are that thermal decomposition may occur during the vaporisation of the sample and only a limited mass range is accessible, (<10 3 AMU). Other techniques such as electrospray and fast atom bombardment (FAB), possess an upper mass limit of 9000 AMU and do not require the samples to be volatile. However, these techniques offer no advantages for the characterisation of moisture-sensitive seflate-containing compounds as they require the sample to be solvated in either methanol-water, glycerol or nitrobenzyl alcohol. 2.5.4. X-ray crystallography and EXAFS spectroscopy. W hile single crystal X-ray crystallography offers the ideal method with which to determine molecular structures, the only successful crystal structure determination of a seflate containing compound to date is that of xenon bis(seflate)J2°l Isolating suitable single crystals is the problem. Single crystals of seflate derivatives ought to be best prepared by vacuum sublimation. However, the technique is notorious for the disorder it produces and the problem is enhanced by the spherical shape of the seflate ligands. Even in the absence of systematic disorder, the peripheral fluorine atoms appear with very large vibrational parameters, caused by a combination of molecular vibrations and disorder. This problem can be reduced by performing the experiments at low temperature, but 27 varying the temperature may result in a phase transition or powdering of the crystal. It seemed likely, therefore, that EXAFS spectroscopy might be the ideal technique for the determination of element-element distances as explained in Chapter One. 2.6. Covalent Bonding. An atom in a high oxidation state requires strong covalent bonds to stabilise it. However, the ligands which can do this must possess a high electronegativity, otherwise, a redox reaction will take place. The seflate ligand is able to stabilise high oxidation states and compounds such as I(OSeF 5 ) 5 and Xe(OSeF 5 ) 2 have few analogues outside of fluorine chemistry. Using electron diffraction a structural investigation was carried out on bis(pentafluoroselenium) oxide, F 5 SeOSeF 5 J 35,36^ The structure consists of octahedra linked via an oxide bridge (Figure 2.3). Figure 2.3. The gas phase structure of F 5 SeOSeF5. The gas phase structure of F 5 SeOSeF 5 indicates a large Se-O-Se angle, 142.4°, and an eclipsed conformation of the fluorine atoms. This is sterically unfavourable and a slight twist of the Se-O-Se linkage would certainly reduce the strain. 28 The bridge angle is large and constant (about 143°) for the three chalcogen species F 5 SOSF5, F 5 SeOSeF 5 and F 5 TeOTeF5. This is at variance with the fact that steric interactions between the equatorial fluorines diminishes considerably in the sequence F 5 SOSF 5 > F 5 SeOSeF 5 > F 5 TeOTeF5. Characterisation of F 5 SOSF 5 ^35,36^ shows the equatorial fluorines are bent 2.1° away from the octahedral orientation, but this effect diminishes in F 5 SeOSeF 5 (1.1°) and disappears for F 5 TeOTeF5. Steric interactions would be expected to cause a lengthening of the O-X bond together with an increase in the bond angle. Therefore, delocalisation of the oxygen lone pairs is resulting in a pn-dn contribution to the O-X bond. This is evidenced by a shortening of the O-X distance. The Se-O bond distance of 1.697(13) of a double and a single bond value (Se 0 selenite, (Se-O) = 1.80(2) 2 A for F 5 SeOSeF5, is between that (Se= 0 ) = 1.61(1) A)J35,361 Therefore, A and ethylene on the basis of the short bonds, large E-O-E angle, as well as the sterically unfavourable eclipsed manner of the equatorial fluorines, one can assume a considerable amount of double bond character for the E-O bond in 0 (E F 5)2, (E = chalcogen). The shortening of the Se-O bond may also be explained in terms of hyperconjugation. These resonance modes (Figure 2.2) would give rise to a shortening of the oxygen bond, and a corresponding lengthening of the fluorine bonds, especially the axial bond. However, no lengthening of the fluorine bonds is observed. Thus, pn-dn bonding is favoured as an explanation for the structural character of F 5 SeOSeF5, F 5 SOSF 5 and F 5 TeOTeF5. 2.7. Ionic Bonding. Attempts have been made to isolate alkali group metal teflate salts in order to determine the electronic and molecular properties of the uncoordinated teflate anionJ 5 4 1 Salts such as Cs[OTeF5] [11] and [NBun4 ][OTeF5] [55] were 29 initially put forward as models for the free teflate anion. These exhibit the highest tellurium-oxygen stretching frequencies known, 873 and 867 cm - 1 respectively. However, structural analysis is difficult due to similarities in the covalent and van der Waals' radii of oxygen and fluorine and fluorine-oxygen site disorder. The compound [(PS)H]+[OTeF5]" [(PS)H+ = protonated 1,8- bis(dimethylamino)naphthalene] J 55,56^ was examined by X-ray crystallography. Unlike the other salts of the OTeF5' anion, it does not exhibit any oxygenfluorine disorder. The spectroscopic data, v(Te-O) = 865 cm ' 1 and r(Te-O) = 1.803(3) A, closely match that for [NBun4 ][OTeF5], and it was concluded that this structure contains the best approximation to that of the free OTeF5' anion. This work confirmed that, as the negative charge of the teflate is localised on the oxygen, the tellurium-oxygen bond shortens, the corresponding stretching frequency increases and the 19F NMR chemical shift of the fluorine trans to the oxygen, shifts to higher frequency. In 1984, Strauss et a l successfully made the first low valent transition metal teflate complex [Mn(CO)5 (OTeF5)], by the reaction of [Mn(CO)5 (CH3)] and HOTeF 5 J 9 , 1 0 , 5 7 1 Fluorine-19 NMR and infrared spectral data were consistent with the compound having a considerable degree of ionic character. Single crystal X-ray analysis showed a short Te - 0 distance of 1.751(11) A, which is indicative of Te-O n bonding and reflects the highly ionic character of this species. The staggered confirmation of the OTeF 5 group with respect to the M n(CO ) 5 moiety precludes O-Mn n bonding. Seflate compounds, in accordance with their scarcity, have been less well studied. The closest model to uncoordinated seflate is [N 0 2 ][0 S eF 5 ] , t 2 3 1 for which v(Se-O) = 918 cm - 1 (this compares with v(Te-O) = 848 cm " 1 in [Mn(CO)5 (OTeF5)]) and its 19F NMR spectrum showed 8Fa108.9, 8 Fe 72.1 ppm and Aab -36.8. Some indication of the ionic nature of the bonding present within a seflate derivative can be derived by comparison with this data. 30 2.8. Xenon Bis(seflate). Xenon difluoride in organic solvents has been successfully used to oxidise low-valent transition metal compounds to produce the corresponding metal fluorides. Recent work at Leicester showed that Xe(OTeF 5 ) 2 can be used in a similar fashion to generate low-valent transition metal teflate complexes. By direct analogy with these reactions, we have attempted to use xenon bis(seflate), Xe(OSeF5)2, as a reagent for the introduction of the seriate group into a metal co-ordination sphere. Xenon bis(seflate), Xe(OSeF5)2, was originally prepared according to the following metathetical reaction^37](Eqn. 2.18). XeF 2 + 2H O S eF 5 -» Xe(OSeF 5 ) 2 + 2 H F Eqn. 2.18. This involves the use of seflic acid, HOSeF 5 which is both difficult to prepare and handle. The synthesis of seflic acid^1,27,48^ is based upon the equilibrium reaction shown in Equation 2.19. In accordance with Le Chatelier’s principle the reaction is shifted to the right by removal of the volatile components, HF, S e 0 2 F 2 and HOSeF 5 from the involatile H 2 S e 0 4. 3 S e 0 2 F 2 + 4 HF H 2 S e 0 4 + 2 HOSeF 5 Eqn. 2.19. The seflic acid product is difficult to isolate as a crystalline solid at room temperature, due to HF impurities which are extremely hard to remove. Yields are variable but generally in the region 19 to 6 8 %. A more convenient and cleaner route to xenon bis(seflate) is the oxidation of selenium oxide difluoride, SeOF2, by xenon difluoride[58] (Eqn. 2 .20). 31 2 SeOF2 + 3 XeF2 Xe(OSeF5)2 + 2 Xe Eqn. 2.20. The crystal structure of Xe(OSeF 5 ) 2 has been reported by Templeton et alS20] (Figure 2.4) using crystals grown by sublimation in FEP tubing under dynamic vacuum. The bond angles are listed in Table 2.8. The F-Se-F angles in each seflate group, other than the two constrained to be 180°, are approximately 90°, and thus correspond to a regular octahedral configuration. The 0-Se-F(2) angle deviates by 1 0 ° from linearity, a deviation which although outside the accuracy limits has rather doubtful significance in view of the constrained nature of the model. The reported bond distances are Xe-O = 2.12(5), Se-O = 1.53(5) and Se-F = 1.70(2) A uncorrected for thermal motion, and Se-F = 1.77 A corrected for thermal motion. Figure 2.4. The X-ray crystal structure of xenon bis(seflate). FI' 3 Figure 2.4 represents the dumbbell shaped molecule which packs into a pseudo-rhombohedral unit cell. From the vibrational and NM R spectroscopic data it is evident that the xenon compound is not simply ionic, since the Xe-O distance of 2.12(5) A is less than one would anticipate for a Xe (II) cation [- OSeF5]" anion contact. 32 Table 2.8. Bond angles, (°), for xenon bis(seflate). 180a 92(3) (8 ) F(2)-Se-0 170(2) F(l)-Se-F(2) 88(3) Xe-O-Se 125(2) F(l)-Se-F(3) F (l)-S e -F (l’) 8 8 0 X) F(2)-Se-F(3) 00 r— H O-Xe-O’ F(3)-Se-F(3’) 8 8 (8 ) F (l)-Se-F(3’) 92(8) F(3)-Se-0 95(2) F (l)-S e -0 85(2) - - a By symmetry. b Assumed value. The 129Xe NMR spectrum of xenon bis(seflate)[38] shows nine resonances at 5 129Xe 3131 ppm, 3J (Xe-Fe) = 38; no coupling to the axial fluorines, Fa, was observed. 2.9. Preparation and Properties of Xenon Bis(seflate). S e 0 2 + SF 4 —> SeOF 2 + SOF 2 S e 0 2 + 2 SF 4 —) SeF 4 + 2 SOF 2 Eqn. 2.21. Eqn. 2.22. The systems described in Equations 2.21 and 2.22 are intimately connected and the products formed depend only on the ratio of the starting materials. Thus, if an excess of S e 0 2 is used, SeOF 2 is formed in high yield. This system was utilised to produce the compound SeOF2, which was used as a starting material for the synthesis of Xe(OSeF5)2. The following procedure describes the synthesis. In a typical reaction SF 4 was condensed on to S e 0 2 (molar ratio 0.9:1). The reaction vessel was then sealed and under constant stirring was heated to 120°C for 12 hours. Selenyl fluoride, SeOF2, was the least volatile product and 33 was collected by pumping under dynamic vacuum into a trap cooled to -78°C. Xenon difluoride was loaded into a prepassivated FEP trap and attached to the Monel line. The SeOF 2 was then condensed on to the XeF 2 and, upon warming to room temperature, a steady reaction occurred (Eqn. 2.20) xenon being evolved for around two hours. The mixture was allowed to equilibrate by stirring overnight. The volatile materials were removed at room temperature by pumping under dynamic vacuum for three hours, after which time crystals of Xe(OSeF 5 ) 2 were obtained. Xenon bis(seflate ) [ 2 7 , 3 7 1 is a colourless solid at room temperature. It is extremely moisture sensitive, hence, reactions and storage must be carried out in prepassivated FEP, Kel-F or other fluoroplastic apparatus. Scorching may occur on contact with susceptible materials, and explosive reactions may occur with unsaturated organic solvents. Melting point 69°C Boiling point Thermal stability < 130°C Molecular weight 511 Vapour pressure 0.05 torr @ 0°C 0.35 torr @ 25°C Xenon bis(seflate) is readily characterised by its 19F NMR spectrum (Figure 2.5). Using dichloromethane as the solvent and D20 as the external lock substance a second-order AB 4 pattern was obtained; 5Fa 81.0 ppm, 8 Fe 70.1 ppm, 2 /( F a-Fe) = 234 Hz, V( 7 7 Se-Fa) = 1323 Hz, and V ^ S e -F e ) = 1318 Hz (Figure 2.5). In addition, 129Xe satellites ^5 8 1 were observed for the equatorial fluorines Fe, 3J (Xe-Fe) = 38 Hz. 34 Figure 2.5. Fluorine-19 NMR spectrum of xenon bis(seflate). 64 62 80 78 76 72 70 66 66 Infrared spectra of the solid showed the following bands, and compare well with those published^27,37^ in the literature: 787 (m), 725 (vs), 725 (vs), 700 (s), 612 (s), 550 (m) and 430 (s) c m '1. As was highlighted in Section 2.5.4, it was anticipated that obtaining single crystals of seflate derivatives would be a problem. However, extended X-ray absorption fine structure (EXAFS) spectroscopy does not require the sample to be in a crystalline form and internal bond distances can be readily obtained on powdered samples. To check the suitability of EXAFS spectroscopy for structure analysis, selenium edge EXAFS data were collected for the crystallographically characterised xenon bis(seflate). Transmission selenium K edge EXAFS data were collected out to k = 15 A'1 (k = photoelectron wave vector). This was later truncated to 13.5 A'1 because of increased noise at higher k values. Three data sets were averaged and the data multiplied by k 3 to compensate for a decrease in intensity at higher k. Fourier filtering was not applied and the fit discussed was compared with the average raw (background subtracted) EXAFS data. The data was modelled 35 using EXCURV92t59] to two shells, atom at 3.07(1) A. 6 fluorine atoms at 1.69(1) A and a xenon Each shell was tested for statistical significance EXCURV92 failed to produce reliable data when modelled for 3 shells of 1 oxygen atom, 5 fluorine atoms and a xenon atom. The EXAFS data is presented in Figure 2.6 and Table 2.9. Table 2.9. EXAFS and crystal data for Xe(OSeF5)2. Parameter Mean rf(Se-X) / Af X-ray EXAFSe 1.67(5)g 1.69(1) 1.73(5)h 2o2/ Ab d(Se-X e) / 2o2/Ab A - 0.008(2) 3.24(2) 3.07(1) - 0.014(2) Fit index 0 Rd 2.5 0.064 19.1 a Standard deviations in parentheses. b Debye-Waller factor. c Fit index = Xi[(% T -X E ) k i3 ] 2 . d R = [s(%T-%E)k 3 dk/s%Ek 3 dk] x 100 %. e E0 12.8 (4), AFAC 0.86 and VPI -4.71. f Mean bond length, X = F and O. g Uncorrected for thermal m otion.h Corrected for thermal motion. EXAFS spectroscopy, as highlighted in Section 1.2.1, has been particularly useful for providing structural data on extremely reactive or unstable materials. In particular this approach has proven capable of distinguishing between M = 0 and M-F. However, in the case of Xe(OSeF5)2, the Se-O bond is not expected to possess a large degree double bond character. The crystal data for Xe(OSeF 5 ) 2 (c f Section 2.8) indicates a Se-O bond length of 1.53(5) A which, in the words of the authors “is unrealistically small because o f the constraints imposed by our model.” 36 Figure 2.6. (a) Background-subtracted EXAFS and (b) the Fourier transform spectra for Xe(OSeF5)2. M coX -5 -15 k /A* -‘1 lb] 3.0 C D >. >- 2.0 (0 k. 13 < 1.0 0 -1 r/A aEXAFS ( experimental x k3, bFourier transforms ( curved-wave theory x k3) experimental, — theoretical) 37 In comparison, determined the Se = 0 a microwave spectroscopic distance to be 1.576(3) A, and in study of SeOF2^61^ Section 3.13, the Br = 0 distance for Cs[BrOF4], which should be similar, was calculated to be 1.58(1) A. Therefore, the Se-O distance of 1.53(5) A, determined using X-ray crystallography, is too low, and the inability of EXAFS spectroscopy to distinguish between the oxygen and fluorine atoms is a reflection of the similarities in the size of oxygen and fluorine atoms, and the Se-O and Se-F bond lengths. Therefore, the first shell of six fluorine atoms represents the average Se-F and Se-O bond lengths present with in the seflate ligand. The value of 1.69(1) A is in satisfactory agreement with the crystallographic study where the average Se-O and Se-F bond distances are calculated to be 1.67(5) (uncorrected) and 1.73(5) A (corrected for thermal motion). A covalent interaction between the seflate anion and the xenon atom is expected to result in a large Xe-O-Se angle. As outlined in Section 2.5.1, xenon bis(seflate) definitely possesses some covalent character, Aab = -11.1, c f [N 0 2 ][0 S e F 5] Aab = -36.8 and F 5 SeOSeF 5 Aab = 13.3. Covalent interactions in the case of F 5 XOXF 5 (X = S, Se or Te, Section 2.6) were observed to result in an increase in the oxygen bridging angle. The Xe-Se distance determined using X-ray crystallography was 3.24(2) EXAFS spectroscopy was 3.07(1) A. A, whereas, the value obtained using The bond lengths are significantly different, and the shorter value obtained using EXAFS spectroscopy implies a reduced Xe-O-Se angle of 109(5)° {cf 125(2)° in the X-ray structure). To determine the bond angle it was necessary to assume a Se-O distance of 1.61.69 A, as already discussed the Se-O distance of 1.53(5) A (X-ray structure) is too short. As can be seen the discrepancy in the Xe-Se distance results in a significant decrease in the Xe-O-Se angle. The lack of information makes a detailed discussion inappropriate, however, with the continued expansion of this area and the use of EXAFS, 19F NMR and vibrational spectroscopies, it may in the future, be possible to establish a trend between bridging angles and the degree of covalent nature. 38 2.10. The Reaction Between [Re2 (CO)io] and Xe(OSeF 5 )2 - The reaction of [Re2 (CO)10] with dilute fluorine-nitrogen mixtures in a flow system leads only to the formation of ReF6.[62] Although, the carbonyl halides [Re(CO) 5 X] (X = chlorine, bromine or iodine) are known,![63] w as thought unlikely that fluorine would stabilise the Re (I) oxidation state as it has no available orbitals to permit n back bonding. However, XeF 2 in solution is a mild fluorinating agent and reaction of xenon difluoride, XeF2, with [Re 2 (C O )10] in anhydrous HF or Genetron 113, does lead to the low-valent rhenium fluoride complex [Re(CO)5F ReF5] ^ (Eqn. 2.23). If the Xe and CO are not vented from the reaction then the ionic [Re(CO ) 6 ReF6] is produced. [Re 2 (C O )10] + 3 XeF 2 -> [Re(CO)5F ReF5] + 5 CO + 3 Xe Eqn. 2.23. The reaction between [Re2 (CO)10] and Xe(OTeF 5 ) 2 also leads to a lowvalent metal teflate complex [Re(CO)5 (OTeF5)] ^ (Eqn. 2.24) and the same complex is formed by the reaction of [Re(CO)5 (CH3)] with HOTeF5. [Re 2 (CO )10] + Xe(OTeF 5 ) 2 2 [Re(CO)5 (OTeF5)] + Xe Eqn. 2.24. In an attempt to prepare a Re (I) seflate derivative the reaction between [Re 2 (CO )10] and Xe(OSeF 5 ) 2 was investigated. Colourless [Re 2 (CO )10] was solvated in dichloromethane and then decanted at -78°C on to an equimolar quantity of Xe(OSeF5)2. No immediate reaction occurred but, upon warming to 0°C, a steady reaction commenced and a gas was evolved. Analysis of the gas by infrared spectroscopy showed that no carbon monoxide was present, and thus the gas was presumed to be xenon. The reaction continued for about three 39 minutes during which time the colour of the solution changed to yellow. The volatile materials were removed in vacuo and an orange solid was isolated. The 19F NM R spectrum of this solid was recorded in a 4 mm FEP tube, inside a 5 mm glass NMR tube, using D20 as the external lock substance and dry CH 2 C12 as the solvent. The spectrum showed an AX 4 pattern; 5Fa 98.9 ppm, 5Fe 64.1 ppm, 2 /(F a-Fe) 232 Hz, V(Fe-Se) 1277 Hz, *7(Fa-Se) 1202 Hz and Aab = -34.8 (Figure 2.7). This shows a high frequency shift of the A part of the AX 4 system, and is consistent with a high degree of ionicity in the Re-O bond. The 13C NM R spectrum contained two resonances at 5180.5 and 5178.9 ppm (Figure 2.8). The ratio of the intensities was approximately 4:1 and is consistent with one axial and four equatorial carbonyl groups, however, Tj effects have not been accounted for: metal carbonyls possess long spin-lattice relaxation times. The infrared spectrum was recorded as a Nujol mull of the solid and the following bands were observed:- 2168 w, 2045 s, 1986 w, 856 s, 722 m, s, 6 6 6 6 8 6 sh, 592 s, 555 s, 505 w and 492 s c n r 1. The high v(Se-O) of 856 cm " 1 is exceeded only by those of [N 0 2 ][0 S e F 5] and the alkali metal salts. This compliments the 19F NM R data and indicates a strong Se-O bond, furthermore, this infers a strong ionic interaction between the Re and O atoms. Using group theory to calculate the number of bands to be expected in the carbonyl region of the infra-red spectra of [Re(CO) 5 (OSeF5)], the following irreducible representation is obtained: r*co = 2 A1 + Bj + Ej Only the A] and Ej modes are infra red active, which suggests that 3 bands should be observed in the carbonyl region of the spectrum. Indeed, three bands were found at, 2168, 2045 and 1986 cm '1, consistent with the proposed structure. 40 Figure 2.7. Fluorine-19 NMR spectrum for the products of the reaction between [Re2(CO)10] and Xe(OSeF5)2 (ppm) Figure 2.8. Carbon-13 NMR spectrum for the product of the reaction between [R e^C O )!0] and Xe(OSeF 5 ) 2 182.5 182.0 181.5 181.0 180.5 180.0 179.5 179.0 178.5 178.0 (ppm) The material proved sufficiently stable to obtain an electron-impact mass spectrum. The correct isotope pattern was obtained for the parent ion [Re(CO)5 (OSeF5)]+, m/z 518 (for 185Re and spectrometry was used to 82 Se) (Figure 2.9). Accurate mass unequivocally identify the presence of [Re(CO)5 (OSeF5)], and no fragments derived from loss of CO, F or seflate were observed. The work carried out in Section 2.9 established that the Se edge EXAFS data for Xe(OSeF 5 ) 2 was satisfactory, and final analysis of [Re(CO) 5 (OSeF5)] was therefore attempted by EXAFS spectroscopy. Transmission selenium K edge EXAFS data were collected out to k = 15 A'1 (k = photoelectron wave vector). This was later truncated to 12.5 A-1 because of increased noise at higher k values. Three data sets were averaged and the data multiplied by k 3 to compensate for a decrease in intensity at higher 42 Figure 2.9. Electron-impact and accurate mass spectrum for [Re(CO)5(OSeF5)]. !(X)% = 247462 A D C 518 l(X) — 0 — 455 460 Mass 517.83400 515.83100 515.83480 513.83180 519.83420 513.83670 514.83740 517.83120 Centroid 100-1 480 500 520 540 580 6(X) Abundance 12 C 13 C 16 0 18 0 17 0 19 F 187 Re 185 Re 80 Se 78 Se 82 Se 29.3811 17.6287 13.8708 8.3225 5.4198 5.3195 4.4703 3.2519 5 5 5 5 5 5 5 5 0 0 0 0 0 0 0 0 6 6 6 6 6 6 6 6 0 0 0 0 0 0 0 0 0 0 0 0 0 0 0 0 5 5 5 5 5 5 5 5 1 0 1 0 1 1 1 0 0 1 0 1 0 0 0 1 1 1 0 0 0 0 0 0 0 0 1 1 0 0 0 0 0 0 0 0 1 0 0 1 Llnowldth = 100.0 ppn lOOX = 330225 I n t 50 H 460 470 480 490 500 Mass 510 520 530 540 5 17 .8 34 03 lOOn 908070- s 60i u m 50 40- 30^ 2010517.97534 60 100 120 Channel 43 140 160 180 200 220 240 k. No Fourier filtering was applied, and the fit discussed below was compared with the average raw (background subtracted) EXAFS data. As with the model compound the data was modelled using EXCURV92 to 2 shells of atoms at 1.71(1) A and 1 rhenium atom at 3.55(1) A (Table 2 .1 0 6 fluorine and Figure 2.10). Each shell was tested for statistical significance J 60] In order to obtain all the internal bond distances, EXAFS data were also recorded for the rhenium edge. Rhenium L(IIi) edge EXAFS data were collected for the crystallographically characterised rhenium carbonyl complexes [Re 2 (C O )10] and [Re(CO) 5 Cl], which were used as model systems to test the reliability of data collection and treatment. However, the results were not in satisfactory agreement with the single crystal data. The modelling program EXCURV92^59^ failed to produce reasonable and realistic values for the Re-C and C-O bond lengths, the reasons for which are not understood. As a consequence, Re edge EXAFS data is not reported for the complexes [Re 2 (CO )10], [Re(CO) 5 Cl] and [Re(CO)5 (OSeF5)]. Table 2.10. EXAFS data for [Re(CO)5 (OSeF5)]. Parameter d(Se-X) / l a 2/ Af 1.71(1) Ab rf(Se-Re) / l a 2/ EXAFSe 0.007(2) A 3.55(1) Ab 0 .0 1 2 Fit indexc 4.1 Rd 25.7 (2 ) a Standard deviations in parentheses. b Debye-Waller factor. c Fit index = Xi[(% T-XE)ki3]2. d R = [s(%T-%E)k 3 dk/s%Ek 3 dk] x 100 e E0 5.0 (4), AFAC 0.86 and VPI -4 .7 1 .f Mean bond length (X = F and O). 44 Figure 2.10. (a) Background-subtracted EXAFS and (b) the Fourier transform spectra for [Re(CO)5 (OSeF5)]. -5 -16 (b) 3 2 1 0 O r/ A 1 aEXAFS ( experimental x k3, — curved-wave theory x k3) bFourier transforms ( experimental, — theoretical) 45 The spectroscopic data presented in this section conclusively shows that the reaction of [Re 2 (CO)10] with Xe(OSeF 5 ) 2 produces [Re(CO)5 (OSeF5)]. Mass spectrometry showed the presence of the parent ion [Re(CO) 5 (OSeF5)]+ and infrared and 19F NMR spectroscopies indicated that the interaction between the seflate group and the rhenium centre is highly ionic in nature. This is to be expected when one considers the high electronegativity of the seflate group and the low oxidation state of the rhenium carbonyl moiety [Re(CO)5]+. Hence, it has been demonstrated, for the first time, that the seflate ligand is compatible with low valent metal carbonyl complexes. The Se edge EXAFS data collection and treatment was satisfactory, and a comparison of the Re-Se distance (3.55 A) to that of the Xe-Se distance (3.07 A) in Xe(OSeF5)2, highlights the covalent nature present in the Xe-seflate bond. The inability to collect and interpret the Re edge EXAFS data is unfortunate because it prohibits the calculation of the bridging Se-O-Re bond angle. 2.11. The Reaction Between [Mn2(CO)io] and Xe(OSeF 5 )2 - A survey of the literature indicates that F 2 and XeF 2 do not react with [Mn 2 (C O )10] to produce any stable carbonyl fluorides. The reaction between AgF and [Mn(CO) 5 Br] [{Mn(CO) 4 F}2] or, with was originally excess reported of AgF, to yield the dimer [Mn(CO)3 F3].t65^ However, reinvestigation by Horn et. alS66^ revealed that the overall reaction produced four structurally related clusters, [Mn4 (CO) 1 2 Fx(OH)4_x] (x = 0-4). The hydroxyl contamination resulted from moisture within the system which is extremely difficult to remove (N.B. AgF is hydrated). This reaction demonstrates the intrinsic differences between fluorine and the heavier halides, in which [Mn(CO) 5 X] (X=C1, Br and I) are all stable solids. Recent work carried out at L eicester^ has shown that [Mn 2 (CO)10] will react with xenon bis(teflate) to produce manganese pentacarbonyl teflate, 46 [Mn(CO)5 (OTeF5)]; which is, however, unstable in the presence of excess of Xe(OTeF5)2. A second paramagnetic species is also formed, and the infra-red data point towards the product being a d 5 Mn (II) species, believed to be cis[Mn(CO) 4 (OTeF5)2]. Any further excess of Xe(OTeF 5 ) 2 leads to decomposition and no identifiable products. Manganese pentacarbonyl teflate, [Mn(CO) 5 (OTeF5)], can also be produced by a methyl exchange reaction using teflic acid^9,10^ (Eqn. 2.25). [MeMn(CO)5] + HOTeF 5 -> [Mn(CO)5 (OTeF5)] + CH 4 Eqn. 2.25. The 19F NM R and infra-red spectral data for the products formed using the two different routes match exactly. These data, along with a crystal structure, offers conclusive evidence for the formation of [Mn(CO)5 (OTeF5)] from the reaction between Xe(OTeF 5 ) 2 and [Mn 2 (CO)10]. In view of these results the synthesis of [Mn(CO)5 (OSeF5)] was attempted via the reaction between [Mn 2 (CO)10] and Xe(OSeF5)2. Manganese carbonyl dimer, [Mn 2 (CO)10], was dissolved in dichloromethane to give a yellow solution. The solution was decanted on to an equimolar quantity of Xe(OSeF5)2, at -78°C. No immediate reaction occurred, but as the solution reached room temperature a vigorous reaction commenced which necessitated cooling with an acetone / C 0 2 bath. This cycle of warming and quenching was repeated until the reaction appeared to be complete. Analysis of the gas produced by infrared spectroscopy showed no carbon monoxide to be present, and it was inferred to be xenon. The solution changed to orange over the course of the reaction {ca. five minutes), and when the reaction was complete all volatile materials were removed to yield an orange solid. 47 The 19F NMR spectrum of this solid was recorded in a 4 mm FEP tube using D20 as the external lock substance and dichloromethane as the solvent. The spectrum showed an AX 4 pattern: 8 Fa 101.7 ppm, 8 Fe 69.2 ppm, 2 /( F a-Fe) 227 Hz , VCFg-Se) 1265 Hz and Aab = -32.5 (Figure 2.12). No 77Se satellites were resolved for Fa. The 13C NMR was recorded and this showed a single broad resonance at 5204.6 ppm, Avi/z 8 6 Hz c f [Re(CO)5 (OSeF5)] AVi/ 2 7 Hz (Figure 2.13). The 19F and 13C NMR spectra were poorly resolved compared with those for the [Re(CO)5 (OSeF5)] experiment. There may be two possible explanations:- 1) There may be stereochemical fluxionality within this system. However, running the NMR experiments at low temperature gave little improvement in the spectral resolution. 2) The manganese 55 nucleus has I = 5/2 and is 100% abundant. When I is greater than a Vi the nucleus possesses an electric quadrupolar moment, Q, which is due to a non-spherical charge distribution J 67^ This can interact with electric field gradients arising from electric charge distributions within the molecule. This interaction provides a means by which the nucleus can relax rapidly, and consequently can dramatically affect NMR spectra. In the 19F NMR spectrum the axial fluorine, Fa, showed a significant shift towards high frequency, and the Aab value of -32.5, demonstrates that the bonding between manganese and oxygen possesses a large degree of ionic character (c f Table 2.4). The infra-red spectrum was recorded as a Nujol mull of the solid and showed the following absorptions:- 2164 s, 2064 s, 2029 s, 864 s, 683 s, 624 s, 593 w and 543 s. 48 Figure 2.11. Fluorine-19 NMR spectrum for the product of the reaction between [Mn2(CO)10] and Xe(OSeF5)2. 100 (ppm) Figure 2.12. Carbon-13 NMR spectrum for the products of the reaction between [Mn 2 (CO)10] and Xe(OSeF5)2. 208 202 206 200 198 The three carbonyl absorptions expected for [Mn(CO)5 (OSeF5)] were clearly visible at 2164, 2064 and 2029 cm-1. The highly ionic character of the complex was also reflected in the selenium-oxygen stretching frequency of 864 cm-i, cf. 856 cm ' 1 for [Re(CO)5 (OSeF5)]. Electron-impact mass spectrometry met with limited success and the only identifiable fragments were due to [Mn(CO)2 (OSeF5)]+ m/z 302, [Mn(CO)(OSeF5)]+ m/z 274, [Mn(OSeF5)]+ m/z 246 and [Mn(CO)5]+ m/z 195 (for 55Mn and 80 Se). Attempts were made to record manganese K edge EXAFS data for [Mn 2 (CO)10] and [Mn(CO)5 (OSeF5)], the former being used as the model compound. Although data sets were successfully recorded for both compounds, the spectra were unusually noisy and no useful information was obtainable from them. 50 Transmission selenium K edge EXAFS data were collected for [Mn(CO)5 (OSeF5)], out to k = 15 later truncated to 13.5 A due A (k = photoelectron wave vector). This was to increased noise at higher k values. Three data sets were averaged and multiplied by k 3 to compensate for a decrease in intensity at higher k. The AFAC and VPI values were taken from the model compound [Xe(OSeF5)2] . No Fourier filtering was applied and the fit described below was compared to the average raw (background subtracted) EXAFS data. As with the model compound, the data was modelled using EXCURV92^59^ to two shells, (Table 6 2.11 fluorine atoms at 1.70(1) and Figure 2.11). A and Each 1 manganese atom at 3.38(1) shell was tested for A statistical significance Table 2.11. EXAFS data for [Mn(CO)5 (OSeF5)]. Parameter d(Sc-X) / 2 c 2/ Af a 2/ 1.70(1) Ab </(Se-Mn) / 2 EXAFSe 0.008(2) A 3.38(1) Ab 0.008(2) Fit index 0 2 .0 Rd 17.0 a Standard deviations in parentheses. b Debye-Waller factor. c Fit index = XjtQc T-XE)ki3]2. d R = [s(%T-%E)k 3 dk/s%Ek 3 dk] x 100 %. e E0 12.7 (4), AFAC 0.86 and VPI -4 .7 1 .f Mean bond length (X = F and O). From this evidence it has been shown that xenon bis(seflate) does indeed react with [Mn 2 (CO)10] to produce [Mn(CO)5 (OSeF5)]. Unlike the teflate analogue, [Mn(CO)5 (OSeF5)] is stable in the presence of excess of xenon bis(seflate), as determined using 19F NMR. 51 Figure 2.13. (a) Background-subtracted EXAFS and (b) the Fourier transform spectra for [Mn(CO)5 (OSeF5)]. (a) * CO -5 -15 (b) 3.0 CO c D >. i_ C6 i— 2 4 6 8 r/ A 1 aEXAFS ( experimental x k3, — curved-wave theory x k3) bFourier transforms ( experimental, — theoretical) 52 10 2.12. The Reaction Between [Ru(CO)3 (PPh3 )2] and Xe(OSeF 5 )2. As a result of the successes with similar systems, numerous reactions have been undertaken between XeF 2 and a range of ruthenium (0 ) complexes.[68,69^ O f these reactions, that between [Ru(CO)3 (PPh3)2] and XeF 2 has been the most fully investigated. The stepwise oxidative addition of xenon difluoride to [Ru(CO)3 (PPh3)2] occurs readily at low temperature. The mechanism ^6 9 , 7 0 1 involves oxidation by [XeF]+, then nucleophilic attack of a fluoride anion at the co-ordinated CO, followed ultimately by elimination of carbon monoxide to yield the stable octahedral complex, [OC-6- 13][RuF2(CO)2(PPh3)2] (Eqn. 2.26). XeF 2 + [Ru(CO)3 (PPh3)2] -> [RuF 2 (CO)2 (PPh3)2] + Xe + CO The products and intermediates formed in this Eqn. 2.26. reaction were characterised using multinuclear NMR and infrared spectroscopy, and X-ray crystallography. Although the reaction between [Ru(CO)3 (PPh3)2] and Xe(OTeF 5 ) 2 has also been r e p o r t e d t h e evidence for the formation of the analogous [Ru(CO)2 (PPh 3 ) 2 (OTeF5)2] was not conclusive. To establish whether the seflate group and low-valent metal-phosphine carbonyl compounds are compatible, the reaction between [Ru(CO) 3 (PPh3)2] and Xe(OSeF 5 ) 2 was investigated. Pale yellow [Ru(CO)3 (PPh3)2] was dissolved in dichloromethane and decanted, at -78°C, on to an equimolar quantity of Xe(OSeF5)2. On warming to room temperature a reaction commenced, as evidenced by the evolution of a gas, and continued at a steady rate for 10 minutes. Analysis of the gas by infrared spectroscopy showed the presence of carbon monoxide [v(CO) at 2143 c m '1]. All the volatile materials were removed and a brown solid was isolated. 53 The reaction was also repeated at lower temperatures in an analogous manner to that described above. The reactions were performed at 0, -5, -15 and -20 °C. At -20 °C, the rate of reaction decreased dramatically, however, it was evident that the temperature of the reaction does not alter the products produced. Separation of the products was attempted using various solvents, but this met with no success. The 19F NMR spectrum (Figure 2.14) was recorded in a similar manner to that described for the rhenium and manganese experiments. The spectrum showed an AX 4 pattern, 5Fa 105.1 ppm, 8 Fe 76.6 ppm, 2 7(FaFe) 235 Hz, lJ(FaSe) 1265 Hz and Aab = -28.5. No 77Se satellites were observed for Fe. Also observed were three singlets at 5-284.3, 8-313.8 and 5-340.2 ppm (Figure 2.15). The resonances at 8-284.3 and 5-313.8 ppm were broad and had half widths of 150 and 100 Hz respectively. The 19F NMR spectrum showed the presence of one seflate environment. Also present between 8 6 8 and 575 ppm were unresolved multiplet resonances, which possibly originated from selenium-fluorine containing decomposition products. The 31P NMR spectrum did not contain any resonances due to that of the triphenylphosphine or the starting material: [Ru(CO)3 (PPh3)2] (8 52.8 ppm). The following resonances were observed, a triplet at 526.2 ppm (7 18 Hz), doublets at 824.6 ppm (718 Hz) and 520.7 ppm (7 17 Hz) and a singlet at 820.4 ppm (Figure 2.16). By comparison of this data with the 31P NMR data for [RuF 2 (CO)2 (PPh3)2],[69] the triplet at 826.2 ppm is in accord with the value already recorded for this compound, the coupling constant being confirmatory of a cis 2 7(P-F) interaction J 69^ The reported chemical shift of the metal-bound fluorine atom in the 19F NMR data for [RuF 2 (CO)2 (PPh3)2] is 5-318 ppm, this is comparable with the resonance observed at 5-313.8 ppm (Figure 2.15). Therefore, it is proposed that a ruthenium seflate complex has been generated, but it has subsequently undergone decomposition as described in Section 2.3. 54 Figure 2.14. Fluorine-19 NMR spectrum for the products of the reaction between [Ru(CO)3 (PPh3)2] and Xe(OSeF5)2. 105 100 (ppm ) Figure 2.15. Fluorine-19 NMR spectrum for the products of the reaction between [Ru(CO)3(PPh3)2] and Xe(OSeF5)2. (ppm ) Figure 2.16. Phosphorus-31 NMR spectrum for the products of the reaction between [Ru(CO)3(PPh3)2] and Xe(OSeF5)2. 27 26 25 24 23 (ppm) 56 22 21 The 13C NMR spectrum showed only phenyl ring carbon resonances between 8127 and 8136 ppm; no carbonyl resonances were detected even after digital filtering and 24,000 scans. The infrared spectrum was recorded for a Nujol mull of the solid, however, this provided no definitive information about the nature of the reaction. The following absorptions were observed 2098 s, 2069, 2220 s, 1996 sh, 1096 s, 8 6 8 s, 752 w, 723 w, 712 w, 698 s, 667 w, 597 w and 566 w. FAB mass spectrometry met with limited success and a fragment possessing the correct isotopic [Ru(CO)(PPh 3 )2 F]+ m/z 673 (for 102 distribution was detected for Ru), using FAB. 2.13. The Reaction Between Xenon Bis(Seflate) and Iodine. In 1862, Kammerer reported ^7 ^ that at 70°C an iodine fluoride was liberated from the reaction between silver fluoride and iodine. It was proven, some years later, that iodine pentafluoride had been formed^72! (Eqn. 2.27). 5 AgF + 3 I 2 -> 5 Agl + IF 5 Eqn. 2.27. Subsequently, Moissan^73,74^ and Prideaux^75^ showed that iodine pentafluoride could be readily prepared by the direct combination of the elements at room temperature. Whilst this is still the preferred method employed today, IF 5 is also accessible from I2, HI or I 2 0 5 using a variety of fluorinating agents. 57 The chemistry of IF 5 hasbeen quite extensively studied but, relevant to the present work, in 1978, Seppelt et al. reported 1 1 7 , 2 8 1 that IF 5 and F 2 POOSeF 5 undergo a metathetical reaction to produce I(OSeF 5 ) 5 (Eqn. 2.28). IF 5 + F 2 PO-OSeF 5 -> FjCI(OSeF5)5.JC + POF 3 Eqn. 2.28. (jc = 0-4) The products, as described earlier (Section 2.4), were used to compare the electronegativity of the seflate group with that of fluorine. The reaction was monitored by observing the equatorial and axial fluorines of IF 5 using 19F NMR spectroscopy. Iodine tris(seflate), I(OSeF5)3, can be synthesised by the following reaction (Eqn. 2.29). 3 Cl-OSeF 5 + IC13 -> I(OSeF 5 ) 3 + 3 Cl2 Eqn. 2.29. Both I(OSeF5) and I(OSeF 5 ) 3 are highly unstable species and have never been isolated, their existence in solution being proven only by 19F NMR spectroscopy. This mirrors the instability of the lower fluorides of iodine. Iodine trifluoride, IF3, is thermally unstable 1 7 6 1 and disproportionates above 35°C (Eqn. 2.30). Iodine monofluoride, IF, is similarly unstable and disproportionates to iodine and iodine pentafluoride below room temperature (Eqn. 2.31). 2 IF 3 5 IF -> -> IF + IF 5 Eqn. 2.30. 2 I2 + IF 5 Eqn. 2.31. The first reaction was intended to synthesis I(OSeF 5 ) 5 and was conducted using a 1:5 molar ratio of iodine to xenon bis(seflate). Iodine was dissolved in dichloromethane, and then decanted on to the xenon bis(seflate) in 58 a 4 mm FEP tube at 25°C. An immediate reaction commenced and continued, at a steady rate, for 15 minutes. After this time no further evolution of a gas was observed. The FEP tube was heat sealed for analysis by NMR spectroscopy. A further sample was prepared in a similar manner, the solvent was removed, and the yellow-orange liquid that remained was submitted for mass spectral analysis. The 19F NMR spectral data for the above reaction is summarised in Table 2.12, and the spectra are shown in Figures 2.17-2.19. Only two AX 4 patterns were observed in the seflate region. A third AX 4 pattern, and a singlet with 77Se satellites were observed at lower frequency. Table 2.12. Fluorine-19 NMR spectral data for the products of the reaction between iodine and five molar equivalents of xenon bis(seflate). 2J (FaFe) ‘/( S e p ,) ppm Hz Hz 9.5 (d)b 90 42.8 (s) - 862 56.3 (q)b 91 - 59.7 (d) 224 1384 (d) 231 1367 8 6 6 .8 - 73.6 (q)b 2 2 1 - 78.3 (q)a - - a The quintet at 78.3 ppm was broad and coupling could not be resolved. b lJ(SeFe) could not be accurately measured due to the large number of peaks present. 59 Figure 2.17. Fluorine-19 NMR spectrum for the products of the reaction between I2 and five molar equivalents of Xe(OSeF5)2. 80 75 70 65 (ppm ) 60 Figure 2.18. Fluorine-19 NMR spectrum for the products of the reaction between I 2 and five molar equivalents of Xe(OSeF5)2. (ppm) Figure 2.19. Fluorine-19 NMR spectrum for the products of the reaction between I 2 and five molar equivalents of Xe(OSeF5)2. (ppm) 61 Although the 19F NMR spectrum is complicated by the large number of peaks, two AX 4 patterns are visible in the seflate region. If I(OSeF 5 ) 5 had been formed, then two seflate environments would be anticipated due to four equatorial and one axial seflate groups. No coupling between the fluorine atoms of the equatorial and axial seflate groups would be expected, as this would involve coupling through six bonds. The two sets of AX 4 patterns should have an integration ratio of 4:1, and this is precisely what is observed for the doublets at 566.8 and 559.7 ppm. Therefore, it is assumed that these two signals are due to the equatorial fluorines of I(OSeF5)5. Furthermore, on the basis of integration it can be seen that the doublet at 5Fe 59.7 ppm is associated with the quintet at 5Fa 73.6 ppm, and the doublet at 5Fe 6 6 . 8 ppm is associated with the quintet at 5Fa 78.3 ppm. This leads to values of Aab (axial seflate) = -13, and AAb (equatorial seflate) = -11.5; on comparison to the Aab values presented in Table 2.4 it can be seen that the interaction possesses some covalent nature, as would be expected an iodine (V). No 19F NMR chemical shifts have been previously reported for I(OSeF5)5, so no comparison is possible. However, the 19F NMR data provides conclusive evidence for the formation of I(OSeF5)5, and therefore extends the use of xenon bis(seflate) into the area of high valent non-metal seflate species. Electron impact mass spectrometry did not show the presence of any seflate containing species. Presumably decomposition had occurred, as previously described in Section 2.3. However, the spectrum did contain peaks which were assigned to [IF5]+ m/z 222, [IF4]+ m/z 203, [IF3]+ m/z 184, [IF2]+ m/z 165, [IF]+ m/z 146 and [I2]+ m/z 254. The presence of the IF 5 moiety in the products of reaction were also confirmed by the 19F NMR spectrum (Figures 2.18 and 2.19) which showed the presence of an AX 4 pattern, 5Fe 9.5 ppm and a quintet at 5Fa 57.0 ppm. This is in excellent agreement with the literature data for IF5.[i8] 62 A subsequent reaction was carried out using iodine and three molar equivalents of xenon bis(seflate). Iodine was dissolved in dichloromethane and then decanted on to the xenon bis(seflate) in a 4 mm FEP tube at 25°C. An immediate reaction commenced. After 10 minutes gas evolution had ceased and the reaction was presumed to be complete. The FEP tube was sealed and a 19F NMR spectrum was recorded. Unreacted iodine was present in the solution, as evidenced by its purple colour. Three sets of AX 4 patterns with accompanying 77Se satellites were observed, while, at lower frequency, a singlet with satellites and a doublet were observed (Table 2.13 and Figure 2.20). Two of the AX 4 patterns showed the same basic features as those for the I 2 and five molar equivalents of Xe(OSeF5)2, namely I(OSeF 5 ) 5 (Figure 2.17). However, the third AX 4 pattern, was indicative of another seflate environment. The signals at 5Fa 83.3 ppm and 5Fe 70.0 ppm are not due to the presence of any I(OSeF 5 )J 23,26^ which at 5Fe 92.0 ppm would be clearly visible. The 19F NMR data reported^23,26] for I(OSeF 5 ) 3 was obtained from a neat sample. It is not unreasonable to assume that solvent effects will alter the observed chemical shifts, as demonstrated for Ti(OTeF 5 ) 4 (Table 2.5). It is concluded here that the signals at 8 Fa 83.3 ppm and 5Fe 70 ppm, Aab = -13.8, arise from the presence of I(OSeF5)3. This is supported by the observation that on addition of more Xe(OSeF 5 ) 2 this signal disappeared and only resonances assignable to I(OSeF 5 ) 5 were observed. The presence of iodine in the reaction mixture, coupled with the fact that no Xe(OSeF 5 ) 2 was observed, would also suggest that a mixture of products was present. 63 Figure 2.20. Fluorine-19 NMR spectrum for the products of the reaction between I 2 and three molar equivalents of Xe(OSeF5)2. (ppm) Table 2.13. Fluorine-19 NMR spectral data for the products of the reaction between Iodine and three molar equivalents of Xenon bis(seflate). 6 ppm 2J (FaFe) lJ( SeFe) Hz Hz 9.5 (d) 90 - 42.2 (s) - 8 6 8 59.6 (d) 226 66.7 (d)b 2 1 0 70.0 (d) 229 73.5 (q)b 2 2 0 - 77.7 (q)a - - 83.3 (q) 227 - 1383 - 1343 1 L1 ~~ a The quintet was broad and no coupling could be resolved. D J (SeFe) could not be accurately measured. In both reactions singlet resonances were observed at either -642.8 ppm, /(Se-F) = 862 Hz or 642.2 ppm, / ( Se-F) = 867 Hz, and these were presumably due to the same species. An inspection of the literature indicates that it is not a selenium (VI) oxide fluoride or fluoride, cf. SeOF4,[3^ Se 2 0 2 F8,[77] S e 0 2 F2[78^ and SeF 6 [ 7 9 1 all of which possess Se-F coupling constants in the region of 1500-1300 Hz. The 19F NMR spectrum of neat SeOF 2 p r o d u c e s ^ a single resonance at 633.5 ppm, /(Se-F) = 837 Hz. Considering how solvent effects can considerably alter the observed chemical shifts {cf. Table 2.5), and the similarity in the magnitude of the coupling constants, it appears the reaction between iodine and xenon bis(seflate) generates SeOF 2 as a decomposition product. 65 2.14. Discussion. Table 2.14. A comparison of the v(CO), v(Se-O) and v(Te-O) values for various carbonyl derivatives. Complex v(CO) / cm ' 1 [Re 2 (CO)10] 2070, 2013, 1975 [Re(CO)5 (OSeF5)] 2168, 2045, 1986 856 b [Re(CO) 5 (OTeF5)] 2164, 2055, 1998 843 8 [Re(CO) 5 Cl] 2155, 2046, 1983 - 63 [Re(CO)5 Br] 2151,2043, 1985 - 63 [Re(CO)5 I] 2144, 2041, 1989 - 63 [Mn 2 (CO)10] 2046, 2013, 1983 - [Mn(CO) 5 (OSeF5)] 2164, 2064, 2029 864 b [Mn(CO) 5 (OTeF5)] 2155,2070, 2016 848 8 [Mn(CO) 5 Cl] 2139, 2055, 1999 - 63 [Mn(CO)5 Br] 2134, 2050, 2001 - 63 [Mn(CO)5 I] 2125, 2043, 2003 - 63 v (X -0 )a /c m _ 1 - Reference b b a X = Selenium or tellurium. b This work. Table 2.14 lists the v(CO), v(O-Se) and v(O-Te) frequencies for a range of halide, seflate and teflate carbonyl complexes, for which a number of trends are apparent:- 66 i) On going from Re (0) to Re (I) a shift towards higher frequency is observed for v(CO). This is the direct result of the increase in the oxidation state which leads to a decrease in the electron density available for % back bonding. ii) The v(CO) data for Re(CO)5X and Mn(CO)5X (X = seflate or teflate) are similar, suggesting that the ionicity of the seflate derivatives is virtually the same as that for the teflate derivatives. iii) The carbonyl stretching frequencies for the seflate and teflate compounds are higher than those for the respective halide analogues. This reflects the high electronegativities of the seflate and teflate ligands, which, as highlighted in Section 2.4, are very similar to that of fluorine. iv) Due to the mass difference between Se and Te, v(Se-O) is greater than v(Te-O). It has been demonstrated that xenon bis(seflate) will oxidise zero valent transition metal carbonyl compounds to produce low valent transition metal carbonyl seflate species. Table 2.15 highlights the similarities in nature of seflate and teflate ligands. Although the reaction between xenon bis(teflate) and [Ru(CO) 3 (PPh 3 )2 ] may produce [Ru(CO)2 (PPh 3 )2 (OTeF5)2], the spectroscopic evidence is not conclusive. The reaction between xenon bis(seflate) and [Ru(CO)3 (PPh 3 )2] produced a seflate containing species, however, several species were generated and identification of the products was difficult. The only product identified was [RuF 2 (CO)2 (PPh 3 )2 ], and this was presumably a decomposition product. Although we have shown that the seflate ligand is compatible with low valent metal carbonyl species, further work is needed to determine whether the seflate group and phosphine donor ligands are compatible. 67 Table 2.15. The comparative chemistry of XeL2, L = fluoride, seflate or teflate. Reactant Reagent Products Reference [Mn 2 (CO)io] XeF 2 No fluoro products - Xe(OTeF 5 ) 2 [Mn(CO)5 (OTeF5)] 8 4 [Mn(CO)4 (OTeF5)2] 4, Decomposition [Re 2 (CO)10] Xe(OSeF 5 ) 2 [Mn(CO)5 (OSeF5)] XeF 2 [Re(CO)5 .JuF.ReF5] c 64 [Re(CO)6 ][ReF6] [Ru(CO) 3 (PPh3)] h Xe(OTeF 5 ) 2 [Re(CO)5 (OTeF5)] Xe(OSeF 5 ) 2 [Re(CO)5 (OSeF5)] XeF 2 [OC-6-13] [RuF 2 (CO) 2 (PPh3)2] 69 Xe(OTeF 5 ) 2 [Ru(CO)2 (PPh 3 )2 (OTeF5)2] a 8 Xe(OSeF 5 ) 2 [RuF 2 (CO)2 (PPh3)2]b XeF 2 if5 Xe(OTeF 5 ) 2 I(OTeF 5 ) 5 Xe(OSeF 5 ) 2 I(OSeF 5 ) 5 a Postulated.b The only one of several products which could be identified. c This work. 68 8 c c - 17,28 c The teflate group and the fluoride ion are highly electronegative ligands, and have been closely compared. Indeed, in the area of high valent transition metal and main group chemistry, the teflate and fluoride ligands are virtually interchangeable as demonstrated by Table 2.15. However, the Table highlights one major difference. The isolation of [Mn(CO)5 (OSeF5)] and [Mn(CO)5 (OTeF5)] clearly distinguishes the seflate and teflate ligands from the fluoride ion. This difference may be a consequence of the higher electronegativity of the fluoride ligand which results in an unstable manganesecarbonyl environment. However, further work is needed to expand the chemistry of the seflate ligand in order to determine fully the similarities between the fluoride and seflate/teflate ligands, and the factors which underlie any differences. 69 References Chapter Two [1] K. Seppelt, Angew. Chem., Int. Ed. Engl., 1972, 15, 630. [2] A. Engelbrecht and F. Sladky, Angew. Chem., Int. Ed. Engl., 1964, 3, 383. [3] K. Seppelt, Angew. Chem., Int. Ed. Engl., 1974, 13, 91. [4] K. Seppelt, Angew. Chem., Int. Ed. Engl., 1976, 15, 44. [5] W. Porcham and A. Engelbrecht, Montash. Chem., 1971, 102, 333; W. Porcham and A. Engelbrecht, Z. Phys. Chem. (Leipzig), 1971, 248, 111 . [6 ] A. Engelbrecht and F. Sladky, Montash. Chem., 1965, 96, 159. [7] A. Engelbrecht and F. Sladky, Inorg. Nucl. Chem. Lett., 1965,1,15. [8 ] M. C. Crossman, Ph.D. Thesis, University of Leicester, 1995. [9] S. H. Strauss, K. D. Abney, K. M. Long and O. P. Anderson, Inorg. Chem., 1984, 23,1994. [10] K. D. Abney, K. M. Long, O. P. Anderson and S. H. Strauss, Inorg. Chem., 1987, 26, 2638. [11] F. Sladky, H. Kropshofer, O. Leitzke and P. Peringer, J. Inorg. Nucl. Chem. Supplement, H. H. Hyman Memorial Volume 1976, Pergamon Press, ed. J. J. Katz and I. Sheft. [12] H. P. A. Mercier, J. C. P. Sanders and G. J. Schrobilgen, J. Am. Chem. Soc., 1 9 9 4 , 116, 2921. [13] H. Kropshofer, O. Leitzke and F. Sladky, J. Chem. Soc., Chem. Commum., 1973, 134; J. F. Sawger, G. J. Schrobilgen, Acta Crystallogr., Sect. B, 1982, 38, 1561. [14] H. Kropshofer, O. Leitzke, P. Peringer and F. Sladky, Chem. Ber., 1981, 114, 2644. [15] H. Kropshofer and F. Sladky, Inorg. Nucl. Chem. Lett., 1972, 8 , 195. [16] G. W. Fraser and G. D. Meikle, J. Chem. Soc., Dalton Trans., 1975, 1033. 70 [17] G. W. Fraser and J. B. Millar, J. Chem. Soc., Dalton Trans.,1974, 2029. [18] D. Lentz and K. Seppelt, Angew. Chem., Int. Ed. E ngl, 1978,17, 355 and references cited therein. [19] E. Jacob, D. Lentz, K. Seppelt and A. Simon, Z. Anorg. Allg. Chem., 1981,472,7. [20] L. K. Templeton, D. H. Templeton, N. Bartlett and K. Seppelt, Inorg. Chem., 1976, 15, 2718; Inorg. Chem., 1976, 15, 2720 [21] K. Seppelt, Chem. Ber., 1972, 105, 2431. [22] K. Seppelt, Angew. Chem., Int. Ed. Engl., 1974, 92, 13. [23] K. Seppelt, Chem. Ber., 1973,106,1920. [24] K. Seppelt, Chem. Ber., 1977,110,1470. [25] A. Clouston, R. D. Peacock and G. W. Fraser, J. Chem. Soc., Chem. Commum., 1970, 1197. [26] K. Seppelt, Chem. Ber., 1978, 106, 157. [27] K. Seppelt and D. Nothe, Inorg. Chem., 1973, 12, 2727. [28] D. Lentz and K. Seppelt, Z. Anorg. Allg. Chem., 1980, 460, 5. [29] P. Huppmann, D. Lentz and K. Seppelt, Z. Anorg. Allg. Chem., 1981, 472, 26. [30] J. E. Smith and G. H. Cady, Inorg. Chem., 1970, 9, 1442. [31] K. Seppelt, Chem. Ber., 1972,105, 3131 and references cited therein. [32] F. Sladky and H. Kropshofer, J. Chem. Soc., Chem. Commum., 1973, 600. [33] K. Seppelt, Z. Anorg. Allg. Chem., 1974, 406, 287. [34] R. Damerius, P. Huppman, D. Lentz and K. Seppelt, J. Chem. Soc., Dalton Trans., 1984, 2821. [35] H. Oberhammer and K. Seppelt, Inorg. Chem., 1978, 17, 1435. [36] H. Oberhammer and K. Seppelt, Angew. Chem., Int. Ed. Engl., 1978, 17, 69 and references cited therein. [37] K.Seppelt, Angew. Chem., Int. Ed. Engl., 1972, 11, 723. [38] K. Seppelt and H. H. Rupp, Z. Anorg. Allg. Chem., 1974, 409, 338. 71 [39] K. Seppelt, Chem. Ber., 1975, 108, 1823. [40] P. Huppmann, J. Labischinski, D. Lentz, H. Pritzkow and K. Seppelt, Z Anorg. Allg. Chem., 1982, 487, 7. [41] T. Birchall, R. D. Myers, H. Waard and G. J. Schrobilgen, Inorg. Chem., 1982, 21, 1068. [42] G. J. Schrobilgen, J. H. Holloway, P. Granger and C. Brevard, Inorg. Chem., 19 7 8 , 17, 980; G. J. Schrobilgen, NMR and the Periodic Table, eds. R. Harris and B. Mann, Academic Press, London, 1978, ch. 14. [43] G. J. Schrobilgen, R. Bums and P. Granger, J. Chem. Soc., Chem. Commun., 1978, 957. [44] A. Engelbrecht and P. Peterfly, Angew. Chem., Int. Ed. Engl., 1969, 8 , 768. [45] K. O. Christe and H. Oberhammer, Inorg. Chem., 1981, 20, 296. [46] K. Seppelt, Angew. Chem., Int. Ed. Engl., 1982, 21, 877. [47] J. A. Pople, W. G. Schneider and H. J. Bernstein, High Resolution Nuclear Magnetic Resonance, McGraw - Hill Book Co., New York, 1959. [48] K. Seppelt, Inorg. Synth., 1980, 20, 38. [49] R. K. Harris and K. J. Packer, J. Chem. Soc., 1961, 4736. [50] P. Bladen, D. H. Brown, K. D. Crosbie and D. W. A. Sharp, Spectrochim. Acta, Part A, 1970, 26, 2221. [51 ] Param eter Adjustment in NMR by Irerative Calculation; NMR Program Library, Quantum Chemistry Program Exchange. [52] K. Schroder and F. Sladky, Chem. Ber., 1980, 113, 1414. [53] E. Mayer and F. Sladky, Inorg. Chem., 1975, 14, 589. [54] P. K. Miller, K. D. Abney, A. K. Rappe, O. P. Anderson and S. H. Strauss, Inorg. Chem., 1988, 27, 2255. [55] S. H. Strauss, K. D. Adney and O. P. Anderson, Inorg. Chem., 1986, 25, 2806. [56] P. J. Kellet, Ph.D. Thesis, Colorado State University, 1989. 72 [57] K. D. Abney, K. M. Long, O. P. Anderson and S. H. Strauss, Inorg. Chem., 1987, 26, 2638. [58] K. Seppelt, D. Lentz and G. Kloter, Inorg. Synth., 1986, 24, 27. [59] EXCURV92, SERC Daresbury Laboratory Program, N. Binstead, J. W. Campbell and S. J. Gurman, 1992. [60] R. W. Taylor, K. J. Martin and P. Meehan, J. Phys. C, 1987, 20, 4005; N. Binstead, S. L. Cook, J. Evans, G. N. Greaves and R. J. Price, J. Amer. Chem. Soc., 1987, 109, 3669. [61] I. C. Bowater, R. D. Brown and F. R. Burden, J. Mol. Spectrosc., 1968, 28, 461. [62] D. M. Bruce, J. H. Holloway and D. R. Russell, J. Chem. Soc., Dalton Trans., 1978, 64, 1627. [63] H. D. Kaesz, R. Bau, D. Hendrickson and J. M. Smith, J. Am. Chem. Soc., 1967, 89, 2844. [64] J. H. Holloway, J. B. Senior and A. C. Szary, J. Chem. Soc., Dalton Trans., 1987, 741. [65] M. K. Chaudhuri, M. M. Kaschani and D. Winkler, J. Organomet. Chem., 19 7 6 , 113, 387. [6 6 ] E. Horn, M. Snow and P. Zeleny, Aust. J. Chem., 1980, 33, 1659. [67] R. K. Harris, Nuclear Magnetic Resonance Spectroscopy, Longman, 1st edn., 1986, 134-137. [6 8 ] A. J. Hewitt, J. H. Holloway, R. D. Peacock, J. B. Raynor and I. L. Wilson, J. Chem. Soc., Dalton Trans., 1976, 579 and references cited therein. [69] S. A. Brewer, K. S. Coleman, J. Fawcett, J. H. Holloway, E. G. Hope, D. R. Russell and P. G. Watson, J. Chem. Soc., Dalton Trans., 1995, 1073. [70] K. S. Coleman, Ph.D. Thesis, University of Leicester, 1996. [71] H. Krammer, J. Prakt. Chem., 1862, 85, 452. [72] G. Gore, Proc. R. Soc., 1870, 19, 235. 73 [73] H. Moissan, Ann. Chim. Phys., 1891, 24, 244. [74] H. Moissan, C. R. Seances Acad. Sci.y Paris 1902,135,563. [75] E. B. R. Prideaux, Trans. Chem. Soc., 1960, 89, 316. [76] V. Gutmann, Halogen Chemistry, Academic Press, London and New York, 1967, vol. 1, 133-206. [77] K. Seppelt, Angew. Chem., Int. Ed. Engl., 1974,13, 92. [78] K. Seppelt and D. D. Desmarteau, Inorg. Synth., 1980, 20, 36. [79] E. L. Muetterties and W. D. Philips, J. Am. Chem. Soc., 1959, 81, 1084. [80] K. Seppelt, D. Lentz and G. Kloter, Inorg. Synth., 1986, 24, 25. 74 CHAPTER THREE Bromine Oxide Fluoride Chemistry 3.1. Introduction. The halogen oxide fluorides and their complexes are of fundamental importance to inorganic chemistry as examples of unusual, discrete, molecular geometries. When compared with their chlorine and iodine counterparts, the bromine oxide fluorides have been little investigated, and those compounds which are known have been poorly characterised. Our goal was to develop new synthetic routes to novel bromine oxide fluoride species, and to use low-temperature infrared and EXAFS spectroscopies as primary characterisation techniques. To put the work into context an overview of the halogen oxide fluoride chemistry is presented below. However, hypofluorite compounds will not be included. 3.2. Structures of the Oxide Fluorides. The shapes of the halogen oxide fluorides can be rationalised using valence shell electron pair repulsion theoryf1,2^ (VSEPR). This states that the geometry of a molecule AXmEn is determined by the repulsion between the pairs of bonding electrons linking A to the ligands, X, and the non-bonding electron pairs, E, in the valence shell of the central atom A. Multiple bonds between A and X are considered to behave like a single pair of electrons, and thus do not change the overall arrangement of the ligands (although they will affect the angles between them). The sum of the number of bonding sets of electrons (m) and the non-bonding pairs of electrons (n) is the criterion which determines the structure of the molecule. Hence, the geometries are (m+n) = 2 (linear), 3 (triangular), 4 (tetrahedral), 5 (trigonal bipyramidal) and 6 (octahedral). Consideration of the numbers and types of repulsions between the various electron pairs allows one to predict, which positions will be occupied by ligands, and which will be taken up by non-bonding electron pairs. 75 According to VSEPR theory, bonds of order greater than one, in this case the X = 0 bonds, and unshared electron pairs have a strong preference for the equatorial positions of a trigonal bipyramid and for trans positions in an octahedron. Table 3.1, shows the predicted structures for the bromine (V) and (VII) oxide fluorides, including the related anions and cations formed in the reaction with an appropriate Lewis acid or base. Table 3.1. Structures of the known and possible oxide fluoride compounds of bromine (V) and bromine (VII). Molecule _____ Number of Number of bonds lone pairs (m) (n) _ _ (m+n) Arrangement Approximate of bonds and Symmetry shape lone pairs _ Pseudo “ trigonal a[Br03]+ 3 0 3 Trigonal jj o B r02F 3 1 4 /X D 3h o Pseudo tetrahedral Cs I Br o^llN O [BrOF2]+ 3 1 4 Pseudo .. Cs tetrahedral Br c/'IN F B r03F 4 0 4 Pseudo F tetrahedral O 76 F '3v a[Br02F2]+ 4 0 4 c2v F Pseudo tetrahedral Br o ^ l0 l \F a[Br03F2]- 5 0 5 F Pseudo trigonal bipyramidal D 3h V lBr = 0^1 0 F aBr02F3 5 0 5 F Pseudo trigonal bipyramidal C2v v j. B r ----- F o^| F a[BrOF4]+ 5 0 5 trigonal bipyramidal BrOF3 [Br02F2]- 4 4 1 1 5 5 c2v F Pseudo 's L o p / iF Pseudo F trigonal v j B r ----- 1 bipyramidal f/ Pseudo trigonal bipyramidal Cs |F F C2v J . B r ----- : I F [BrOF4]- 5 1 6 Pseudo octahedral l[Br02F4] 0 F\ II_F Rr -- Pseudo O octahedral Br: I C4v D4h -F 'F O aBrOF< Pseudo O octahedral Br: 'Aw -F ‘F F Species are unknown to date. 77 3.3. The Halogenyl Fluorides, XO 2 F. 3.3.1. Chloryl fluoride. Chloryl fluoride, C10 2 F, has been prepared by several methods. Early syntheses used shock-sensitive chlorine oxides such as C120 and C10 2 J 3,4^ The danger involved with the use of these materials can be avoided by the use of equimolar amounts of C1F3 and Na[C103] which also gives rise to C10 2 F.[5,6] Chloryl fluoride, a colourless gas at room temperature, is a powerful oxidising and fluorinating reagent which is sta b le ^ up to 300°C. Studies using microwave spectroscopy have provided information on the molecule's internal p aram eters'8,9^ and shown that it has the predicted Cs symmetry. The reaction of C102F with dried CsF at -80°C affords Cs[C10 2 F2],^10,11^ and analysis by infrared sp e c tro sc o p y ^ produces data consistent with the anion being a pseudo-trigonal bipyramid of C2v symmetry. Chloryl fluorides react with the Lewis acids BF3, PF5, AsF5, SbF 5 and VF5^ to form salts comprising of the cation [C102]+ and the anions [BF4]', [PF6]‘, [AsF6]', [SbF6]' and [VF6]' respectively. The reaction with PtF 6 gives a mixture of [C10 2 F2]+[ PtF6]‘ and [C10 2 ]+[PtF6] \ [13,14] 3.3.2. Bromyl fluoride. Bromyl fluoride, B r0 2 F, has been known since the early 1950's, although its synthesis has aroused much controversy. In 1957, Schmeisser reported that the reaction between BrF 5 and K [B r03] at -60°C yielded B r0 2F among its p r o d u c t s H o w e v e r , Bougon reported that K [B r0 2 F2], BrF 3 and 0 2 are produced, ^ whilst Gillespie suggested that the product mixture consisted of K [B r0 2 F2], K[BrOF4] and B r0 2 F j 17^ Evidently no clear and established route exists to bromyl fluoride. This may be a result of the highly 78 reactive character of B r0 2 F, or a reflection of the variety of products that may be produced in the reaction. Bromyl fluoride is a colourless solid at low temperatures with a melting point of -9 ° C .^ At room temperature, it slowly decomposes and the liquid produced is generally yellow owing to the presence of BrF3. The liquid decomposes violently above 56°C to BrF3, Br 2 and 0 2. The Raman spectra of bromyl fluoride, ^18-2°] as neat compound and as an HF solution, are consistent with a monomeric, pseudo-pyramidal molecule of Cs symmetry. It forms adducts with AsF 5 and SbF 5 J 21,22^ which are, however, of low stability and decompose near room temperature. The reaction of PtF 6 and B r0 2F at -120°C yields a mixture of [BrOF 2 ]+[PtF6]' and [B r0 2 ]+[PtF6] '.[23] The attempted oxidative fluorination of B r0 2F using KrF 2 does not yield bromine (VII) oxide fluorides, but instead, proceeds via BrOF3, to yield BrF 5 as the only product. Bromyl fluoride will form adducts with Lewis bases such as KF. However, a preferable route to K [B r0 2 F2] is the fluorine-oxygen exchange reaction between K[BrF6] and K [B r03] in CH 3 CN.[19] A mixture of K [B r0 2 F2] and K[BrOF4] is obtained, and separation of the products exploits the solubility of K[BrOF4] in CH 3 CN, compared with the insolubility of K[BrOF4]. Potassium difluorobromate, K [B r0 2 F2], is a white solid which is stable at room temperature. Using infrared and Raman spectroscopy, Bougon et al. have shown[24] that the structure is a pseudo-trigonal bipyramid of C2v symmetry. 3.3.3. Iodyl fluoride. Iodyl fluoride, I 0 2 F, was first prepared in 1953[251 by the thermal decomposition of IOF3, which also produces IF5. The reaction is reversible, and consequently, IOF 3 is formed by refluxing I 0 2F with IF5. Iodyl fluoride is also produced by the fluorination of I 2 0 5 in AHF, at 20°CJ3^ Iodyl fluoride is a stable colourless solid at room temperature, which slowly evolves HF in moist 79 air.t16^ Vibrational spectroscopic s tu d ie s ^ were hampered by the complexity of the spectra, which is indicative of extensive couplings between the different vibrational modes. This, together with low volatility at high temperatures suggests that I 0 2F is a polymeric species. Iodyl fluoride reacts with Lewis acids to form [I0 2 ]+[AsF6]" salts such as On the other hand, reaction of I 0 2F with Lewis bases affords complexes involving the anion [I0 2 F2]'. Thus, KF and I 0 2F react in AHF to yield K [I0 2 F2] and vibrational spectroscopic studies indicate that it is isostructural with [C10 2 F2]' and [B r0 2 F2]'. 3.4. Halogen Oxide Trifluorides, XOF3 . 3.4.1. Chlorine oxide trifluoride. Chlorine oxide trifluoride, C10F3, was first synthesised in 1 9 6 5 ^ by the fluorination of C12 0 , a route limited by the explosive nature of C12 0 . Later, Bougon et a l prepared CIOF 3 by the UV irradiation of a mixture o f C1F3 and OF 2 J 28] A large scale preparation^29^ involves the fluorination of chlorine nitrate, C 1 0 N 02, at -35°C, which affords a mixture of C10F 3 and F N 0 2. These are separated by virtue of a large difference in their vapour pressures. Chlorine oxide trifluoride is a colourless compound with a melting point of -43°C and a boiling point of 28°C. Gas phase electron diffraction showed the molecular structure to be a distorted pseudo-trigonal bipyramid,[30^ with a doubly bonded oxygen and a lone pair lying in the equatorial plane. Chlorine oxide trifluoride possesses a stability intermediate between that of C1F3 and CIF 5 , and reacts with glass, quartz and most metals causing both fluorination and oxygenationP 1^ Its reaction with organic substances, even at low temperatures can be explosive. As a powerful oxidant, it has proved to be a useful supporter of combustion for rocket fuels such as N 2 H4. 80 Chlorine oxide trifluoride forms stable 1:1 adducts with a variety of Lewis acids,t27,32J e.g. BiF5, SbF5, AsF5, TaF5, NbF5, VF5, PF 5 and BF3. The vibrational spectra observed for the [C10F2]+ salts[33] showed the presence of six fundamental vibrations, which is consistent with them being pseudotetrahedral molecules of Cs symmetry. Chlorine oxide trifluoride forms stable adducts with strong Lewis bases^32,34^ such as CsF, RbF and KF; but no reaction was observed with the weaker base NOF. Vibrational spectroscopic data was used to infer the molecular structure of [C10F 4 ] 'J 32^ however, it appears that splitting of the degenerate modes led to an inconclusive assignment. Further reactions include attempts to isolate a [C10F4]+ salt, utilising the reactions of CIOF 3 with SbF 5 -F 2 or PtF 6 J 14,34^ These failed but the latter reaction produced [C10F 2 ]+[PtF6]". 3.4.2. Bromine oxide trifluoride. Bromine oxide trifluoride was first prepared^35] by the reaction of K[BrOF4] and [ 0 2 ]+[AsF6]" in a solution of BrF5. It can also be made using the reaction of K[BrOF4] and the weak Lewis acid HF.[17^ The HF is removed at low temperature to leave K[HF2] and BrOF3. The BrOF 3 cannot simply be distilled from the BrOF 3 -K[HF2] mixture since the reaction is reversible. Instead, BrF5, into which the BrOF 3 dissolves, is distilled on to the mixture. The solution can then be decanted off to leave the solid behind. In 1987, Wilson and Christe developed a new, high yield, one-step synthesis of BrOF 3 which involves the reaction of L i[N 0 3] and an excess of BrF 5 J 36^ Bromine oxide trifluoride is a colourless liquid ^3 7 1 or solid, with a melting point range of -5 to 0°C. At room temperature it slowly decomposes to produce BrF 3 and 0 2. Vibrational spectroscopic data has provided conclusive evidence that its structure is pseudo-trigonal bipyramidal,[19,38] analogous to that of C10F3. 81 Bromine oxide trifluoride possesses a similar amphoteric nature to that of CIOF 3 . The reaction between BrOF 3 and the Lewis acids BF3, AsF 5 and SbF 5 affords 1 3 9 1 the adducts [BrOF 2 ]+[BF4]-, [BrOF 2 ]+[AsF6]’ and [BrOF 2 ]+[SbF6] ' respectively. The stability of the complex formed increases considerably with the increasing strength of the Lewis acid employed. The adduct [BrOF 2 ]+[SbF6]" can also be prepared from the reaction between I 0 2 F 3 *SbF5 and BrF5. The vibrational spectroscopic data from these complexes^23’37,39,40] are consistent with their containing a cation of pseudopyramidal geometry (Cs symmetry) and the assignments made are in good agreement with those for the isostructural species SeOF2, SOF 2 and [C10F2]+. Salts of the anion [BrOF4]" are easily synthesised using the method described by Wilson and ChristeJ36^ The reaction between BrF 5 and the alkali metal nitrates M [N 0 3], M = Na, K, Rb or Cs, yields the corresponding anionic salts. From Na+ —» Cs+, smaller excesses of BrF 5 and shorter reaction times are required. The salts are stable white solids at room temperature, and the vibrational data suggest that the anion possesses C4v symmetry. However, as with [C10F4]‘, the assignments were made difficult due to splittings of the degenerate modes. 3.4.3. Iodine oxide trifluoride. Iodine oxide trifluoride was first claimed to have been synthesised by Ruff and Braida in 1934. Later, in 1953, this claim was c o n firm e d ^ when IOF 3 was prepared by refluxing a saturated solution of I 2 0 5 in IF5. On cooling, colourless needles of IOF 3 are formed. Iodine oxide trifluoride is stable up to 110°C at which temperature it decom poses^ to IF 5 and I 0 2F and, as explained earlier, this is reversible (see Section 3.3.3). Iodine oxide trifluoride has been characterised using X-ray crystallography.^ The molecular structure is a pseudo-trigonal bipyramid with axial fluorines and a lone pair, a doubly bonded 82 oxygen and a fluorine lying in the equatorial plane. The compound is isostructural with CIOF 3 . The anion [IOF4]‘ is accessible from the indirect reaction of KF and I 2 0 5, in a 5:1 molar ratio, with a large excess of IF5/ 42! The mixture is refluxed for one hour and the white solid isolated is stable up to 200°C. Quenching with water produces HF and K [I0 3]. X-ray crystallography^42! has shown that [IOF 4 ]' is a square based pyramid with the four fluorine atoms in the equatorial plane; the vibrational spectroscopic data is in accord with the predictions that the molecule would have C4v symmetry. 3.5. Perhalogenyl Fluorides, XO 3 F. 3.5.1. Perchloryl fluoride. Perchloryl fluoride, C10 3 F, was first synthesised in the early 1950's and has been extensively investigated since. Perchloryl fluoride is readily prepared by several different ro u tes/19! including the fluorination of K[C103] using F2, in the super-acid medium H S 0 3 F-SbF5. The electrolysis of a saturated solution of NaC10 4 in AHF also yields CIO 3 F. Perchloryl fluoride is a stable colourless gas with a melting point of -47 °C. The physical properties are well docum ented/19! Its inertness relative to the other halogen oxide fluorides is a consequence of its energetically favourable pseudo-tetrahedral configuration. Perchloryl fluoride hydrolyses slowly in water and is thermally stable up to 400°C. As a consequence of the low polarity of CIO 3 F, it is soluble in a wide range of non-polar solvents/43! and at elevated temperatures it is a powerful oxidising agent. Gas-phase electron diffraction studies confirm that it has a pseudo-tetrahedral geometry of C3v sym m etry/44! Applications include its selective fluorinating properties in organic chem istry/45^ e.g. the replacement of the hydrogen atoms of a CH 2 group by fluorine. It is also possible to introduce chlorate groups, [CIO3 ], into organic 83 molecules, e.g. the reaction of C 6 H5Li and C103F produces C 6 H 5 C10 3 J46^ Perchloryl fluoride has also been extensively used, alone or mixed, with other halogen fluorides as an oxidant for rocket fuels . [ 2 7 1 The UV photolysis of CIO 3 F and CIF 3 , CIF 5 , OF 2 or F 2 produces C10F 3 J47^ Perchloryl fluoride behaves as a mild fluorinating agent and converts UF 4 to UF 6 via an unknown uranium oxide fluoride,f48^ but it is inert towards both Lewis acids and bases. 3.5.2. Perbromyl fluoride. Perbromyl fluoride, B r0 3 F, is prepared by the action of a powerful fluorinating agent such as SbF5, AsF5, [BrF 4 ]+[AsF6]' or BrF 5 on [K B r04] in anhydrous H FJ49,50^ Perbromyl fluoride is a colourless gas with a melting point of -110 °C. Electron diffraction studies confirm the structure is pseudotetrahedral and vibrational studies are in agreement with the molecule having C3v symmetry. The chemical behaviour of B r0 3F is similar to that of CIO 3 F. However, the difficulty involved in making B r0 3 F, and its lower stability, means its chemistry is less diverse. No adducts of B r0 3F with Lewis acids or bases have been reported to date. 3.5.3. Periodyl fluoride. Periodyl fluoride, I 0 3 F, can be prepared by passing fluorine through a solution of H I0 4 in HF, or by the reaction of K [I0 4] with H S 0 3 f J 48,52,53^ It is a white solid which is stable up to 100°C. Vibrational analysis has been attempted, however, polymerisation appears to occur and this has prevented a satisfactory assignment. Periodyl fluoride possesses some fluoride ion donor properties and a solution of the compound in HF reacts with AsF 5 or BF 3 to give compounds containing the cation [I0 3 ]+. 84 3.6. Halogen Dioxide Trifluoride, XO2 F3 . 3.6.1. Chlorine dioxide trifluoride. Chlorine dioxide trifluoride is prepared by a multi-step synthesis.1[54^ The oxidation of C102F using PtF 6 produces [C102 F 2 ]+[PtF6]" and [C102 ]+[PtF6] \ The reaction of these cations with F N 0 2 or FNO, at -78°C, produces C10 2 F 3 and C102F respectively. The C10 2 F 3 is then separated from the C102F by fractional condensation. Any remaining C102F can be removed by the addition of BF 3 to the mixture, and this produces the adducts [C10 2 F 2 ]+[BF4]' and [C10 2 ]+[BF4]". The species [C10 2 ]+[BF4]‘ is unstable above 20°C, and can be removed as it is the only volatile product. The final step involves the reaction of [C10 2 F 2 +][BF4]" with F N 0 2 which liberates the volatile C10 2 F3. Chlorine dioxide trifluoride is a stable g a s ^ with a melting point of -81°C and a boiling point of -22°C. Vibrational studies combined with 19F NMR spectroscopic data^54! are consistent with C10 2 F 3 having the structure of a pseudo-trigonal bipyramid of C2v symmetry; this corresponds to two fluorines occupying the axial positions, as would be predicted using VSEPR theory. Chlorine dioxide trifluoride is a strong oxidative fluorinator which reacts explosively with organic materials and fluorinates metal surfaces, producing c io 2 f. The synthesis of [C10 2 F2]+ has already been highlighted above, and salts with the corresponding counter ions [PtF6] \ [BF4]' and [AsF6]" are solids which are stable at 25°CJ55^ They all react violently with water and organic materials, and dissolve in AHF without decomposition. Characterisation of the adducts by 19F NMR and vibrational spectroscopy give rise to the conclusion that the structure of the cation is pseudo-tetrahedral with C2v symmetry. Validation of the assignment is possible because of the similarity of the spectrum of [C10 2 F2]+ to the spectrum of the isostructural SQ 2 F2. 85 3.6.2. Bromine dioxide trifluoride. Although there have been numerous attempts to synthesise bromine dioxide trifluoride it has never been isolated. Unsuccessful routes include: the fluorination of B r0 2F using KrF2,[19J the fluorination of B r0 3F by [KrF]+[AsF6] '[37l and the hydrolysis of BrF 5 in HF at low temperatures.[56] Mass spectrometry of the hydrolysis products of BrF 5 and BrF 3 has suggested the presence of [B r0 2 F2]+. However, this seems unlikely considering the high energy barrier associated with the conversion of bromine (V) to bromine (VII). 3.6.3. Iodine dioxide trifluoride. Iodine dioxide trifluoride was first obtained by Engelbrecht^57^ in 1969. A twenty-fold excess of H S 0 3F was allowed to react with [Ba3 H 4 (I0 6)2] . The mixture was then distilled under reduced pressure, and the fraction obtained contained HOIOF 4 and H S 0 3F in a 2:1 molar ratio. The addition of oleum converted the HOIOF 4 to I 0 2 F3, which was then separated from the mixture by sublimation, under reduced pressure, on to a cold finger. Iodine dioxide trifluoride is a yellow crystalline s o l i d ^ with a vapour pressure of 5 torr at 25°C. It attacks glass and quartz slowly at room temperature and is photosensitive, producing IOF 3 and 0 2. It is a strong oxidising agent, reacting explosively with organic molecules at room temperature. Iodine dioxide trifluoride exists in two isomeric form s,*^ the ratio of which is solvent and temperature dependent. This molecule does not obey VSEPR theory, which states that the most electronegative elements occupy the axial positions of a trigonal bipyramid. The C2v isomer has both the axial positions occupied by fluorine atoms, whereas the Cs isomer has an oxygen atom in one of these positions (Figure 3.1). 86 Figure 3.1. The isomeric forms of I 0 2F3. H o F F C2v Cs Iodine dioxide trifluoride readily reacts with Lewis bases^60,6^ to form the anion, [I0 2 F4]'; thus, it is observed in the reaction between I 0 2 F 3 and AHF which produces the acid HOIOF4. Alternatively, C s[I0 2 F4] can be prepared by the reaction of C s[I0 4] with either AHF, BrF5, C1F3, C1F5 or F 2 J 62^ Tetrafluoroortho-periodic acid, HOIOF4, attacks glass and quartz at room t e m p e r a t u r e , a n d reacts explosively with organic compounds. Structural characterisation using 19F NMR and vibrational spectroscopies has shown that this molecule exists as two isomers. The 19F NMR spectrum contains a singlet associated with an isomer with the four fluorines in the plane, and a doublet and quartet due to an isomer with three equatorial fluorines and one axial fluorine. The chemistry of the OIOF 4 group has been extended to xenon (II) derivatives, compounds such as Xe(OIOF 4 ) 2 and FXe(OIOF4) demonstrate the high electronegativity of this group and its pseudo-fluorine properties. Iodine dioxide trifluoride reacts with Lewis acids such as AsF 5 and SbF5, to produce 1:1 adductsJ63^ The spectroscopic data for these adducts are not consistent with the ionic formulations [I0 2 F 2 ]+[MF6]" (M = As or Sb). It appears that I 0 2 F 3 acts as an oxygen donor. This type of bonding has also been observed for IOF5 Lewis acid adducts. 87 3.7. Halogen Oxide Pentafluorides, XOF5 . 3.7.1. Chlorine oxide pentafluoride. Chlorine oxide pentafluoride, CIOF 5 , has been reported as a product from the photochemical reaction of C1F5 and OF 2 in a nickel vessel However, no spectroscopic data has been published. 3.7.2 Bromine oxide pentafluoride Four documented attempts have been made to synthesise BrOF5. These are: i) the UV irradiation of BrF 5 and excess of OF 2 between -60 and -40°C j14^ ii) the heating of a mixture of BrF 5 and 0 2 at 4300 psi to 207°C,[37] iii) the hydrolysis of [BrF 6 ]+[AsF6]+ in HF^37^ and iv) the reaction of BrOF 3 and KrF 2 J 19^ However, BrOF 5 still remains unknown. 3.7.3. Iodine oxide pentafluoride. Iodine oxide pentafluoride, IOF5, is formed by the reaction of IF 7 and water, silica or I 2 0 5 J 65"68^ Another route, proposed by Christe and SchackJ69^ employs the use of the ligand transfer reagent POF3, which reacts with IF 7 to produce IOF 5 and PF5. Iodine oxide pentafluoride is a colourless liquid at room temperature with a melting point of 4°C. The molecule has been characterised using electron diffraction^70^ and the structure confirmed as an octahedron of C4v symmetry. Fluorine-19 NMR and vibrational spectroscopic d a t a ^ are in agreement with this result. Iodine oxide pentafluoride reacts with AsF 5 and SbF 5 to form 1:1 adducts. Vibrational and NMR spectroscopic d a t a ^ suggested that the species formed are not the expected donor-acceptor complexes, but are adducts in 88 which the Lewis acid is bonded to the IOF 5 via the oxygen atom. Iodine oxide pentafluoride readily reacts with the Lewis base [NMe4]+F to form the adduct [IOF 6 ]'[N M e 4 ]+J 72,73^ This colourless crystalline solid has a melting point of 172°C, and the structure of the anion, determined by X-ray crystallography^73^ is a pseudo-pentagonal bipyramid of C5v symmetry. 3.8. Halogen Oxide Fluoride, XOF. 3.8.1. Chlorosyl fluoride. Chlorosyl fluoride, ClOF, is the only known halogen (III) oxide fluoride and was first reported in 1930 by Ruff and KingJ74^ The solid melts to a red liquid at -70°C and is unstable in the gaseous state. Infrared spectroscopy has shown[75] that this compound is the primary hydrolysis product of C1F3. In 1974 Andrews et al. demonstrated^76^ that ClOF was formed during the photolysis of 0 3 and C1F in a krypton matrix at -258°C. Vibrational spectroscopic data and normal coordinate analysis provided the means to determine the bond lengths, and a bond angle, F-Cl-O, of 120°. 3.9. Summary. The oxide fluorides of chlorine and iodine have been extensively studied, and most of the neutral complexes have been shown to exhibit amphoteric behaviour. Thus, an oxide fluoride, XOnFm, reacts with a Lewis acid to produce the cation [XOnFm_1]+, whilst, with a fluoride ion donor, it produces the anion [XOnFm+1]". Several of the structures proposed on the basis of VSEPR theory have been confirmed by the structural characterisation of one of the halogen oxide fluoride species. The only deviations appear to be in the case of iodine, where 89 for example [I0 2 F3] is dimeric and, [I0 3 F] and [I0 2 F] are polymeric. This is associated with the transition from the fourth to the fifth row of elements: the inclusion of the Lanthanide series causing an increase in the co-ordination number. 3.10. The Unusual Nature of Bromine (VII). By analogy with chlorine and iodine, bromine (VII) oxide fluoride complexes are expected to exist, and the ease of accessing this valence state should be intermediate between that of chlorine and iodine. However, out of a possible eight bromine (VII) oxide fluorides only one is known. Perchlorate (VII) salts were first prepared via the oxidation of chlorates using sulphuric acid in 1816J771 Trisodium paraperiodate, Na 3 H 2 [I0 6], was first prepared by the oxidation of sodium iodate using chlorine in 1833J78^ However, perbromate was not successfully prepared until 1968,^79^ and then by the exotic means of a hot atom process namely, the (3 decay of radioactive selenium, ( 8 3 Se) (Eqn. 3.1). The precipitation of the perbromate anion as the rubidium salt led to the first isolation of a perbromate salt. [8 3 Se 0 4]2" -> [83Br04r + P" Eqn. 3.1. The first macro scale preparation of perbromates was performed using the electrolytic oxidation of a neutral solution of L i[B r03] . The yield of this reaction was rather poor at 1%. A more successful method involves the use of aqueous XeF 2 which produces [B r04]" in a 10% yield (Eqn. 3.2). [BrO3]' + H20 + XeF 2 -» Xe + [B r04]’ + 2 HF 90 Eqn. 3.2. The most practical synthesis involves the oxidation of [B r03]' using elemental fluorine in aqueous sodium hydroxide (Eqn. 3.3). This preparation^80^ involves a hazardous fluorination stage and a rather lengthy purification. The yields obtained are around 2 0 [B r0 3]" + F 2 + 2 OH" %. -> [B r04]" + 2 P + H20 Eqn. 3.3. The isolation of potassium perbromate involves the neutralisation of perbromic acid with potassium hydroxide. Perbromic acid is a strong monobasic acid, which is stable up to 6 M (55% H B r04), even at 100°C. Concentrated solutions develop a yellow tinge due to the decomposition of trace amounts of hypobromous acid, HOBr. Above 6 M, perbromic acid tends to be erratically unstable although the decomposition is not explosive. The concentration in vacuo, at room temperature, produces an azeotrope which consists of about 80% perbromic acid (ca. 12 M). However, it usually decomposes during or shortly after preparation. The molecular distillation of this azeotrope is possible if heat is applied rapidly in a high vacuum. The bromate-perbromate electrode potential is 1.74 volts in acid solution,^8^ making perbromic acid a potent oxidant (cfi perchlorate E° = 1.23 V and periodate E° = 1.64 V). Dilute solutions react sluggishly with Br" and I" at room temperature, whilst chloride ions are unaffected. However, 6 M perbromic acid attacks stainless steel and, at 100°C, it oxidises Cl" to Cl2, Cr (III) to Cr (VI), Mn (II) to Mn (IV) and Ce (III) to Ce (IV). Potassium perbromate is stable up to 275°C,^80^ at which temperature it decomposes to potassium bromate and oxygen, however, the impure compound may partially decompose at lower temperatures. The numerous unsuccessful attempts to prepare the perbromate ion[82^ and its long-time status as a "non existent" species is surprising, particularly in view of its considerable stability. The oxidation potential of 1.74 V suggests that the use of ozone (E° = 2.07 V) to oxidise bromate to perbromate should be 91 s u c c e s s f u l H o w e v e r , its rather sluggish oxidising nature suggests that a large activation energy exists between Br (VII) and Br (V). Hence, the overall energy required for the formation of [B r04]' from [B r03]' is the sum of the free energy change for the reaction, plus the activation energy. Thus, to overcome the energy barrier only the strongest oxidising agents are successful. To date only two other bromine (VII) species have been isolated. One is perbromyl fluoride, B r0 3 F, the preparation of which involves the use of the super-acid media HF-SbF 5 J 83^ The other, [BrF6]+, can be prepared^84! by the reaction between [KrF]+-[Kr2 F3]+ and BrF5. Both B r0 3F and [BrF6]+ are quite stable at room temperature although B r0 3F is considerably more reactive than its chlorine analogue. The salts of [BrF6]+ are extremely powerful oxidising agents^84] and will oxidise 0 2 to [0 2]+ and Xe to [XeF]+. Table 3.2. Standard electrode potentials (in acid solution) between highest oxidation states of non metals. Periodic Group 15 16 17 Species Involved h 3x o 3 h 3x o 4 h 3 x o 3- h 3 x o 4- xo3- xo4- p -0.28 S 0 .1 0 C1 1.23 As 0.57 Se 1.09 Br 1.74 Sb 0.72 Te 0.90 I 1.64 It can be seen (Table 3.2) that the perbromate anion is a more powerful oxidant than perchlorate or periodate. This situation is not that different to what is found for the elements of Groups 15 and 16 (Table 3.2). As is highlighted the compounds become less potent oxidants to the left of the periodic table, and this reflects the decreased nuclear charge of the species. However, it is 92 apparent, that a large increase in the oxidising power, of the highest oxidation state oxo acids, occurs on going from the last short period (P, S and Cl) to the first long period (As, Se, Br); whereas, only small changes in oxidising power occur between the next periods Prior to the isolation of perbromate salts several explanations were put forward to account for their anomalous absence. Amongst them were inner sphere electron repulsions^85! a high s to p promotion energy^86! and the presence of a node in the bromine 4d orbitalJ87^ the region most favourable for bonding. However, no satisfactory explanation exists for the experimental difficulties that were encountered in the synthesis of high oxidation state bromine compounds. 3.11. Area of Study. There is a scarcity of bromine oxide fluoride compounds compared to their respective chlorine and iodine analogues. However, following work by Gillespie and Spekkens,^21^ all of the bromine (V) oxide fluorides have been prepared, including the corresponding anions and cations. The structures proposed for these compounds were based on infrared, Raman and 19F NMR spectroscopic data, their reactivity and physical properties precluding crystallographic analysis. The bromine (V) oxide fluorides are thermally less stable than their corresponding chlorine and iodine analogues, the reasons for which are unknown. Perbromyl fluoride, B r0 3 F, although less stable than C103F or I 0 3 F, is the only known bromine (VII) oxide fluoride compound. When one considers that all of the neutral chlorine and iodine (VII) oxide fluorides have been prepared and that several ionic adducts are known, this is somewhat surprising. Bromine (VII) species do exist, however, the search for bromine (VII) oxide fluoride species seems to be following the same course as that laid down during 93 the search for perbromate. In this case, although attempts were made to synthesise perbromate, the lack of results soon led to the labelling of the compound as "non-existent”. In fact, it appears that more effort was expended in rationalising these failures than in finding an alternative route. It was felt that a determined effort would lead to the isolation of new bromine (VII) oxide fluorides. A range of very strong oxidative fluorinating reagents and techniques, such as the photolysis of liquid fluorine, were available, and it was envisaged that these might offer new routes to some of the unknown compounds. The use of extended X-ray absorption fine structure (EXAFS) spectroscopy was thought to be an ideal means for the characterisation of these highly reactive molecules. The initial stage of this work was to synthesise some model compounds and compare the EXAFS spectroscopic results with data from known crystal structures. The next stage was to synthesise the known bromine oxide fluoride compounds and to analyse these using EXAFS spectroscopy. The compounds prepared would also be suitable starting materials for the synthesis of new bromine (VII) oxide fluoride species. In addition, all the compounds prepared were to be characterised by multinuclear NMR and matrix-isolation infrared spectroscopy. 94 3.12. EXAFS Spectroscopic Study of the Bromine Fluorides. Bromine trifluoride, BrF3, and bromine pentafluoride, BrF5, are stable liquids at room temperature. Bromine trifluoride is a pale yellow liquid, and was first prepared in 1900 by Moissan ^8 8 1 by the reaction between bromine and excess of fluorine at room temperature. Bromine pentafluoride is a colourless liquid with a high vapour pressure at room temperature. It was first prepared in 1931 by Ruff and Menzel ^8 9 1 by the reaction of bromine trifluoride and fluorine at 200°C in a platinum apparatus. An explosion occurred, which was caused by corrosion of the platinum vessel, and the subsequent release of the reactants into the oil bath. They finally successfully prepared BrF 5 by the direct combination of bromine and fluorine in a copper apparatus at 200°C. As with the bromine oxide fluorides, the bromine fluorides are amphoteric. They react readily with Lewis acids or bases to produce the corresponding adducts. The adducts [BrF 2 ][AsF 6 ] [ 9 2 , 9 3 1 and [BrF 4 ][Sb 2 F n ] [9 4 ,9 5 1 K[BrF4] J 9 0 1 Cs[BrF6] J 9 1 1 were prepared as described in the experimental. The adduct [BrF 2 ][AsF6] slowly decomposes over a period of time to a red-coloured solid. This may be due to dissociation and the red colour is presumably due to the presence of bromine. The solids Cs[BrF6],^91^ [BrF 2 ][AsF 6 ] [ 9 5 1 an(j [BrF4 ][Sb 2 F u ] [9 4 1 have previously been characterised using single crystal X-ray diffraction, while potassium tetrafluorobromate (III), K[BrF4], has been characterised ^9 0 1 using single crystal neutron diffraction. The compounds K[BrF4], Cs[BrF6], [BrF 2 ][AsF6] and [BrF 4 ][Sb 2 F n ] were loaded into FEP cells and diluted with fully-passivated Teflon. The bromine K-edge EXAFS data were collected at room temperature in A'1(k = photoelectron wave vector). This was later truncated to 13.5 A for [BrF ][Sb F n ] and 13 A"1 for [BrF ][AsF6], Cs[BrF6] and K[BrF4] due to increased noise at higher k. Three data sets were averaged in each case and the data multiplied by k3 to compensate for a transmission mode, out to k= -1 15 4 2 95 2 decrease in intensity at higher k. No Fourier filtering was applied and the fits discussed were compared with the averaged raw (background subtracted) EXAFS data. The analysis was modelled using EXCURV92^96^ to one shell for the anions and two shells for the cations (Figures 3.2 to 3.5). Each shell was tested for statistical significance J 97^ VPI and AFAC were mapped for the compounds and the values obtained were identical for each. These values should be comparable with the values obtained for other bromine species in these types of environment and facilitate the definitive characterisation of any novel species prepared. The parameter VPI takes into account inelastic losses and the core hole lifetime. VPI is always negative and decreases with increasing edge energy. For the first long period of elements it is found to fall in the range -1 to -2. AFAC is the proportion of electrons which perform an EXAFS type scatter and it is usually found in the range 0.7 to 0.9. 3.12.1. Discussion. A comparison of the X-ray crystallographic and EXAFS spectroscopic data (Table 3.3) demonstrates that EXAFS spectroscopy is a suitable technique for determining internal bonding parameters for bromine fluorides. Figures 3.2 to 3.5 show representative examples of the background-subtracted EXAFS and the Fourier transform spectra for the compounds K[BrF4], Cs[BrF6], [BrF 2 ][AsF6] and [BrF 4 ][Sb 2 F 11]. The bond lengths obtained using EXAFS spectroscopy are in good agreement with the values obtained from the single crystal studies. Table 3.3 highlights the expected increase in the Br-F bond lengths on going from Cs[BrF6] to K[BrF4]. This is attributed to the change in oxidation state and shows that the anionic Br-F bond lengths may be expected in the region 1 .8 to 1.9 A. The EXAFS data for the cationic bromine fluorides, [BrF2]+ and [BrF4]+, demonstrates that Br-F bond lengths are in the region of 96 Table 3.3. EXAFS and X-ray crystal data for K[BrF4] and Cs[BrF6], [BrF 2 ][AsF6] and [BrF4 ][Sb 2 F n ]. K[BrF4] Compound Crystal </(Br-F) \ A 1.89(2) Fit Index 0 Rd o2 1.847(1) EXAFSf 1.85(1) 0.006(2) 0.007(2) 2 .8 2 .1 23.4 20.3 [BrF 2 ][AsF6] [BrF4 ][Sb 2 F n ] EXAFSg Crystal EXAFSh A \ Ab 1.69(2) 1.70(1) 1.81(11) 1.69(1) A 2.29(2) (/(Br-F) \ 2 Crystal Crystal (/(Br-F) \ 2 EXAFSe 1.88(1) 2a2 \ Ab Compound Cs[BrF6] <J2 \ Ab Fit Index 0 Rd 0.003(2) 2.35(1) 0.007(1) 2.36(10) 2.40(1) 0.024(2) 0.028(3) 2 .0 2 .0 16.9 18.9 a Standard deviations in parentheses. Debye-Waller factor. c Fit index = Zj[(% T-%E)ki3]2. d R = [s(XT-XE)£3 dk/sXEk 3 dk] x 100 %. e E0 9.75 (0.34), AFAC 0.71 and VPI -1.86. f E0 6.29 (0.29), AFAC 0.71 and VPI -1.86. g E0 4.11 (0.40), AFAC 0.71 and VPI -1.86. h E0 6.14 (0.33), AFAC 0.71 and VPI -1.86. 97 Figure 3.2. a) Background-subtracted EXAFS and b) the Fourier transform spectra for K[BrF4]. -8 -1 2 ° - i k/A 1 (b) 1.8 CO k. < 0.6 2 4 6 8 o_ r/A aEXAFS ( experimental x k3, — curved-wave theory x k3) bFourier transforms ( experimental, — theoretical) 98 10 Figure 3.3. (a) Background-subtracted EXAFS and (b) the Fourier transform spectra for Cs[BrF6]. (a) -10 (b) 3.0 (0 C 2.0 D >. w (0 3 k - r/A aEXAFS ( experimental x k3, — curved-wave theory x k3) bFourier transforms ( experimental, — theoretical) 99 Figure 3.4. (a) Background-subtracted EXAFS and (b) the Fourier transform spectra for [BrF2][AsF6]. (a) 6 0 6 0 k/A 1 (b) 1.0 c 3 (0 1 - 4-» 3 0.5 < v - ■ r/A aEXAFS (----- experimental x k3, — curved-wave theory x k3) bFourier transforms ( experimental, — theoretical) 100 Figure 3.5. (a) Background-subtracted EXAFS and (b) the Fourier transform spectra for[BrF4] [Sb2F xx]. CO -12 (b) 2.0 0) 1.0 < 2 aEXAFS ( 4 6 8 experimental x k3, — curved-wave theory x k3) bFourier transforms ( experimental, — theoretical) 101 10 1.69 A, and the formal oxidation state of the bromine does not significantly affect the bond length. It is noted that the single crystal data presented for [BrF4]+ is not particularly good (R=0.14). The crystallographic experimental data did not allow precise determination of bond lengths and angles, however, at the time this was considered unimportant as the structure determination provided information about the overall structure of this interesting adduct. A comparison of the crystallographic and EXAFS data (Table 3.3) for [BrF4]+ shows that the crystallographically determined bond length of 1.81(11) A was not very precise and, more importantly, is indicative of an anionic Br-F bond distance. The value obtained using EXAFS spectroscopy, 1.69(1) A, provides a far more meaningful representation of the Br-F bond length. For the cations, [BrF2]+ and [BrF4]+, EXAFS spectroscopy is able to verify the presence of bridging fluorines at 2.35(1) and 2.40(1) A respectively. As can be seen in the EXAFS spectra presented in Figures 3.2 to 3.5, extra shells were observed in each case at distances above 3 A from the central bromine atoms. Attempts were made to fit these shells and, in each case, the result showed an improvement in the R value. The shells, however, were found to be statistically in sig n ific a n t^ and were not included in the fits presented in this Chapter. The hexafluorobromine (VII) cation was synthesised by Gillespie and Schrobilgen[98' 10°l in 1976. Although no crystallographic data are available, characterisation by 19F NMR and Raman spectroscopy is reported. The 19F NMR spectra recorded in HF at room temperature showed two overlapping 1 : 1 : 1 :1 quartets at 8339.4 ppm. The two quartets are assigned to [7 9 BrF6]+ and [8 1 BrF6]+ and arise from spin-spin coupling of the six equivalent fluorines with 79Br and 81 Br, both with 7=3/2. The equal intensities of the quartets is in accordance with the natural abundance of the two bromine isotopes (7 9 Br, 50.57% and 81Br 49.43%) and the ratio / ( 1 9 F- 8 1 B r):/( 1 9 F-7 9 Br) is in agreement with the gyromagnetic ratios y8 1 Br:y79Br = 1.0778. 102 The hexafluorobromine (V) anion has been characterised recently ^ 1 0 1 ^ and the low temperature (-40°C) 19F NMR spectrum shows the same features as those described above. This is contrary to previous NMR experiments^102^ which failed to observe Br-F coupling even at temperatures of -60°C. The fluorine resonances occur at lower frequency, 5100.6 ppm, which is a result of the lower oxidation state of the central bromine. The observation of this coupling indicates an octahedral structure implying that the valence electrons are occupying the sterically inactive As orbital. Bromine is thought to have a maximum co-ordination number of six, the same as chlorine, whereas iodine has a maximum co-ordination number of 8 . This was used to explain the reactions of [BrF6]+ and [IF6]+ with NOF (Eqn.’s 3.4 and 3.5). [BrF 6 ]+[AsF6]' + NOF [IF 6 ]+[AsF6]" + NOF -> -> BrF 5 + F 2 + [NO]+[AsF6]’ IF7 + [NO]+[AsF6r Eqn. 3.4. Eqn. 3.5. However, the crystal structure of [BrF4 ]+[Sb2 F 11]' in d ic a te s ^ that the bromine centre has a co-ordination number of seven. Distortion within the cation indicates the presence of a sterically active lone pair of electrons, therefore, the crystal structure showed the bromine to be pseudo-hepta coordinate. The crystal data showed four fluorines bound at an average distance of 1.81(11) A and two bridging fluorines, to the neighbouring [Sb2 Fn]" molecules, at a distance of 2.36(10) A. The EXAFS data has shown that the BrF terminal distances are 1.69(1) A. This seems to be a more realistic value in the context of the nature of the charge on the species. For the cations [BrF2]+ and [BrF4]+ the lone pairs of electrons appear to adopt sterically active roles, contrary to what is found in [BrF6] \ This maybe a reflection of the inability of bromine to exist with a co-ordination number of seven. Further work is needed to establish whether bromine does exist with a 103 co-ordination number of seven. The steric crowding in BrF 7 would presumably be less than that for [BrF4]+ as, according to VSEPR theory, a non-bonding pair of electrons would exhibit greater repulsion than a bromine fluorine single bond. 3.13. The Synthesis and EXAFS Characterisation of Caesium Bromine-Oxide Tetrafluoride. The existence of K[BrOF4], has been reported by both B o u g o n ^ and G illespieJ17^ Bougon obtained K[BrOF4] by the reaction of K [B r03] with a large excess of BrF 5 at 80°C in the presence of fluorine. Although this method reportedly yields a pure product, the course of the reaction is difficult to control and K[BrF4] is usually obtained as the only product. Gillespie reported that the reaction between K[BrF6] and K [B r03] in CH3CN solution produces a mixture of K [B r0 2 F2] and K[BrOF4] (Eqn. 3.6). The separation of the two products relies on the solubility of K[BrOF4] in CH3CN compared with the insolubility of K [B r0 2 F2]. K[BrF6] + K [B r03] -> K[BrOF4] + K [B r0 2 F2] Eqn. 3.6. In 1978, Christe reported an improved synthesis of [BrOF4]‘ salts The reaction of K [B r04] or C s[B r04] with BrF 5 and F 2 led to K[BrOF4] and the previously unknown Cs[BrOF4]. Although this method results in essentially pure products in high yield, the required [B r04]" salts are difficult to prepare. The reaction of the caesium salt, C s[B r04], occurs at room temperature over thirty hours with 1 0 0 % conversion, whereas, the potassium salt requires ninety five hours at 80°C with only 70% conversion. In 1987 Christe and Wilson proposed a new one-step synthesis to [BrOF4]‘ saltsP 6^ They reported that the reactions of an excess of BrF 5 with 104 the alkali-metal nitrates M[NOs] (M = Na, K, Rb or Cs) provided a new, simple high yield route to the corresponding [BrOF4]' salts and F N 0 2. The heavy alkali metal salts (K, Rb and Cs) of [BrOF4]' form at temperatures as high as 100°C. For the N a[N 0 3 ]-BrF 5 system at 0°C some BrOF 3 was obtained along with Na[BrOF4]. At 25°C any BrOF 3 formed undergoes either fast decomposition to BrF 3 and 0 2 or further reaction with N a[N 03] to produce B r0 2 F. This then complexes with the NaF to yield N a[B r0 2 F2]. It was noted that the formation of BrOF 3 cannot be the result of the decomposition of Na[BrOF4] as the salt is stable up to 160°C. Instead, it must be formed from a less stable intermediate that is capable of generating BrOF3, MF, F N 0 2 or M[BrOF4] and F N 0 2. It was concluded that the reaction must go via the intermediate [N 0 3 -BrF5]". Decomposition of the resulting M+[N 0 3 -BrF5]‘ complex could involve either F N 0 2 elimination from the anion yielding M+[BrOF4]" or fluoride abstraction from the [N 0 3 *BrF5]‘ anion by M+. The [BrF 4 0 N 0 2] is presumably unstable and would eliminate F N 0 2 to produce BrOF3. If this is the case, then the reaction pathway would depend on the fluoride ion affinity of M + and the thermal stability of M+[N 0 3 *BrF5] \ On the basis of the above reasoning, the mechanism shown in Figure 3.6 was proposed P® Figure 3.6, Proposed reaction scheme for the reaction between bromine pentafluoride and the alkali metal nitrates. M^NOa' + BrF5 105 The salt Cs[BrOF4] was prepared by the reaction of BrF 5 and C s[N 03] and the stable white solid produced was stored in an inert atmosphere dry box. Analysis by infrared spectroscopy showed the presence of a strong absorption at 925 cm - 1 and a very broad band in the range 562-443 cm"1. These absorptions are characteristic of the [BrOF4]' anion and vary only slightly for the different alkali metal saltsP® The sample was loaded into a FEP cell and diluted with passivated Teflon. The bromine K edge EXAFS data were recorded at room temperature in transmission mode out to k = 15 A"1 (k = photo electron wave vector). This was later truncated to 13.5 A"1 due to poor signal-to-noise ratios at higher k. Six data sets were collected, averaged, and then multiplied by k 3 to compensate for the drop-off in intensity at higher k. No smoothing or Fourier filtering was applied, and the fit discussed below was compared with the averaged raw (background subtracted) EXAFS data. The data was modelled using EXCURV92J961 for two shells of one oxygen atom at 1.58(1) A and four fluorine atoms at 1.87(1) A (Table 3.4). Each shell was added stepwise and the fits tested for statistical significance J 9 7 1 As with the EXAFS studies of the bromine fluorides, extra shells are evident at distances greater than 3 A from the central bromine atom. Inclusion of extra shells resulted in an improvement of the R value for the experiment. However, the shells were statistically insignificant and not included in the fit discussed here. The Br-F bond distances compare well with those obtained for the anionic bromine fluorides (Section 3.12). The expected double-bond character o between bromine and oxygen is reflected by the bond distance of 1.58(1) A. A similar EXAFS spectroscopic study was carried out on K [B r04](S), N a[B r03](s), N a[B r02](s) and Na[BrO](aq),^103^ where the Br-O bond distances were 1.61(2), 1.65(2), 1.75(2) and 1.81(2) A respectively. As can be seen, the Br-O distance for Cs[BrOF4], 1.58(1), is slightly shorter than that observed for 106 K [B r04]. The introduction of four electron-withdrawing fluorine atoms is expected to reduce the Br-O bond lengths. Table 3.4. EXAFS data for Cs[BrOF4]. EXAFS Datae rf(Br-0 ) / A 1.58(1) 2 c 2 / Ab 0.007(1) rf(Br-F) / A 1.87(1) 2a2 / Ab 0.006(1) Fit Index 0 2.29 Rd 18.0 a Standard deviations in parentheses. b Debye-Waller factor. c Fit index = Xj[(% T-XE)ki3]2. d R = [s(XT-%E)k 3 dk/s%Ek 3 dk] x 100 %. e E0 7.25 (0.40), AFAC 0.71 and V P I-1.86. 107 Figure 3.7. (a) Background-subtracted EXAFS and (b) the Fourier transform spectra for Cs[BrOF4]. M -6 o - k /A 1 1.8 Z> 12 < 0.6 2 6 4 8 r/V aEXAFS ( experimental x k3, — curved-wave theory x k3) bFourier transforms ( experimental, —- theoretical) 108 10 3.14. The Synthesis of Potassium Perbromate. The synthetic routes to the perbromate anion are outlined in Section 3.10. Oxidation of the bromate anion in sodium hydroxide solution using elemental f l u o r i n e p r o v i d e s the best route (Eqn. 3.7). [BrO3r + F 2 + 2 OH* [B r04]' + 2 F + H20 Eqn. 3.7. Initial attempts necessitated a slow rate of addition of fluorine and led to reaction times of 2 to 3 days. However, it was realised that it is necessary to complete the fluorination and isolation of the product within a single day. Faster flow rates were attained with the construction of a new metal vacuum line in a well vented, low population area and the fluorination stage was cut to a few hours. However, care had to be taken to cool the reaction mixture because the fast rates of addition of fluorine have the potential to lead to exotherms and ignition of the solvent vapour. Potassium perbromate was dried under dynamic vacuum at 150°C for 12 hours, and stored in an inert atmosphere dry box. An infrared spectrum of the solid as a Nujol mull showed its characteristic absorptions to be present at 798 and 410 cm-1. The 81Br NMR spectrum (7=3/2, 49.43% abundance) was recorded for K [B r0 4] and referenced to aqueous KBr (1 mol dm-3), the values being corrected to infinite dilution using the published dataJ103^ The 81Br NMR resonance of [B r04]" was observed at 52470, relative to infinitely dilute aqueous Br", in good agreement with the reported^103^ value of 52476 ppm. 109 3.15. The Synthesis of Perbromyl Fluoride. The synthesis of perbromyl fluoride was first reported by Appleman et al. in 1 9 6 9 and involved the reaction between potassium perbromate and antimony pentafluoride in anhydrous hydrogen fluoride. The compound is highly volatile and possesses a vapour pressure of 56 torr at -50°C. Due to the problems encountered during the synthesis of the other bromine oxide fluoride compounds, the reactions between K [B r04] and SbF5, SbF 5 -AHF, BrF5-AHF or AsF5-AHF were re-examined to establish which provided the most convenient route to B r0 3 F. Using 19F NMR spectroscopy, the reactions were all shown to produce B r0 3 F. The preferred route employed the use of AHF and BrF5, the reason for this was solely the ease of transfer of AHF and BrF 5 as opposed to the less volatile and more viscous SbF5. Perbromyl fluoride decomposes slowly at room temperature and this was evidenced by the formation of Br2. This discouraged us from using trap to trap distillation as a means of purification. Instead, the reaction products were distilled at low temperature, -84°C, at which temperature the vapour pressures of B r0 3 F, AHF and BrF 5 are 5, 2 and 0 torr respectively. The volatile products were condensed, under static vacuum, into a second FEP tube which contained dried NaF. The NaF formed an involatile adduct with trace amounts of HF, whereas, B r0 3F and NaF did not react. As already stated, B r0 3F decomposes slowly at room temperature to produce Br2. The Br 2 can be readily removed by the condensation of F 2 into the FEP tube at 196°C. On warming to room temperature the Br 2 is oxidised to BrF3, which then reacts with NaF to form an involatile adduct. Therefore, the NaF serves two purposes, the removal of HF and BrF3, and facilitates the preparation of pure B r0 3 F. No decomposition was observed to occur if the B r0 3F is stored at liquid nitrogen temperatures. The reaction between K [B r04] and BrF5, using AHF as the solvent has been previously investigated^21^ The reaction apparently occurs according to 110 Equation 3.8. Figure 3.8 shows the 19F NMR spectrum recorded for the reaction medium at -59 °C. The AX 4 pattern generated by the BrF 5 was unresolved at room temperature. However, at -59°C, the multiplicity was resolved: 8 Fa 271.9 and 8 Fe 135.2 ppm. Also apparent was a resonance at 8274.2 ppm due to the presence of B r0 3F (cf. neat B r0 3 F, 8274 ppm, -80 °C )J3TI 2 K [B r0 4] + BrF 5 + 2 HF -> 2 B r 0 3F + B r0 2F + 2 K[HF2] Eqn. 3.8. The presence of B r0 2F was used by Spekkens^2^ to explain the poor resolution observed for the 19F NMR spectrum of the reaction media. Fluorine19 NM R experiments performed by Spekkens failed to detect the presence of B r0 2 F, although, it was detected using Raman spectroscopy. Gillespie and Spekkens report the 19F NMR chemical shift for B r0 2F to be 8210 ppm when recorded as a solution in BrF 5 J 37^ The presence of B r0 2F cannot be verified using low temperature 19F NMR spectroscopy, this seems strange in view of the smooth base line observed in the region of 8210 ppm and the high resolution of the spectrum, discounting the presence of any fluxionality. Therefore, the reaction described here clearly does produce B r0 3 F, however, whether this is exactly as outlined in Equation 3.8 is doubtful. The reaction of K [B r04] with XF 5 -AHF, where X = Sb or As, is thought to go via a different reaction pathway. However, none of the intermediate species have been observed. Work has demonstrated that the reaction is not that shown in Equation 3.9. K [B r0 4] + AsF 5 —» B r0 3F + K[AsOF4] Ill Eqn. 3.9. Figure 3.8. Fluorine-19 NMR spectrum of B r0 3F in BrF5. to 276 274 272 (ppm ) 270 268 140 138 136 134 (ppm ) 132 130 Raman spectroscopy has shown that neither AsOF 3 nor K[AsOF4] are present^21! An alternative mechanism is that proposed for the formation of CIO 3 F from [CIO4 ]" under the same conditions.^ Although numerous speculative opinions have been expressed for this system, it seems unlikely that the formation of B r0 3F involves a mechanism where [B r03]+ is an intermediate. Furthermore, the high yields, > 96 %, would not be expected in view of the likely instability of [B r03]+. It seems more likely that the mechanism involves protonated perbromic acid as shown in Scheme 3.1. Scheme 3.1. Proposed reaction pathway for the formation of perbromyl fluoride. 4 HF + 2 XF 5 -> 2 [H2 F]+ + [B r04r 2 [H2 F]+ + 2 [XF6]‘ -> [H2 O B r03]+ + HF o v erall: [B r04]' + 3 HF + 2 X F 5 [H2 O B r03]+ + 2 HF -» -> B r0 3F + [H3 0 ] + B r0 3F + [H3 0 ] + + 2 [XF6]' X = Sb or As. The formation of perbromyl fluoride is more likely to involve the nucleophilic displacement of H20 by F , rather than the heterolytic cleavage of the Br-OH 2 bond in [H2 0 B r 0 3]+ to give H20 and [B r03]+ followed by reaction of [B r0 3]+ with HF. 113 3.16. The Attempted Synthesis of Bromine Oxide Trifluoride. The reaction between the heavy alkali metal nitrates, M [N 03] (M = K, Rb or Cs), and bromine pentafluoride produces the corresponding tetrafluorobromate (V) salt. The yield for this reaction is quantitative for the caesium reaction and decreases as Group 1 is ascended. It was noted by Christe et al. that the reaction between N a[N 03] and BrF5^36^ did not produce Na[BrOF4] as the only product. They observed that BrOF 3 was also formed, along with some Na[BrF4], which was presumably formed via the decomposition of the BrOF 3 as shown in Equations 3.10 and 3.11 (Figure 3.6). BrOF 3 —» BrF 3 + V i0 2 NaF + BrF 3 -> Eqn. 3.10. Na[BrF4] Eqn. 3.11. The route by which the above reaction occurs (see Section 3.13) depends solely on the fluoride ion affinity of the alkali metal. On the basis of hard-soft acid-base principles, it was reasoned that L i[N 03] should form BrOF 3 in the highest yield due to the high fluoride affinity of the lithium cation. Indeed, Christe et al. reported that when L i[N 03] was allowed to react with an excess of BrF 5 at 0 °C over twenty days, BrOF 3 was formed in essentially quantitative yield. The reaction was attempted at 0°C over a period of twenty days in a nickel reaction vessel which was shaken several times daily. The vessel was attached to a metal line and the volatile materials were distilled through a trap at -64°C under dynamic vacuum. At this temperature, BrF 5 and F N 0 2 do not collect and any solid trapped should have been due to the presence of BrOF3, but no solid was collected. Five attempts were made to synthesis BrOF3. However, 19F NMR spectroscopy of the residual solid dissolved in CH3CN 114 showed only the presence of [BrF4]', suggesting that BrOF 3 may have been formed but had subsequently decomposed according to Equation 3.10. Similar results were observed when the reaction was carried out at -10 and -20°C and with larger excesses of BrF5. A second approach for the production of bromine oxide trifluoride was attempted. This involved the reaction between K[BrOF4] and a weak Lewis acid J17^ Bromine oxide trifluoride can be obtained by dissolving K[BrOF4] in AHF at low temperature. On wanning the K[BrOF4]-HF mixture to room temperature BrOF 3 and K[HF2] are formed (Eqn. 3.12). K[BrOF4] + HF —» BrOF 3 + K[HF2] Eqn. 3.12. The HF can be removed by pumping the mixture at -60°C to leave a mixture of BrOF 3 and K[HF2]. The BrOF 3 cannot simply be removed under dynamic vacuum as the reaction is reversible. This problem can be overcome by condensation of BrF 5 on to the mixture which solubilises the BrOF 3 but not the K[HF2] . Decantation of the solution affords separation. The BrF 5 can then be removed at low temperature to leave BrOF3. The problem with this preparation is that a pure source of K[BrOF4] is required. Gillespie and Spekkens described that a mixture of K[BrOF4] and K [B r0 2 F2] could be obtained from the fluorine exchange reaction between K [B r0 3] and K[BrF6] in CH 3 C N J17^ The separation of the two products relies on the slight solubility of K[BrOF4] over the apparent insolubility of K [B r0 2 F2] in CH 3 CN. A mixture of K [B r0 2 F2] and K[BrOF4] and CH3CN was shaken for two hours at room temperature. The liquid was then filtered into a second vessel. On the basis of what Gillespie and Spekkens reported removal of the CH3CN should have produced a white solid. In our hands, however, no white solid was deposited on any of the several times the reaction was repeated. 115 3.17. The attempted preparation of Bromyl Fluoride. Two synthetic procedures were investigated as a possible route to bromyl fluoride, B r0 2 F. The reaction between K [B r03] and BrF 5 and a catalytic amount of HF reportedly produces K [B r0 2 F2], K[BrOF4] and B r0 2 F J 15,16,17^ The above mixture was stirred at room temperature for two hours. The resultant mixture was brown, consistent with Gillespie's observations, and was presumably so due to the formation of bromine. The volatile materials were pumped under a dynamic vacuum through a trap at -48 °C (n-hexyl alcohol / C 0 2 (S)), at which temperature B r0 2F should have condensed. No material was collected in the trap indicating that B r0 2F had not been formed. During the reactions a large volume of gas was evolved which was presumably oxygen since it was not totally condensable at liquid nitrogen temperatures. Analysis of the solid product by 19F NMR spectroscopy showed only the presence of [BrF4] \ The solvent used was CH 3 CN, in which K [B r0 2 F2] is insoluble and K[BrOF4] is only slightly soluble. Infrared analysis of the solid was uninformative due to the large number of absorptions which coincided and the broadness associated with them. The reaction was repeated four times with no change in the result. The second synthetic route relies on the isolation of K [B r0 2 F2] . As was the case in Section 3.16, the reaction between K [B r0 2 F2] and HF yields B r0 2F and K[HF2].[17] Unlike the K[HF 2 ]-BrOF 3 mixture, B r0 2F can be removed from the K[HF 2 ]-B r0 2F mixture without the reverse reaction occurring; indicating that BrOF 3 is a stronger fluoride ion acceptor than B r0 2 F. As a consequence of not being able to separate the mixture of K [B r0 2 F2] and K[BrOF4] this reaction could not be attempted. 116 3.18. Conclusion. Little progress has been made in the synthesis of new bromine oxide fluorides. However, it was felt that significant knowledge had been gained about approaches that had the potential to be successful with a concerted effort. However, as was envisaged, the use of EXAFS spectroscopy provided an excellent means by which to obtain structural data on this fundamentally important class of compounds. EXAFS spectroscopy was successfully used to characterise the model compounds K[BrF4], Cs[BrF6], [BrF 2 ][AsF6] and [BrF 4 ][Sb 2 F 11]. The expected trends in Br-F bond lengths have been clearly illustrated and the improvement in the data available for [BrF4 ][Sb 2 F n ] demonstrates the problems encountered when trying to characterise these type of compounds using single crystal techniques. The use of EXAFS spectroscopy has permitted the anion [BrOF4]" to be structurally characterised for the first time. The Br-F bond lengths are in good agreement with established trends highlighted for the bromine fluoride ions and the Br-O bond length also seems very reasonable. The isolation of K [B r04] proved difficult. However, once isolated, the synthesis of B r0 3F opened a potential route to new bromine (VII) oxide fluorides. The possibility of performing EXAFS spectroscopy on low temperature bromine oxide fluorides also offered the potential to obtain structural data on otherwise, highly reactive, gaseous materials. Two major problems were encountered mid-way through the second year of research. The use of EXAFS spectroscopy relies on the allocation of time at the synchrotron radiation source in Daresbury. No time was allocated to the Leicester Fluorine Group during this year and, consequently, the characterisation of new species which had been made was not possible. Secondly, new regulations about the transportation of BrF5, which is a vital reagent in this area of chemistry, meant that this was no longer accessable, and 117 the delivery of a new cylinder would have taken in excess of a year. As a consequence of this, and the very slow progress being made within this area, a decision was made to stop the work and move into the areas described in Chapters Two and Four. However, as outlined in Section 3.11, a lot of work is still needed to determine whether this class of compounds can be expanded. A determined effort would undoubtedly unearth new synthetic routes to bromine (VII) oxide fluorides. The generation of new bromine (VII) oxide fluorides will take severe conditions, but the use of [KrF]+ or [Kr2 F3]+ and the laser photolysis of liquid fluorine are still to be investigated. These reagents and techniques such as low temperature infrared and EXAFS spectroscopy offer the most realistic chances of success. 118 References Chapter Three [1] R. 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Mahjoub, X. Zhang and K. Seppelt, Angew. Chem. Int. Ed. Engl., 1995,1, 261. [102] K. O. Christe and W. W. Wilson, Inorg. Chem., 1989, 28, 3275 and references cited therein. [103] W. Levason, J. S. Ogden, M. D. Spicer and N. A. Young, J. Chem. Soc., Dalton Trans., 1990, 349 and references cited therein. 124 CHAPTER FOUR Displacement and Oxidation Reactions using Fluorosulphonic Acid 4.1. Introduction. Fluorosulphonic acid, HSO 3 F, was first prepared by Thorpe and Kirman in 1 8 9 2 ^ by the reaction of S 0 3 with HF (Eqn. 4.1). S 0 3 + HF H S 0 3F Eqn. 4.1. Fluorosulphonic acid has found extensive uses as a catalyst and reagent in the fields of inorganic and organic chemistryJ2'4^ A number of inorganic fluorides can be prepared by the reaction of oxides, hydroxides and the salts of oxo acids with fluorosulphonic acid, for example, P 4 O 10, B(OH ) 3 and K M n0 4 react with it to produce POF3, BF 3 and M n 0 3F respectively. Within the area of organic chemistry, many of the processes involving fluorosulphonic acid have been patented. Its diversity of uses is exemplified by its role as a catalyst in alkylation, polymerisation and isomerisation reactions. A number of organic derivatives have been prepared and these include aryl and alkyl fluorosulphates. It is also used in various refining processes such as the removal of organofluorine compounds from hydrocarbons, the refining of lubricating oils and the removal of metals from crude petroleum. 4.2. Properties. Fluorosulphonic acid is isoelectronic with H 2 S 0 4, one fluorine atom replacing an -OH group. As a consequence of one less hydroxyl group, the degree of molecular association, in the form of intermolecular hydrogen bonds, is considerably lower and this markedly affects its physical properties (Table 4.1). Its melting point is lower than that of sulphuric acid, and this has enabled NMR spectra to be recorded at temperatures low enough for protonation 125 reactions to be observed. An example of this is the protonation of acetam ide,^ (see Figure 4.1). Table 4.1. Physical properties of H 2 S 0 4, HSO 3 F, S 0 2 F2, CF 3 SO 3 H and HF. P ro p erty ^ h 2so4 HSO 3 F s o 2f 2 CF 3 SO 3 H HF Mw 98.1 1 0 0 .1 1 0 2 .1 150.0 2 0 .0 M. pt. / °C 10.4 -87.3 -136.7 162 -83.1 B .p t./° C 338 165.5 -55.4 34 19.5 Density / gem " 3 1.841 1.743 - 1.69 0.987 -H 0 1 2 .1 15.1 - 15.1 15.1 Figure 4.1. The protonation of acetamide. H HSOgF^ CH3C. CH 3C nh 2 NH At room temperature the lH NMR spectrum consists of two peaks which originate from the methyl and amide protons. On cooling to -92°C a third peak appears in the spectrum at high frequency. The three peaks have the relative intensities 1:2:3 and are assigned to OH, NH 2 and CH 3 groups, demonstrating that protonation occurs at the oxygen rather than at nitrogen. Fluorosulphonic acid, along with anhydrous hydrogen fluoride and trifluoromethanesulphonic acid (commonly referred to as triflic acid), are the strongest monoprotic acids known (Table 4.1). To differentiate between these acids is difficult, mainly because their acidities are extremely solvent dependent and the precise measurement of Hammett acidity functions is difficult J7-9^ 126 Compared to hydrogen fluoride, fluorosulphonic acid and triflic acid possess large liquid ranges and are easily purified by distillation at atmospheric pressure. They are also compatible with glass which means synthetic procedures and spectroscopic characterisation are straightforward. Although fluorosulphonic acid and triflic acid share comparable physical properties, differences in their chemical behaviour has led to a far greater use of triflic acid in the areas of synthesis and c a t a l y s i s T r i f l i c acid and its conjugate base possess an extremely high thermal stability and a resistance to both reductive and oxidative cleavage. This non-oxidising behaviour minimises the number of potential side reactions, and reduces the hazards associated with the use of strong oxidising acids such as perchloric acid. Although triflic acid fumes in moist air, it is completely miscible in water and many polar organic solvents. These properties and the highly labile nature of the triflate group have resulted in an extensive amount of literature, and this was most recently reviewed in 1977 HO] Fluorosulphonic acid undergoes an immediate and vigorous reaction when hydrolysed and this occurs via two pathways Part of the acid hydrolyses rapidly to HF and H 2 S 0 4, the extent of which depends on the rate of addition and the temperature. The remainder of the acid forms the hydroxonium salt, [H 3 0 ]+[S0 3 F]", which then undergoes slow hydrolysis. W oolf et alSn ^ have prepared aqueous solutions of the almost unhydrolysed acid. They also report that when excess water is added to fluorosulphonic acid, it does not completely react to form HF and H 2 S 0 4 but instead, a stable equilibrium is rapidly achieved. 127 Figure 4.2. Variety of fluorosulphate derivatives. B inary M ono & B identate A '' v ’i'V -V / ' 1 t V [Zn(S03F)2][14! T e rn a ry [Nb(S03F)5]ll5) Cs[Sb(S03F)6][|6) 3A /3J [Br(S03F)3][>3] Cs2[Pt(S03F)6]l16l B identate H etero-bim etallic [c-Pd2(n-C0)2][(S03F)2][>7] 1 7 l l l A V * A C 111 1UO I O SU I .............. * Vi n tc pna [Sn(S03F)2(CH3)2][181 [AgSn(S03F)6]P>l [CuSn(S03F)6]P» T rid e n ta te M onodentate [Fe(S03F)2]P21 [Br(S03F)4]li3] [Au(S03F)4]-[|61 ........ H igh O xidation S tate [Re02(S03F)3][19l [XeF5(S03F)]P°l [Ni(S03F)2]P21 4.3. Synthetic Routes to Metal Fluorosulphate Complexes. Derivatives containing the fluorosulphate group occur across most of the periodic table and the extent of this has been highlighted in several reviews.^2,3,23,24] This ability to show such diversity is attributed to its relatively high thermal stability and its versatile co-ordinating ability (Figure 4.2). The majority of these complexes can be prepared using S2 0 6 F 2 -HS 0 3F mixtures, although other routes are known. 4.3.1. Syntheses involving S20 6F2 or S20 6F2-H S03F. Fluorosulphate derivatives are most commonly synthesised using bisfluorosulphuryl peroxide, S 2 0 6 F2, which itself is prepared in large quantities by the catalytic (AgF2) fluorination of S 0 3 using fluorine[25J (Eqn. 4.2). F2 + 2 SO 3 2 » S 20 6F2 Eqn. 4.2. Bis-fluorosulphuryl peroxide is a thermally stable solid or liquid which may be stored in glass vessels. Its boiling point is 67°C and, at this temperature, reaction times are prohibitively long. For example, A g(S 0 3 F ) 2 is formed by the reaction, at 67°C, between a ten fold excess of S 2 0 6 F 2 and silver powder over seven days P 4^ A mixture of S 2 0 6 F 2 in H S 0 3F is of particular synthetic use as it combines the oxidising power of S 2 0 6 F 2 and the ionising ability of H S 0 3 F. Several distinct advantages arise from the use of this s y s te m :-^ i) Bis-fluorosulphuryl peroxide is completely miscible in fluorosulphonic acid, and a mixture of the two produces a stable, clear, colourless solution. 129 ii) The boiling point of the mixture is 160°C, and this can reduce reaction times and encourage reactions which may not occur with S2 0 6 F 2 alone. iii) Depending on the temperature, it is possible to maintain a high concentration of F 0 2 S0- radicals, formed by the reversible dissociation of S2^6p2- iv) Due to its strong ionising properties fluorosulphonic acid may dissolve products formed at the surface of a reactant, thereby discouraging passivation. The use of S2 0 6 F 2 in HSO 3 F has been employed during the synthesis of a wide range of binary main group and transition metal fluorosulphates. These syntheses usually involve the reaction between excess of S2 0 6 F 2 and an elemental powder. This area of chemistry has been reviewed extensively and a few typical reactions are shown in Equations 4.3-4.5. Au + S 2 0 6 F 2 M + n/2 S 2 0 6 F 2 HS° 3F ► [Au(S0 3 F)3][12] HSQ3F ► Eqn. 4.3. [M (S0 3 F)4] Eqn. 4.4. M = Ti, Zr, Hf, Sn, Pt, Ir, Pd and Ru.[24] HSO3F 2 M + 5 S2 0 6 F 2 ► 2 [M (S0 3 F)5] Eqn. 4.5. M = Nb, Ta and Sb . [ 2 6 1 Metal oxidation using S 2 0 6 F 2 and H S 0 3F in the presence of stoichiometric amounts of M [S 0 3 F] (M = alkali metal) has been used to produce ternary fluorosulphates. Caesium fluorosulphate^27*29^ is generally the 130 alkali metal fluorosulphate of choice in these reactions and there are several reasons for this:- i) Caesium fluorosulphate is formed in situ by the reaction of CsCl and H S 0 3 F.[26] Hydrogen chloride, the by-product, must be removed prior to the addition of the S2 0 6 F2, otherwise oxidation leads to derivatives of the type [C 1 0 J+ (jc = 1 or 2). These result in the formation of a red-orange impurity which may interfere with the reaction. ii) The caesium cation possesses a low reduction potential and is also the weakest electrophile and least polarising alkali metal cation. iii) Generally, caesium salts are highly soluble in many strong- or super-acid systems. iv) Unreacted C s[S 0 3 F] is easily recognised by the occurrence of an infrared absorption at 728 cm"1, assigned to the v 2 sulphur-fluorine stretch; v 2 is found at gradually increasing wavenumbers for the other alkali metal fluorosulphates. This may lead to an unambiguous assignment because these higher frequencies can coincide with coordinated fluorosulphate stretches. 4.3.1.1. Limitations o f the S2O^F2-HSO^F system. Bis-fluorosulphuryl peroxide was first isolated and characterised by Dudley and Cady in 1957.[301 It was prepared by the fluorination of S 0 3 in a flow reactor,*^ and was catalysed by AgF 2 at 100 to 170°C. Since then, a number of alternative routes to S2 0 6 F 2 have been published and these include the low temperature electrolysis of dilute solutions of K [S 0 3 F] in fluorosulphonic acid,^31^ the reaction of CrF 5 and S 0 3 in a 5:1 molar ratioJ32^ the photolysis of C 10S 03F at ambient temperatures for 2-4 ho u rsj33^ the 131 reaction of Cs[AgF4] with S 0 3 J 34^ the low temperature combination of HSO 3 F and 0 2 +[AsF 6 ]V 35] and finally, the reaction of F 0 S 0 2F and S 0 3 J 23^ However, none of these methods have provided a more synthetically viable route to bisfluorosulphuryl peroxide than the original method. Over the years, a number of flaws were found in the original reactor design. Serious concerns by Cady himself led him to publish a waming^36^ about the presence of the potentially explosive by-product fluorine fluorosulphate, F 0 S 0 2 F. Other problems have also come to lightP 1^ These include the S 0 3 delivery system, the copper reactor which proved insufficiently resistant to fluorine at elevated temperatures, the glass-to-metal inlet systems, the use of asbestos and finally the lack of any disposal tower for the highly volatile effluent gases, which included F 2 and F 0 S 0 2F and present serious hazards. Aubke et al. have recently published details of a new but similar catalytic reactor^37! (AgF 2 on copper turnings as a support) for the fluorination of SO 3 in a flow system. The new method eliminates the hazards previously described, but it still requires the building and construction of a large, specialised and expensive reactor. This, in turn, means that S 2 0 6 F 2 is not readily available in bulk quantities and, unless it is produced commercially, then the number of people involved in this area of research will remain small. A further drawback of the S2 0 6 F 2 -H S 03F system is the reaction conditions and times. Table 4.2 summarises some of the reaction conditions involved in the formation of fluorosulphate derivatives. The temperatures used are in the region 25-150°C, while the reaction times vary from half a day to four weeks. This highlights the fact that this route into fluorosulphate derivatives is neither quick nor convenient. 132 Table 4.2. Reaction times and temperatures involved in the formation of fluorosulphate derivatives. Compound Temperature Reaction time /° C /D ays [C d(S0 3 F)2] 90 28 24 [Z n(S0 3 F)2] 90 2 1 24 [Pd(S0 3 F)3] 25-120 3 28 [R u(S0 3 F)3] 60 1 28 [T i(S 0 3 F)4] 25-60 1 0 28 [Z r(S 0 3 F)4] 25-120 2 1 28 [H f(S0 3 F)4] 25-120 2 1 28 [Sn(S0 3 F)4] 25 0.5 28 [Ir(S0 3 F)4] 60-140 6.5 28 [N b(S0 3 F)5] 40 3 15 [T a(S0 3 F)5] 40 3 15 Cs 2 [T i(S 0 3 F)6] 25 5 28 Cs 2 [Sn(S0 3 F)6] 25 0.5 28 Cs 2 [Pt(S0 3 F)6] 80 3 28 Cs 2 [Ir(S0 3 F)6] 150 1 0 28 Cs 2 [G e(S0 3 F)6] 50 2 28 C s[Sb(S0 3 F)6] 70 0.5 29 Reference In order to avoid contamination by other metal fluorosulphates, metal reactors (Monel or nickel) are often replaced by glass vessels. At temperatures above 100°C in these systems several destructive processes can occur. At temperatures around 140°C fluorosulphonic acid dissociates^37^ according to Equation 4.6. This produces HF which itself attacks glass and forms SiF4. Studies by Cady et al. showed that, at 120°C, the thermolysis of S 2 0 6 F 2 133 produces O 2 and S 2 0 5 F 2 ,t38] hence reaction vessels must be vented regularly to avoid explosions. At certain temperatures, the catalytic decomposition of S2 0 6 F 2 to 0 2 and S 2 0 5 F 2 may occurJ37^ This was observed during the formation of Ir(S 0 3 F)4. At 130°C, Ir(S 0 3 F ) 4 decomposes to Ir(S 0 3 F)3, S2 0 5 F 2 and 0 2, the Ir(S 0 3 F ) 3 is then re-oxidised by S 2 0 6 F 2 to form Ir(S 0 3 F)4. H S 0 3F HF + S 0 3 Eqn. 4.6. The reaction of bis-fluorosulphuryl peroxide and H S 0 3F with a variety of metal carbonyls and chlorides has been investigated J 24^ However, with metal carbonyls, the CO is not only substituted but it is also oxidised to C 0 2 whilst, with metal chlorides, oxidation of the chlorine leads to chlorine (I) to (VI) fluorosulphate derivatives, therefore necessitating the use of large excesses ° f s 2 o 6 f 2. 4.3.2. Displacement reactions. In 1967, a route for preparing transition metal fluorosulphates was required and W oolf described the preparation of bivalent fluorosulphate derivatives o f Mn, Fe, Co, Ni, Cu, Zn and C dJ39^ Initially, copper salts were investigated and it was found that copper (II) chloride, sulphate and acetate undergo displacement in boiling fluorosulphonic acid to produce copper bisfluorosulphate. Copper (II) fluoride was found to undergo incomplete substitution, the ease of replacement following the sequence [CH 3 C 0 2]‘ > [S 0 4]2- > Cl" > F \ The divalent fluorosulphates of Mn, Fe, Ni, Cu, Zn and Cd were all formed in an identical manner, and the ease of replacement followed the above sequence. More recently, the solvolysis of FeCl3 J 40^ Z r(C 0 2 CF3)4^41^ and A g (C 0 2 CF3)^42^ in H S 0 3F has been shown to produce F e(S 0 3 F)3, Z r(S 0 3 F ) 4 and A g S 0 3F respectively. Cady et al. noted that the solvolysis of metal chlorides in HSQ3F appeared to be facilitated by the addition of 134 K [S 0 3 F ]J 4 3 1 which may aid the removal of fluorosulphate containing products from the surface of the reactant. The use of S2 0 6 F 2 and S2 0 5 F 2 in displacement reactions does not produce binary fluorosulphates, but leads to heteroleptic complexes such as oxo fluorosulphates ^2 4 1 (Eqn.’s 4.7 - 4.9). TaCl 5 + S 2 0 6 F 2 -> [TaO(S0 3 F)3] Eqn. 4.7. WC16 + S 2 0 6 F 2 [W 0 (S 0 3 F)4] Eqn. 4.8. Ti(OCH 3 ) 4 + S 2 OsF 2 -> Eqn. 4.9. [Ti(OCH 3 )2 (S 0 3 F)2] An alternative approach is the reaction of metal carbonyls with S2 0 6 F 2 which first substitutes and is then oxidised. However, there are only three examples of this type of reaction which appears to be restricted to the first row transition metals (Eqn.’s 4.10 - 4.12)J 2 4 1 Also as explained earlier, the use of S 2 0 6 F 2 and CO is not ideal. [Mn 2 (CO)10] + S2 0 6 F 2 Co 2 (CO ) 8 + S2 0 6 F 2 Cr(CO ) 6 + S2 0 6 F 2 —» —» —^ [M n(S0 3 F)4] [C o(S0 3 F)2] [Cr(S0 3 F)3] Eqn. 4.10. Eqn. 4.11. Eqn. 4.12. The use of these types of displacement reactions appears to be restricted mainly to the first row transition elements. 135 4.3.3. Syntheses involving B rS 0 3F. Bromine fluorosulphate, B rS 0 3 F^43^ has not found many applications within synthetic chemistry, the reasons for this appear to be the excessively long reaction times and the need for S2 0 6 F 2 for its synthesis. Its limited use includes the synthesis of the noble metal fluorosulphates [Pd(S0 3 F)2] [P t(S 0 3 F)4] / 45^ [Ag 3 (S 0 3 F)4] / 46^ and [A u(S0 3 F) 3 ] J 45^ which are prepared by oxidation of the respective noble metal. 4.3.4. Insertion reactions. The insertion of S 0 3 into a metal-fluorine bond has produced fluorosulphate derivatives in a few instances (Eqn.’s 4.13-4.17). The reaction between AgF 2 and S 0 3 produces [A g(S0 3 F)2],[42^ and this is expected when one considers the catalytic nature of AgF 2 during the formation of S2 0 6 F 2 from S 0 3 and F 2 (Section 4.3.1.1. and Eqn. 4.14). The reaction between Ag, BrF 3 and S 0 3 involves the fluorination of silver followed by the insertion of S 0 3 J 47^ whilst the reaction between S 0 3 and CrF5, at -22°C, AgF 2 + S 0 3 Eqn. 4.13. —> [A g(S0 3 F)2] [A g(S 0 3 F)2] + F 2 —> AgF 2 + S2 0 6 F 2 3 Ag + 2 BrF 3 + 3 S 0 3 offers a route to —> 3[ A g(S 0 3 F)2] + Br2 Eqn. 4.14. Eqn. 4.15. W F6 + S 0 3 [WF 2 (S 0 3 F)4] [48] Eqn. 4.16. CrF 5 + S 0 3 —> [C r(S0 3 F)3] + S 2 O^F2 Eqn. 4.17. 136 4.3.5. Oxidising reactions involving H S 0 3F. Fluorosulphonic acid is a mild oxidising reagent Copper and bismuth both react with the boiling acid to produce fluorosulphates derivatives, while Ag, As and Sb all dissolve to give colourless solutions. Niobium, Ta, U and Pb all dissolve to produce green solutions and Na, K, Ca, In, T1 and Sn react to produce green solutions over white precipitates. Woolf et a l established that the white precipitates are fluorosulphates or their decomposition products. They also noted that the ESR behaviour and UV spectra of the green solutions resembled those of sulphur in oleum. This correlates with the fact that the metals which produced the green solutions are good reducing agents in acidic solutions, and hence reduce the sulphur to its elemental form, whereas, the less potent reducing metals can only reduce HSO 3 F to S 0 2. Aubke et a l have recently re-investigated the oxidation of Sb by HSO 3 F / 50! their intention being to prepare the univalent antimony salt, S b (S 0 3 F). An earlier communication had reported the formation and isolation of this sa lt/51! however, no spectroscopic or analytical data has been published. Aubke’s attempt to isolate the antimony (I) fluorosulphate salt failed, and no univalent antimony compounds were identified. However, they were able to structurally characterise some previously unknown antimony (III) fluorofluorosulphate compounds. 4.4. Decomposition of Fluorosulphates. The decomposition of fluorosulphate derivatives is observed to occur via four different pathways, and is outlined as follows:- i) The elimination o f sulphur trioxide. The oxidation of niobium metal using S 2 0 6 F2, in HSO 3 F results in the formation of the unstable pentafluorosulphate, 137 [N b(S 0 3 F)5], this loses S 0 3 and eventually forms the solid [NbF 2 (S 0 3 F)3]^15^ (Eqn. 4.18). N b (S 0 3 F ) 5 ^ N bF(S0 3 F ) 4 + S 0 3 NbF 2 (S 0 3 F ) 3 + S 0 3 Eqn. 4.18. ii) The elimination o f S20 5F2. This results in the formation of an oxofluorosulphate complex. As noted earlier, substitution reactions between metal chlorides or carbonyls and S 2 0 6 F 2 usually produces an oxo-fluorosulphate as the p r o d u c t^ (Eqn.’s 4.19 and 4.20). WC16 + S 2 0 6 F 2 W (CO ) 6 + S 2 0 6 F 2 -» [W 0 (S 0 3 F)4] + S 2 0 5 F 2 ^ [W 0 (S 0 3 F)4] + S2 0 5 F 2 Eqn. 4.19. Eqn. 4.20. Further examples of this type of reactivity include the formation of [V 0 (S 0 3 F)3], [N b 0 (S 0 3 F)3], [M o0 2 (S 0 3 F)2] and [R e0 2 (S 0 3 F)3].[24] The elimination of S2 0 5 F 2 does not always result exclusively in the formation of an oxo-fluorosulphate complex. For example, the decomposition of Ir(S 0 3 F ) 4 yields S 2 0 5 F2, but Ir(S 0 3 F ) 3 and 0 2 are also produced.^37! iii) The reductive elimination o f SO3F radicals. The thermal decomposition of Pdn [PdIV(S 0 3 F)6] leads to the elimination of S 2 0 6 F 2 (Eqn. 4.21).[52] This is not common and highlights the strong oxidising ability of the palladium fluorosulphate complex. The addition of bromine to the mixture accelerates the rate of decomposition which produces Pd(S 0 3 F ) 2 and B rS 0 3 F. A similar process was observed for A g(S 0 3 F ) 2 which decomposes at 210°C to produce A g (S 0 3 F) and S2 0 6 F2.[53] o Pdu [PdIV(S 0 3 F)6] 2 Pd(S 0 3 F ) 2 + S 2 OeF 2 138 Eqn. 4.21. iv) The decomposition of alkaline and alkaline earth metal fluorosulphate salts, follows one or both of the paths outlined in Equations 4.22 and 4.23. M (S 0 3 F)x —) 2 M (S 0 3 F)a; MFjp + x SO 3 —^ Eqn. 4.22. + x SO 2 F 2 Eqn. 4.23. The actual decomposition generally reflects small differences in the structure, for example the co-ordination number for Ca(II) and Ba(II) is six and eight respectively. However, the polarising power of Ba(II) is lower than that of Ca(II). Consequently, C a(S 0 3 F ) 2 is pyrolysed to CaF 2 whereas B a(S0 3 F ) 2 produces B a(S 04). These types of decomposition reactions have been studied extensively by Muetteries et alS54,55^ 4.5. Spectroscopic Characterisation of Fluorosulphate Compounds. 4.5.1. Vibrational spectroscopy. Vibrational spectroscopy is a powerful tool for the identification of fluorosulphate derivatives. The most useful information is found in the region 1500 - 700 cm '1, and reveals the type of interaction between the cation and anion, as well as the mode of coordination of the fluorosulphate ion (Figure 4.3). 139 Figure 4.3. The bonding modes of the fluorosulphate ligand. Free ion Q ° O-Monodentate 0,0'-B identate M '— . _ \ / > 0 ^ 0 o/S\ C 3va t S 0 , 0 ’,0"-Tridentate ___ __ / 0 ^A , Cs a t S Cs at S C3vat S For potassium fluorosulphate, K [S 0 3 F], the anion possesses C3v symmetry and this leads to six vibrational modes which are both infrared and Raman active (Table 4.3) three of these are E modes and are doubly degenerate. Table 4.3. Infrared vibrational d a t a ^ and assignments for potassium fluorosulphate. Infrared / cm - 1 Assignment 1280 s v 4 (E) S03 asym str 1080 s Vl (A) S°3symstr 750 s v2 (A) SFstr 590 s ^ 5 570 m ^3 (A) S03Symdefn 480 m V6 (E) SF defn (E) S03 asym defn For C s[S 0 3 F] the anion is d i s t o r t e d t h i s lowers its symmetry from C3v to Cs and results in a splitting of the E modes of 11-24 cm-1. This arises from crystal packing and the molecular structure has shown the S -0 bonds to differ in length (Table 4.8). 140 For tridentate fluorosulphate anions, as was shown in Figure 4.3, the symmetry of the anion remains C3v. Cobalt (II) fluorosulphate gives an infrared spectrum which exhibits six absorptions, and these correspond to the six fundamental modes of a fluorosulphate group possessing C3v symmetry (Table 4.4); no splitting of the degenerate modes was observed Table 4.4. Infrared vibrational d a t a ^ and assignments for [Co(S0 3 F)2], [Fe(S0 3 F)2] and [Ni(SQ 3 F)2]. Infrared spectra / cm " 1 [C o(S0 3 F)2] [Fe(S0 3 F)2] [N i(S0 3 F)2] Assignment 1265 vs 1261 vs 1262 vs v 4 (E) 1109 s 1118s 1 1 2 0 850 s 862 s 859 s v 2 (A) 610 s 610 s 619 s v5 (E) 568 m 568 m 568 m v 3 (A) 420 m 419(m 422 m v6 (E) s Vi (A) For the ionic salts of fluorosulphate, particularly the large alkali metals, the cation-anion interaction is small. However, the shift of the v 2 frequency by ~100 cm ' 1 implies some cation-anion interaction. Therefore, the maintenance of C3v symmetry and the significant anion-cation interaction indicates that all three oxygens are attached in an equivalent manner to the cobalt ions. The same explanation applies for F e(S 0 3 F ) 2 and N i(S 0 3 F ) 2 and, in view of their lack of volatility and insolubility in fluorosulphonic acid, it is apparent that these compounds are polymeric. 141 Table 4.5. The Raman vibrational data^1^ and assignments for K [B r(S0 3 F)4] and K [I(S0 3 F)4]. K[S03F] Mode K[Br(S03F)4] K[I(S03F)4] Mode 1287 s v4 (E) 1424 m 1409 m v7 (A") S03 as 1082 s Vj (A) 1407 w - 786 s v2 (A) 1237s 592 s v5 (E) 570 m V3 (A) 970 m 409 m V6 (E) 834 m 837 m - - 578 ms 582 ms V5 - - 553 ms 554 ms v8 (A") - - 615 vs 620 vs V3 - - 406 w 407 w V9 (A") - - 239 vs 239 vs V6 (A') 1 2 2 0 w sh - 1250 s 1 2 2 2 Vi (A1) S03 w sh 1 0 0 2 m - v4 (A') S03as v2 (A') SF (A') (A’) The fluorosulphate group can act as a monodentate ligand and to demonstrate, this the Raman spectra of K [B r(S0 3 F)4] and K [I(S 0 3 F)4] are presented in Table 4.5 A number of differences are apparent when the vibrational data are compared to that of K [S 0 3 F ]. The S-F stretching mode is observed at 834 cm '1, which is higher than that observed for K [S 0 3 F] and indicates a covalent interaction between the fluorosulphate group and the halogen centre. There are nine absorptions and this indicates a lowering of symmetry which implies either a mono- or bi-dentate interaction. The magnitude of the splitting of the E modes is proportional to the degree of covalency, [Br(S 0 3 F)4]- 454 cm " 1 [v4 (E) v 4 (A') + v 7 (A")] and [I(S 0 3 F)4]' 407 c m '1: the smaller splitting of the latter indicates a slightly more polar bond. For anions of this type the vibrational trends are interpreted as being due to a 142 covalent monodentate interaction between the fluorosulphate group and the halogen centre. Covalent bridging fluorosulphate groups are found in [Fe(S0 3 F)3] [Sn(S0 3 F)2 Me2] ^ and [Sn(S0 3 F) 2 Cl2 ] J 56^ the tin compound having been crystallographically characterised. As a result of the 0 ,0 -b id en tate interaction, the symmetry is lowered to Cs and the degeneracy is removed. The infrared spectra of the iron and tin fluorosulphates indicates the presence of only one type of fluorosulphate group (Table 4.6). The S-F stretching mode occurs in the same region as that observed for K [B r(S0 3 F)4], this is indicative of a covalent interaction between the fluorosulphate group and the metal centre. It is the two higher S-O stretching modes, found between 1355-1385 and 1130-1180 cm '1, which identify the bidentate ligand. These two values are found to be intermediate between those of K [S 0 3 F] and K [B r(S0 3 F)4], whereas the position of the third S-O mode remains virtually the same as that for K [S 0 3 F]. Table 4.6. Infrared vibrational data, cm '1, for the -S 0 3F group in [Fe(S0 3 F)3], [Sn(S0 3 F) 2 Me2], [Sn(S0 3 F)2 Cl2], K [B r(S0 3 F)4] and K [S 0 3 F]. [Fe(S03F)3] [Sn(S03F)2Me2] [Sn(S03F)2Cl2] K[Br(S03F)4] K [S03F] 1360 m 1350 m 1385 s 1416 mw 1285 s 1137 s 1180 s 1130 vs 1229 mw - 1090 s 1072 s 1087 s br 970 ms 1079 s 850 m 827 s 864 s 834 m 745 s 630 m 620 m 628 m 615 vs - 579 w 590 s 586 s 578 ms 590 s 551 m 554 m 555 s 553 w 570 m 419 w 417 w 420 w sh 406 w 407 m 143 Bridging and terminal fluorosulphate groups may be present in the same molecule. This [Sn(S 0 3 F)4] J is the case for [Br(S0 3 F)3],[13) [I(S 0 3 F)3] [13] and ' The infrared spectra are very complex and in the case of B r(S 0 3 F ) 3 a total of 27 bands and shoulders are observed between 1500-200 cm"1. Considering just the v(S-O) region, 1500-900 cm '1, one set of absorptions are in approximately the same place as those in [Br(S0 3 F)4]', and the second set are assigned to bridging fluorosulphate groups. The presence of bridging fluorosulphate groups leads to low volatility, high decomposition temperatures and a reluctance to dissolve in H S 0 3 F, as is the case for [Sn(S0 3 F)4] and B r(S 0 3 F)3. Finally, weakly coordinating highly ionic fluorosulphate groups have been observed for [M (S0 3 F)2 (C 0 )2][57] (M = Pt or Pd), [M (S0 3 F)(C 0)5] [19] (M = Mn or Re) and [Sn(S0 3 F)2] For example, the vibrational data for [R e(S0 3 F)(C 0)5] ^ 1 9 1 is consistent with a weakly coordinating monodentate fluorosulphate group. Several important features highlight this (Table 4.7). i) In the S-O and S-F stretching region, the S-F stretch and the symmetric S-O stretch for ionic fluorosulphates are found in nearly identical regions (e.g. K [S 0 3 F] v(S-F) = 745 and v(S-O) = 1079 cm '1), ii) The splitting of the asymmetric S 0 3 stretch, which is indicative of a departure from C3v symmetry, is only slight (-6 0 cm " 1 for Re) and this splitting is too large to be solely due to site affects (for example in [N O ][S0 3 F] site effects result in a splitting of 1020 cm_1 )J58^ iii) The position of v(SO—M), 1030 cm"1, is consistent with the highly ionic character proposed for the M -O S02F interaction, iv) The position of v(CO), 2160-1980 cm"1, is also indicative of the highly ionic nature of -S 0 3F (for example see rhenium pentacarbonyl seflate, Table 2.14). The spectral features described here differ markedly from the patterns displayed by other covalent monodentate fluorosulphate groups (e.g. K [B r(S0 3 F)4], Table 4.5). 144 Table 4.7. Comparison and assignment of the infrared vibrational data, cm '1, for K [S 0 3 F] and [R e(S0 3 F)(CO)5]. k s o 3f Mode [Re(S0 3 F)(C0)2] Assignment - - 2160 w sh v(CO) - - 2141 vs v(CO) - - 1980 vs v(CO) 1280 s v4 (E) 1315 m V a sy m (S 0 2) - - 1255 m V a sy m (S 0 3) 1080 s V , 1170 w V Sy m ( S 0 3) - - 1 1 2 0 w V sym (S O s ) - - 1030 m V(S 786 s v2 (A) 760 m v (S-F) (A) O --M ) 4.5.2. X-ray crystallography. The polymeric nature of many transition metal fluorosulphates has prevented their crystallographic analysis, although recently compounds such as C s[A u(S0 3 F)4],[29) Cs2 [Pt(S0 3 F)6],|29) [Sn(S0 3 F)2 Me2][59] and C s[Sb(S0 3 F) 6 ]t29l have been characterised. The number of molecular structures being reported is increasing and, as will be shown in Section 4.6, this allows comparisons to be made of the different bonding situations. Bond lengths vary considerably depending on the nature of the interaction talcing place; typical values are, sulphur-oxygen (bridging) -1.41-1.51 -1.37-1.45 A and sulphur-fluorine -1.45-1.57 A. 145 A, sulphur-oxygen (terminal) 4.5.3. Fluorine-19 NMR spectroscopy. The resonances generated by a fluorine bonded to the sulphur of a fluorosulphate group appears to be of little use as a diagnostic tool. Although shifts have been observed, they are usually small and the lack of multiplicity offers no additional information. Very few fluorine shifts appear in the literature and those reported are shifted little from that of fluorosulphonic acid. The largest shifts (5compiex-§Hso3F) are undoubtedly those which arise from covalently bound fluorosulphate groups. Hohorst and Shreeve conducted 19F NMR studies on a range of fluorosulphate derivatives The majority of shifts were observed in the range -40 to -50 ppm, however, they were unable to relate the observed shifts to any single factor. 4.5.4. Mdssbauer spectroscopy. Mossbauer spectroscopy is a useful structural tool which can be used to define oxidation states and local symmetry, however, it is limited to a few nuclei such as 57Fe and 119 Sn. Mossbauer spectroscopy has successfully been used to define the geometry in a number of tin (II) and (IV) fluorosulphate compounds . 1 5 6 ’6 1 ' 6 3 1 4.5.5. Magnetic studies and electronic spectroscopy. The difficulty in obtaining single crystals and the limitations of other spectroscopic tools has meant that the characterisation of new complexes has depended heavily on vibrational spectroscopy. Magnetic susceptibility measurements and electronic spectroscopy have, consequently, played an important role in the understanding of the nature of metal fluorosulphate derivatives J 64,65^ The use of electronic spectroscopy and magnetic susceptibility enabled chemists to understand more about the bonding occurring 146 at the metal centre. This information, coupled with the vibrational data often permitted the coordination within a fluorosulphate derivative to be unequivocally assigned. 4.6. Single Crystal X-Ray Analysis of Fluorosulphate Compounds. Caesium fluorosulphate has recently been crystallography c h a r a c t e r i s e d T a b l e 4.8 lists the internal bond lengths and angles for the molecule. This can be considered to represent the fluorosulphate group in a totally ionic environment. It must be recognised that contacts to the very weak electrophile Cs+ will be rather long and weak. As more electrophilic cations are encountered, so these interatomic contacts are expected to shorten and to increase in strength. Table 4.8. Bond lengths and angles for C s[S 0 3 F]. Bond Length (A) Bond Angle (deg) S-O (1) 1.458(2) F(l)-S-0(1) 102.3(1) S-O (2) 1.437(2) F(l)-S-0(2) 106.8(1) S-O (3) 1.436(2) F(l)-S-0(3) 107.8(2) S-F 1.569(2) 0 (l)-S -0 (2 ) 113.6(1) 0 (l)-S -0 (3 ) 113.2(1) 0(2)-S-0(3) 112.7(1) The fluorosulphate anion departs slightly from C3v symmetry towards the point group Cs. The cation is nine coordinate with eight sites occupied by 147 oxygen and one by fluorine. The above bond parameters are comparable with those obtained for K [S 0 3 F] and [NH4 ][S 0 3 F]. The addition of fluorosulphonic acid to C s[S 0 3 F] results in the formation of a monosolvate of the composition Cs[H (S0 3 F)2]. The presence of a rather short O—H—O hydrogen bond was confirmed by X-ray studies, and the hydrogen atom was located. The hydrogen atom was found at the inversion centre and the O—H—O bond was linear and symmetrical. The 0--H bond for [H(S 0 3 F)2]- is 1 .2 1 0 (2 ) A, which is rather short when compared to [H(OTeF 5 )2 r , 1.297(8) A, is still considerably longer than that for [HF2]' which contains the strongest and shortest hydrogen bond, 1.13(1) A. It was noted by Aubke et a l that the formation of a hydrogen bond resulted in several changes The S-O bonds involved in hydrogen bonding were lengthened to 1.471(2) A whereas the remaining S-O, 1.399(3) and 1.406(2) A, and S-F bonds, 1.531(2) A, were shortened relative to Cs[S 0 3 F]. It appears that bonding to the peripheral oxygen and fluorine atoms in the pair [S 0 3 F ]' and [H (S 0 3 F)2]" strengthens for the binuclear species at the expense of the bond strength in the bridging region. The coordination number of the caesium in this species is 1 2 , and it appears that the contacts are slightly longer than those for C s[S 0 3 F]. Overall, the bonds involving peripheral atoms strengthen slightly from C s[S 0 3 F] to Cs[H (S0 3 F)2], therefore, the basicity of the peripheral atoms decreases. Hence, the ability of these atoms to coordinate to Cs+ is reduced and the fluorosulphate ions appear to be more nucleophilic. The trends described above are demonstrated by the molecules C s[A u(S0 3 F)4] and C s[Sb(S0 3 F)6] (Table 4.9 and Figure 4.4) where the coordination geometries of the central atoms are square planar and octahedral respectively. The strength of the cation-fluorosulphate interaction is reflected by the strong bonds between the central cation and the oxygen bridging atoms, Ob. Due to variations in the atomic number and oxidation state, the M-Ob bond distances are not strictly comparable. Strong coordination of the fluorosulphate groups, via oxygen, to the central ion has two, indirect, secondary effects:-^29^ i) 148 The bond between sulphur and the bridging oxygen will lengthen, ii) Due to increased multiple bonding in the approximately tetrahedral fluorosulphate groups, the bonds between sulphur and the peripheral oxygen and fluorine atoms will shorten relative to C s[S 03F]. The expected trends may be slightly modified by inter atomic contacts to the cation Cs+. Table 4.9. Bond lengths and angles for Cs[A u(S03F)4] and Cs[Sb(S03F)6]. C s[Au(S03F)4] Cs[Sb(S03F)6] A d(S- Ob) / A 1.968(4) 1.955(2) 1.508(4) 1.516(2) Z(M O bS) 125.2(3) 136.6(1) 1.393(5) 1.396(3) 1.402(6) 1.409(4) 1.523(6) 1.486(3) Parameter </(M-Ob) / d{ S-Ot) / A d( S -F ) / A Ob = bridging oxygen, Ot = terminal oxygen Figure 4.4. Molecular structures1^ of a) [Au(S03F)4]" and b) [Sb(S03F)6] ' . 149 The anion [A u(S0 3 F)4]" possesses Q symmetry and the Au-O bond distances average 1.972(4) in this molecule is 1 0 A. The coordination number of the caesium cation . Within C s[Sb(S0 3 F)6] the six symmetry-related fluorosulphate groups are octahedrally coordinated around the antimony atom. The Sb-O distance, 1.955(2) A, is of the same order of magnitude as the A u-0 distance in [A u(S0 3 F)4] \ The caesium ion is 12 coordinate and the weak inter-atomic contacts involve only peripheral oxygen atoms. These very weak contacts indicate that [Sb(S0 3 F)6]‘ is a very poor nucleophile. From the work done by Aubke et a l a number of interesting conclusions were r e a c h e d :- ^ i) The peripheral atoms of the molecular anion are the most likely to coordinate to a cation, ii) The fluorine atom is least likely to coordinate to a cation, presumably this is due to the higher electronegativity of the fluorine over oxygen and therefore, its lower basicity, iii) These weak inter-atomic contacts may cause distortions from the idealised geometries, iv) These inter ionic interactions may affect and probably slightly weaken the internal bonds of sulphur to the peripheral atoms. For both anions, the S-Ot and S-F bond lengths are considerably reduced when compared to those of C s[S 0 3 F]. This suggests strong multiple bonds and is most striking for [Sb(S0 3 F)6]". The 'onion skin' model was suggested by Aubke et alS29^ to illustrate the low basicity of this anion. The inner coordination sphere consists of six octahedrally arranged Ob atoms, which are strongly bonded to the antimony. There is a rather wide Sb-Ob-S angle of 136.6 (1)° and finally a third sphere containing 18 hard donor atoms. The peripheral atoms (oxygen and fluorine) are even more strongly bonded to the sulphur than are the oxygen atoms of the inner sphere, resulting in a very weak nucleophile. A final feature noted for monodentate covalent fluorosulphate groups is the increase in the M-Ob-S bond angle with increased coordination number of M. The angle increases in the following order 117.2(2)° for [H (S0 3 F)2]" < 125.2(3)° for [A u(S0 3 F)4]‘ < 136.6(1)° for [Sb(S0 3 F)6]-. Steric effects are not 150 considered to be the reason for the increasing angle. Instead, an increase in the M-O bond strength causes a widening of the angle which may indicate delocalisation of lone pair electron density from the bridging oxygen to M: for M = H the most acute angle is observed and angle widening increases with increasing oxidation state of the metal. As already shown, the fluorosulphate group can act as a bidentate ligand. The molecular structures of [c-Pd2(|i-C 0)2][(S 03F)2]t17l and [Sn(S03F)2(CH3)2]t59] are closely related. Two oxygen atoms of the fluorosulphate group are weakly coordinated to the metal centres, the third is not involved in direct coordination to either the Pd or the Sn. The metal-oxygen bonds are long and weak. As expected, the Sn-O distance is longer than the PdO distance, where Sn is noted to have a larger covalent radius. The three sulphur-oxygen bond distances are not significantly different. Overall, although the fluorosulphate group is behaving as a bidentate ligand, the departure from C3v symmetry is only slight, and in both cases the S-O and S-F bond distances are virtually the same as for C s[S 03F]. Figure 4.5. Molecular structure of [Au(S03F)3]. The molecular structure of [A u(S0 3 F)3][12] shows that it is dimeric in the solid-state and possesses both mono- and bi-dentate fluorosulphate groups (Figure 4.5). The presence of bidentate, symmetrically bridging fluorosulphate groups generates an eight-membered centrosymmetric ring which adopts a chair conformation, the two gold atoms being in transannular positions and linked by two S 0 2 moieties. Both gold centres are identical and the geometry around each gold is virtually square planar. The covalent monodentate ligands show bond distances in accord with those already summarised. The bidentate fluorosulphate ligands are not as strongly bonded to the gold centre as the monodentate ligands: Au-Oav = 1.957(8) Au-Oav = 2.018(7) A A for the mono and for the bidentate. Within the bidentate fluorosulphate group the expected trends arise. The sulphur bridging oxygen bonds have increased in length by ~ 0.3 A. More noticeably the S-F and S-Ot bond distances for both the mono- and bi-dentate fluorosulphate groups are considerably shorter than those found in C s[S 0 3 F]. Presumably, the same explanation applies as in the structures described above, where an increased bond strength was observed for the peripheral atoms. No tridentate fluorosulphate complexes have been crystallographically characterised to date. For weakly co-ordinating, ionic fluorosulphate groups, one can assume that none of the internal parameters of the fluorosulphate group will be significantly changed. Covalent tridentate fluorosulphate groups will vary depending on the strength of the interaction between the bridging oxygen and the cationic centre. As the interaction increases so the S-Ob bond would be increased and the S-F bond will presumably become shorter and less basic. These compounds will be extensively polymeric, and recrystallisation from a suitable solvent would be expected to be a major problem. 152 4.7. Recent developments in fluorosulphate chemistry. 4.7.1 Cationic carbonyl metal species. Very recently a number of noble metal carbonyl cations, which are stabilised by fluorosulphate ligands,^66^ have been reported. Little is known about why some of these compounds have relatively high thermal stabilities despite the absence, or near absence, of metal-carbon 7t back-bonding. The strength of a CO bond is readily determined by vibrational s p e c t r o s c o p y T h e carbon-oxygen stretching frequency is very sensitive to changes in the CO bond order, caused by n back-donation from the metal into C-O n* antibonding orbitals. Shifts to lower frequency, that observed for gaseous carbon monoxide is 2143 cm-1, are used not only to detect and estimate the extent of synergic bonding, but also to assign co-ordination modes of the carbonyl ligand. Terminal monodentate CO groups are usually found in the region 2125-1850 cm '1, while bridging bidentate CO groups have u(CO) values between 1860-1700 c m '1. Another group of carbonyl derivatives exists, in which n back-bonding seems to be insignificant in the formation of a metal-carbon bond. The best known examples include metal carbonyl cations and metal carbonyl halides. Metal carbonyl cations have only been discovered recently whereas metal carbonyl halides have a much longer history which stretches back to 1868 when Schtitzenberger discovered the three platinum (II) carbonyl halides [Pt Cl2 (CO)2], [Pt2 Cl4 (CO)2] and [Pt2 Cl4 (CO ) 3 ] . 1 6 8 1 The study of noble metal carbonyl cations began with the isolation of carbonyl gold (I) fluorosulphate, [A u(S0 3 F)(C0)] J69^ This was serendipitous and stemmed from investigations to detect the formyl cation, [HCO]+. Attempts to observe the cation by NMR spectroscopy using 13 C-enriched CO in a superacid solution were unsuccessful. This was presumably due to proton 153 exchange which occurred even at low temperatures. It appears that the stretching force constant of t)(CO) for [HCO]+ may represent the upper limit for a complex ion with a solely a bonded CO: complete absence of n backdonation. The value obtained may have acted as a benchmark for judging the extent of n back-donation in other complexes. Gold trisfluorosulphate was first reported in 1972 by Johnson, Dev and Cady and it was soon realised that A u(S 0 3 F ) 3 should act as a fluorosulphate ion acceptor J 7 0 1 Later, it was shown that A u(S 0 3 F ) 3 does indeed behave as a Lewis acid in fluorosulphonic acid to form the conjugate base [A u(S0 3 F)4 ]'J 7 1 1 In 1990 an attempt to protonate carbon monoxide using the superacid system H S 0 3 F -A u(S0 3 F ) 3 resulted in the discovery of the metal carbonyl cation [Au(CO) 2 ]+J 6 9 1 Gaseous carbon monoxide is found to be virtually insoluble in fluorosulphonic acid, therefore the uptake of the gas could be easily monitored. A colour change was observed for the reaction and the volatile reduction products, C 0 2 and S2 0 5 F2, were isolated. W ork-up yields a white moisture-sensitive solid which is stable up to 190°C and has a melting point of 49-50°C. Characterisation of the solid, using vibrational spectroscopy, identified it as [A u(S0 3 F)(C0)] and the reported v(CO) is 2195 cm-1, well above that for free CO (2143 cm '1). It was suggested that the reductive carbonylation of A u(S 0 3 F ) 3 follows Equation 4.24. The resultant Au (I) species is then stabilised by the complexation of CO according to Equation 4.25, the final product [A u(S 0 3 F)(C 0)] results from a very facile substitution reaction as summarised in Equation 4.26. A u (S 0 3 F ) 3 + CO —> Au+ + C 0 2 + [S 0 3 F] + S2 O^F5 Eqn. 4.24. Au+ + 2 CO [Au(CO)2]+ Eqn. 4.25. -> 154 [Au(CO)2]+ + [SO 3F]- -> [A u(S0 3F)(C0)] + CO Eqn. 4.26. The need for a strong protonic acid during the synthesis of transition metal carbonyl derivatives is apparent. The extension of this to other systems such as H S 0 3 F-[P t(S 0 3 F)4] ^ was soon undertaken,^13,68^ and led to the complexes [cis-P t(S 0 3 F) 2 (C 0 )2] and [cw-Pd(S0 3 F) 2 (C 0 )2].[73] The latter compound was structurally characterised by X-ray diffraction^74^ and the molecular structure showed a square planar geometry at the Pd centre, with terminally bound CO and monodentate fluorosulphate groups in a cis arrangement. The absence of significant Pd to CO 7 t-back donation is highlighted by the high CO-stretching frequencies, v av(CO) of 2218 cm-1. The CO bond lengths are also short, 1.106(6) and 1.114(6) those of gaseous carbon monoxide (cf. 1.12822 A). A, when compared to The X-ray study also revealed a number of intra- and inter-molecular contacts between the carbon atom of the CO group and the oxygens of the fluorosulphate groups, these appear to stabilise the structure. This synthetic approach has been extended to iridium J37^ As was described earlier, Ir(S 0 3 F ) 3 is obtained from the thermal decomposition of Ir(S 0 3 F)4, itself formed during the oxidation of iridium metal by S2 0 6 F 2 in H S 0 3 F. A s with the previous examples, mer-[Ir(S0 3 F)3 (CO)3] forms from the binary fluorosulphate precursor in fluorosulphonic acid, at 60°C for four days, under two atmospheres of carbon monoxide (Eqn. 4.27 and Figure 4.6). Ir(S 0 3 F ) 3 + 3 CO -> raer-[Ir(S0 3 F)3 (C 0)3] Eqn. 4.27. It should be noted that no change in oxidation state has occurred, therefore, this is not a reductive carbonylation reaction as observed for the previous noble metal carbonyl cation derivatives. These carbonylation reactions are usually very fast. However, the formation of mer-[Ir(S0 3 F) 3 (C 0 )3] proceeds over four days. Aubke et al. suggested a gradual stepwise addition of 155 CO, rather than the substitution of CO by [S 0 3 F]' as was previously observed P 7^ Figure 4.6. Crystal structure of mer-[Ir(S 0 3 F) 3 (C 0 )3 ] . Again, as was observed for cw-[Pd(S0 3 F)2 (C 0)2], the CO stretching frequencies (2249, 2208 and 2198 cm-1) suggest significantly reduced n back bonding. It also appears that significant inter- and intra-molecular SO-CO contacts exert a stabilising influence on the structure of mer- [Ir(S0 3 F)3 (C 0)3] . The isolation of mer- [Ir(S0 3 F) 3 (C 0 )3] has expanded the range of cationic metal carbonyl fluorosulphates with significantly reduced n back bonding from Groups 10 and 11 into Group 9. Reductive carbonylation reactions of fluorosulphate derivatives in fluorosulphonic acid has led to similar reactions being carried out in SbF5 under mild conditions and an atmosphere of carbon monoxide. The generated homoleptic cations are stabilised in all cases by [Sb2 F n ]' anions, and again secondary contacts appear to stabilise the resulting salts in the solid phase. The cations [Ru(CO)6]2+ and [Os(CO)6]2+ were isolated^75! by the reductive carbonylation of [R u(S0 3 F)3] and [O s(S0 3 F)3] in SbF 5 under a CO atmosphere. A number of other species have been isolated (c f [Fe(CO)6]2+, 156 [Ir(CO)6]2+, [Hg(CO)2]2+, [Pt(CO)4]2+ and Pd(CO ) 4 ] 2 + ) . 1 7 3 ' 7 6 - 7 8 1 The latter two species are generated by carbonylation of the complexes [d s-P t(S 0 3 F) 2 (C 0 )2] and [c/.s-Pd(S0 3 F) 2 (C 0 )2], in SbF 5 under a carbon monoxide atmosphere. The ability of fluorosulphate ligands to form inter- and intra-molecular contacts appears to be a significant factor in stabilising these types of species. As the expansion of this area continues the role of the fluorosulphate anion, and similar anions (e.g. [Sb 2 F n ]~) should become clearer. 4.7.2. Super acids. Anhydrous hydrogen fluoride and fluorosulphonic acid are the strongest known Brpnsted acids. They each have identical Hammett acidity functions, -H0, of 151, and are considerably more acidic than conventional aqueous acid systems such as nitric or sulphuric acid (i.e. 106 - 10 1 0 times). They are generally referred to as super acids and are essentially non aqueous, this is important since the strongest acid which can exist in the presence of H 2 0 , is [H 3 0 ] +. The Hammett acidity function^7,67! is used as a scale by which the acidity of super acid systems can be gauged, and is defined in Equation 4.28. H 0 = p t f B H + - l o g { [BH+] / [B]} Eqn. 4.28. B is an indicator base and [BH+] is its protonated form, pKBH+is equal to -logK where K is the dissociation constant of [BH+], the ratio of [BH+] to [B] may be determined spectrophotometrically. Early work suggested that anhydrous HF had a value of -H 0 of 11, however, recent studies have indicated that this value is too low . t 7 9 1 It appears that as the acids approach 1 0 0 % purity and become anhydrous there is a rapid increase in -H0. This is due to a rapidly increasing concentration of the solvated proton. Autoprotolysis of HF and H S 0 3F results in the formation of the following species, [H 2 F]+ and [H2 S 0 3 F]+. Due to the problems involved in 157 obtaining 100% anhydrous HF the presence of small concentrations of basic impurities, such as water, drastically decreases the value of -H0. The use of fluorosulphonic acid as opposed to anhydrous HF is far more desirable, this is because fluorosulphonic acid has a wider liquid range, and as stated earlier, specialised equipment is not required to handle it. These acids are employed in industrial processes and academic research.[7,66] As reaction media they have been used to generate a wide range of highly reactive organic and inorganic cations. In solution, they stabilise these otherwise short-lived species by virtue of their high acidities and the low nucleophilicity of the conjugate base ion. The high acidity of fluorosulphonic acid is due to the formation of the [H 2 S 0 3 F]+ (Eqn. 4.29). 2 H S 0 3F [H2 S 0 3 F]+ + [S 0 3 F]' Eqn. 4.29. Amongst the most powerful super acids the highest acidities are observed in conjugate superacid system s,^ which usually consist of a strong protonic acid and a powerful Lewis acid. The conjugate superacid system, H S 0 3 F-SbF5, is often termed ’magic acid'. This system is quite complex, owing to facile fluoride versus fluorosulphate exchange and the presence of concentration dependent solute association via - 0 S 0 2 F- and -F- bridgesJ28^ These factors give rise to the presence of several anions in solution and, as a result, the use of this system during the synthesis of salts with electrophilic cations is a problem. This difficulty can be avoided by using a conjugate super acid system of the type H S 0 3 F -E (S 0 3 F)n, where E (S 0 3 F)n is a high oxidation state binary fluorosulphate acting as a Lewis acid (Eqn. 4.30) typically, E = Au (n = 3), Pt (n = 4), Nb or Ta (n = 5). These systems all have inherent problems which include the high price and oxidising ability of Au(III) and Pt(IV) and the limited thermal stability of the Ta and Nb systems. 158 2 m HSO 3 F + E (S 0 3F)n m [H 2 S 0 3 F]+ + [E (S0 3F)n+m]m- Eqn. 4.30. The usefulness of these superacid media is determined by three general properties which are demonstrated by the system H S 0 3 F-A u(S0 3 F)3:- 1) The proton donor strength should be very high and is limited by the acidity of [H 2 S 0 3 F]+. 2) The nucleophilicity or electron pair donor ability of the conjugate base ion, [A u(S0 3 F)4] \ should be very low. 3) The electron pair acceptor strength or Lewis acidity of the molecular Lewis acid should be high. Some current research is directed towards finding alternative conjugate superacid systems, and this involves the synthesis of a range of binary and ternary fluorosulphates. Such an investigation led to the isolation of [M (S 0 3 F)4] and Cs 2 [M (S 0 3 F)6] species, where M = Ti, Zr or Hf However, all three Group four binary fluorosulphates are insoluble in H S 0 3 F. The acceptor ability of these compounds is demonstrated by the isolation of the ternary fluorosulphates, which are thermally stable up to 260°C. The [M (S 0 3 F)4] systems are polymeric and demonstrate the intrinsic acceptor ability of the metal centre. Only where there is a limited tendency towards polymer formation, e.g. [A u(S0 3 F)3] which is dimeric, is it feasible to use binary fluorosulphates in conjugate H S 0 3F super acid systems. 159 4.8. Area of Study. The use of fluorosulphonic acid as an oxidising agent is an area of chemistry which is little explored. On the other hand, the use of bisfluorosulphuryl peroxide as an oxidative-addition reagent is well established, however, the difficulty involved in its preparation has restricted its use to a few institutions. Brazier and W oolf and Aubke et alS50^ have demonstrated that fluorosulphonic acid possesses some oxidising powers. Whether this oxidising prowess is due to the presence of sulphur trioxide is unknown, S 0 3 will undoubtedly be present in small quantities as an equilibrium product. Initial experiments were carried out in this laboratory to establish the similarities between AHF and H S 0 3 F. This involved the protonation of the carbonyl clusters [Ir4 (CO)12], [Ru 3 (CO)12] and [Os3 (CO)12]. In the present work it was envisaged that the presence of fluorosulphate anions might facilitate the formation of crystals and allow definitive characterisation of these protonated carbonyl clusters. Further reactions were carried out on a variety of metal carbonyl complexes and Group 4 cyclopentadienyl derivatives. Here, it was hoped that the presence of the carbonyl groups would facilitate the oxidation of these species, and that, hopefully, this would establish new synthetic routes to transition metal fluorosulphate complexes. 160 4.9. The Reactions of [Ir4 (CO)i2], [Ru3 (CO)i2] and [Os3 (CO)i2] with H S 0 3 F. The protonation of the carbonyl clusters [Ir4 (CO)12], [Ru3 (CO)12] and [0 8 3 (0 0 )!2] has been previously reported. Early studies by Knight et. al. employed the use of concentrated H 2 S0 4,^8°] whereas more recent work by Hope et al. used A H FJ81^ Anhydrous HF is a convenient solvent for the fluorination of transitionmetal carbonyls. It was observed by NMR spectroscopy that protonated transition-metal carbonyl complexes were present in solution, and this suggested that HF is not an inert solvent in these reactions. The carbonyl clusters [Ru 3 (CO)12], [Os3 (CO)12] and [Ir4 (CO)12] react with AHF to produce [Ru 3 (C O ) 1 2 H ]+, [Ru(CO)5 H]+, [Os3 (CO) 1 2 H]+, [Os(CO)5 H]+ and [Ir4 (CO) 1 2 H2]2+ respectively, and were characterised in solution by a combination of *H and 13C NMR spectroscopy. These reactions were repeated using H S 0 3 F, where it was hoped that the resulting fluorosulphate anions would facilitate the formation of crystals, and so allow the definitive characterisation of these species as solids. Fluorosulphonic acid was condensed separately into three FEP tubes which contained the transition metal carbonyl clusters [Ir4 (CO)12], [Ru3 (CO)12] and [Os3 (C O )12] respectively. All three solids dissolved immediately and, in the case o f [Ir4 (CO)12], this is in stark contrast to the AHF reaction where dissolution occurs only slowly over eight hours. The FEP vessels were sealed and analyses were undertaken by JH and 13C NMR spectroscopy. The above reactions were repeated in an identical manner and the fluorosulphonic acid was slowly removed under vacuum. The intention was to isolate crystals which could be analysed by X-ray diffraction, however, in each case fine powders were obtained and attempts to recrystallise these powders using dried solvents did not result in the formation of any suitable crystals. 161 Triosmium dodecacarbonyl dissolved in HSO 3 F at room temperature to give a yellow solution for which the ]H and 13C NMR revealed the presence of three protonated species. The NMR spectra revealed that these three species were identical to those observed during the dissolution of [Os3 (CO)12] in AHF.[81] A proton resonance at 8-8.2 ppm (cf 8-8.5 ppm in HF) and a 13C resonance at 8160.8 ppm (d 2 /(C H ) = 3.1 Hz) [c f 8159.0 ppm (d 2 /(C H ) = 3.0 Hz) in HF] were assigned to the mononuclear [Os(CO)5 H]+, no resonance could be observed for the CO,ran5 -H, presumably because this species is a minor component in the solution. This species was previously confirmed by a proton study of the reaction of [Os(CO)5] with 98% H 2 S 0 4 J 82^ Also observed were a proton resonance at 8-19.6 ppm (cf 8-20.3 ppm in HF) and five 13C resonances at 8173.7 (d 2 /(C H ) = 3.2 Hz), 8170.8, 8168.9, 8163.9 and 8161.8 ppm (d 2 /(C H ) = 7.1 Hz) (cf 8176.4, 8169.8 (d 2 J(CH) = 2.9 Hz), 8165.1, 8167.0 (d 2 /(C H ) = 2.0 Hz) and 8159.5 ppm (d 2 /(C H ) = 7.0 Hz) in HF), and these were assigned to [Os3 (CO) 1 2 H]+J 81,83^ Finally, a proton resonance at 820.1 ppm and 13C resonances at 8155.2 and 8154.1 ppm were evident, these are consistent with the AHF reaction, however, this third minor complex remains unassigned. The proton NMR spectrum of the golden orange solution of [Ru 3 (CO )12] in H S 0 3F solution showed two hydride signals at 8-7.2 and 8-19.1 ppm (c f 8-7.9 and 8-19.4 ppm in AHF). The latter resonance, characteristic of a bridging hydride, is assigned to [Ru 3 (CO) 1 2 H]+ and also corresponds to the peak at 8-19.4 ppm reported as the only resonance from a solution of [Ru 3 (C O )12] in 98% H 2 S 0 4 J 80^ The first peak corresponds to that at -7.2 ppm reported for [Ru(CO) 5 H]+ from the reaction of [Ru(CO)5] and concentrated H 2 S 0 4 J 82^ In order to observe a complete 13C NMR spectrum for these species, 13CO enrichment is required ^8 ^ and to record the spectrum at low temperature. This was deemed unnecessary as the reactivity of this system appears identical to that observed for the dissolution of [Ru3 (CO)12] in AHF. 162 A yellow solution of [Ir4 (CO)12] dissolved in HSO 3 F exhibited a single hydride resonance at 8-19.6 ppm (c f 8-20.0 ppm in HF). This species has been previously characterised as [Ir4 (CO) 1 2 H2]2+ on the basis of an accurate integral of the hydride resonance against a weighed amount of Me 2 S 0 4 J 80^ The 13C NMR spectra exhibited two resonances at 8146.7 (dd 2 J(CH) = 2.1 Hz) and 8144.6 ppm (d 2 J(CH) = 20.9 Hz) in an approximate 2:1 ratio with coupling constants indicative of COcl5-H and COtrans-H (Figures 4.7 and 4.8). The assignments are in excellent agreement with those reported for the dissolution of [Ir4 (CO)12] in AHF c f 8143.7 (2 J(CH) = 2.2 Hz) and 8142.0 (2 J(CH) = 21.2 H z ).^ Figure 4.7. Proposed structure of [Ir4 (CO) 1 2 H2] 2 + 2+ CO, OC: CO, OC; OC CO. CO, CO, aTrans hydride. bCis hydride. CO, Figure 4.8. Carbon-13 and 147.2 147.0 13C{ 1H} NMR spectra of [Ir4 (CO)12] in HSO 3F. 146.8 146.6 146.4 146.2 146.0 145.8 145.6 145.4 145.2 145.0 144.8 144.6 144.4 144.2 (ppm) 4.9.1. Summary. The reactivity of fluorosulphonic acid with the three metal carbonyls [Ir4(CO)12], [Ru 3 (CO)12] and [Os3 (CO)12], appears to be similar to that in anhydrous HF, which is hardly surprising in view of their identical Hammett acidity functions (H 0 = -15.1). Both acids are considerably stronger than concentrated sulphuric acid and this increased acidity may account for the presence of [Ru(CO)5 H]+ and [Os(CO)5 H]+. The osmium monomer was observed in sulphuric acid, but only after heating to 100°CJ82^ The ruthenium monomer was not observed in this system, and heating of the solution to 100°C 164 only resulted in decomposition. With AHF or HSO 3 F, [Os(CO)5 H]+ and [Ru(CO) 5 H]+ were observed at room temperature, and the ruthenium monomer was observed in a greater abundance than the osmium analogue. This may be a consequence of the metal-metal bond strength, which has previously^84! been noted to increase down the group. In the osmium case it was also noted that strong acid media such as AHF or HSO 3 F resulted in the formation of an unknown bridging hydride species J 81! The presence of a significant resonance in the lH NMR spectrum, for the osmium reaction, suggests that this species must contain more than one equivalent bridging hydride to account for its the relative intensity. Therefore, it appears that the use of very strong protonic acids has two noticeable affects over the less acidic concentrated sulphuric acid, namely:- i) strong acids appear to promote cluster fragmentation and ii) for the [Os3 (CO)12] reaction a third hydride species is noted, as yet unassigned. In the case of [Os3 (CO)12], the formation of a mono hydrogen bridged complex occurred readily. The use of the strong superacids, AHF and H S 0 3 F, may result in the formation of a bis or tris hydride bridged complex, [Os3 (CO) 1 2 H2]2+ or [Os3 (CO) 1 2 H3]3+ respectively. et Deeming demonstrated that the substitution of carbonyl ligands in [0 8 3 (0 0 al )!2] by tertiary phosphines, increased the Lewis basicity of the complex, and made multiprotonation more favourable.^82! The alternative to this, is to increase the acidity of the proton donating species, thereby making multiprotonation more favourable. The formation of a bis hydride complex, which would result in five different carbonyl environments and therefore, five different resonances in the 13C NM R spectrum, does not fit the 13C NMR data obtained. Presently there are two unassigned resonances, however, the formation of the tris hydride would result in just two carbonyl environments, those carbonyls trans to a hydride and an equal number cis. The formation of such a complex would be 165 expected to carry some multiplicity. No multiplicity is observed for the unassigned signals, this may be the result of fluxional behaviour. It seems apparent that the use of a conjugate superacid system may unravel this. The use of a system such as A u(S 0 3 F ) 3 / H S 0 3F would serve two purposes:- i) A system such as this offers increased acidity, which should increase the amount of the monomer or the unknown bridged hydride species, ii) The [A u(S 0 3 F)4]‘ conjugate base is a bulky anion of low nucleophilicity. It may prove ideal for the growth of crystals of [Ir4 (CO) 1 2 H2]2+ and [M 3 (CO) 1 2 H]+ (M = Ru and Os). The obvious drawback of such a system, however, is the potential oxidising ability of the Au(III) centre. 4.10. The Reaction Between [Fe2 (CO)9 ] and HSO3 F. Iron is rapidly oxidised in moist air and in its finely divided form is pyrophoric. It readily dissolves in dilute mineral acids which, in the absence of air or oxidising acids, produces Fe(II)J67,85^ The use of warm dilute nitric acid or the presence of air usually results in the formation of some Fe(III). Strong oxidising media, such as concentrated H N 0 3, passivate the metal and prevent complete reaction. Brazier and W oolf observed that iron did not react with boiling fluorosulphonic acid[49^ which, in view of the low oxidising prowess of sulphuric acid, is surprising. This lack of reactivity may be the result of passivation. The reaction between [Fe3 (CO)12] and H 2 S 0 4, as noted earlier, results in decomposition of the carbonyl cluster As a consequence of this, the reaction between [Fe2 (CO)9] and H S 0 3F was attempted as a convenient route to [F e(S 0 3 F)2] or [Fe(S0 3 F)3]. Fluorosulphonic acid was condensed on to dark yellow platelets of [Fe2 (CO)9] in a FEP tube at -196°C. On warming to room temperature, a reaction commenced as evidenced by the evolution of a gas. Analysis of the gas by gas-phase infrared spectroscopy identified it as carbon monoxide. Removal of the fluorosulphonic acid produced a dark green solid. 166 Analysis of the solid was undertaken by mass spectrometry and infrared spectroscopy. The vibrational spectroscopic data is compared to that previously published in the literature for [Fe(S0 3 F)2],[22] and is presented in Table 4.10. The vibrational data is relatively simple to interpret. Using the information provided in Section 4.5.1, it can be seen that no splitting of the E modes was observed and therefore, the fluorosulphate anion must possess C3v symmetry. The shift of the v(S-F) relative to [K (S0 3 F)], indicates that the anion must be behaving as a tridentate bridging group. A comparison of the vibrational data presented in Table 4.10 offers conclusive evidence that the reaction between [Fe2 (CO)9] and H S 0 3F affords [Fe(S0 3 F)2] . In view of the polymeric nature of [F e(S 0 3 F)2] it is understandable why mass spectrometry failed to produce any identifiable patterns. Table 4.10. Infrared spectroscopic data for K[SQ 3 F] and [Fe(SQ 3 F)2]. K [S 0 3 F]a [Fe(S0 3 F)2]b [Fe(S0 3 F)2]c cm"1 cm'1 cm"1 1280 s 1270 s 1261 vs v 4 (E) 1080 s 1171 s 1181 s Vi (Aj) 750 s 865 s 862 s v2 (Al) 590 s 611 s 610 s v 5 (E) 570 m 573 m 568 m v 3 (Aj) 480 m - 419 m v 6 (E) a Ref. 14. b This work. 0 Ref. 22. 167 Assignment. 4.11. The Reaction Between Re or Mn Carbonyl Derivatives and HSO3F. The complexes [M (C0) 5 (S 0 3 F)] (M = Mn and Re) have been synthesised previously The compounds are produced by the reaction of [M(CO)5 X] (M = Mn, Re; X = Cl, Br) with A g[S 0 3 F] in a suitable solvent. In the case of rhenium the reaction proceeds very smoothly, however, the manganese reaction is rather slower and complete substitution is only observed when the reaction is carried out over five days using [Mn(CO)5 Br]. Four reactions were attempted and these used the readily available starting materials [Mn 2 (CO)10], [MeMn(CO)5], [Re2 (CO)10] and [Re(CO)5 Cl]. All the reactions were carried out in FEP vessels, and the fluorosulphonic acid was condensed into the tubes at -196°C. The reaction between [Mn2 (CO)10] and HS 0 3F commenced upon warming the mixture to room temperature, as evidenced by the production of a gas which was identified by gas phase infrared spectroscopy as carbon monoxide. During the course of the reaction a solid was precipitated. Once the reaction was judged to be complete, the fluorosulphonic acid was removed under reduced pressure and a dark green solid was isolated. Analysis was undertaken using mass spectrometry and infrared spectroscopy. Mass spectrometry failed to produce any identifiable patterns and infrared spectroscopy also met with no success. The infrared spectrum showed no carbonyl absorptions and the sulphur-oxygen and sulphur-fluorine region consisted of several very broad absorptions. No useful information was obtained and repetition of the experiment provided no improvements. The solid was insoluble in a range of solvents, including fluorosulphonic acid, and this suggests a polymeric nature. The reaction between [MeMn(CO)5] and H S 0 3F occurred upon warming the mixture to -78°C, and continued for approximately 30 minutes. 168 Analysis of the gas produced, by gas-phase infrared spectroscopy, showed only the presence of methane. During the course of the reaction, solid was precipitated. On completion of the reaction the fluorosulphonic acid was removed under reduced pressure. Attempts were made to obtain infrared and mass spectral data, however, as observed for the [Mn 2 (CO)10] reaction, no characterisable spectra were produced. The infrared spectrum showed two strong absorptions at 2131 and 2087 cm-1. This indicated the presence of carbonyl groups within the product, and implied a different reaction scheme to that observed for the [Mn2 (CO)10] reaction. A comparison of these carbonyl absorptions to those of the starting material [MeMn(CO)5] (cf. 2082, 1997 and 1947 cm '1) and the anticipated product [M n(C0) 5 (S 0 3 F)] (cf. 2140, 2056, 2030, 2000, 1972 cm '1) indicates neither was present in the product. The region 1500-600 cm ' 1 was dominated by strong, broad absorptions and no information could be obtained. The reactions between [Re2 (CO)10] or [Re(CO)5 Cl] and fluorosulphonic acid occurred steadily at room temperature and, upon removal of the fluorosulphonic acid, produced a brown coloured solid. Analysis of the gas evolved from the [Re(CO)5 Cl] reaction did not show the presence of HC1 but rather HF. This was unexplained, but may be the result of an exchange reaction occurring within the solution. Both reactions produced a solid of the same colour, however, it is noted that the reaction between [Re(CO)5 Cl] and A g [S 0 3 F] produced a white crystalline solidJ19^ The infrared spectroscopic data is presented in Table 4.11 and is compared with the previously published data for [Re(C0) 5 (S 0 3 F)] and K [S 0 3 F]. T w o important features highlight the ionic nature of the fluorosulphate group within the product. The sulphur-fluorine stretch is observed in the same region as that for K [S 0 3 F]: any strong covalent interaction at the oxygens is expected to strengthen this bond. The symmetric 169 Table 4.11. Infrared vibrational data for [Re(C0) 5 (S 0 3 F)]. K [S 0 3 F]b Assignment [Re(C0) 5 (S 0 3 F)]b [Re2 (CO)10] + [Re(CO)5 Cl] + h s o 3f h s o 3f Assignment0 - - 2160 w sh 2169 w 2165 w A v(CO) - - 2040 vs 2044 s 2045 s Ev(CO ) - - 1980 vs 1963 s sh 1967 s sh A' v(CO) - - 1315 m 1376 s 1370 s v 7 (A") 1280 s v 4 (E) 1255 m 1234 s 1232 s v 4 (A ) - - 1170 w 1170 w 1170 w v (M—0 ) - - 1 1 2 0 w 1150 w 1150 w v (M -O ) 1080 s Vl (A’) 1030 m 1050 vs 1055 vs Vi (A ) 750 s v 2 (A’) 760 m 748 vs 749 vs v 2 (A') 590 s V5 (E) 590 s 590 s 590 s v 5 (A ) - - - 578 w sh 579 w sh V8 (A") 570 m v 3 (A ) 560 m 552 m 553 m v 3 (A ) 480 m V6 (E) 340 s - - v 6 (E) a Ref. 14 b Ref. 19 c Using assignments previously made by Aubke et al. S 0 3 stretch is also found in the same region as that for K [S 0 3 F] and the splitting of the asymmetric S 0 3 stretch, observed at 1280 cm ' 1 for K [S 0 3 F], is indicative of a departure from C3v symmetry. Removal of the degeneracy for this E mode is small, but not as small as that observed by Aubke et a l. It is noted here that the splitting is - 1 2 2 cm ' 1 as opposed to that reported earlier of 60 cm"1. As previously stated, these splittings are too large to be due to site effects and are interpreted as indicative of an ionic interaction as opposed to a covalent one, this is apparent by comparison of this data with that in Table 4.5. The carbonyl stretching region contained three absorptions which are expected for a pentacarbonyl derivative. A comparison of this carbonyl stretching data with that in Table 2.14, again highlights the highly ionic nature of the bonding within this complex, similar to that observed for the seflate and teflate derivatives. The data provided here thus indicates that the reaction between fluorosulphonic acid and [Re2 (CO)10] or [Re(CO)5 Cl] produces [R e(C 0) 5 (S 0 3 F)] as the major product. Confirmation of the formation of [Re(CO)5 (S 0 3 F)] was provided by El and FAB mass spectrometry. For both experiments, the following species were identified:- [R e(C 0) 5 (S 0 3 F)]+ m/z = 426, [Re(C0) 4 (S 0 3 F)]+ m/z = 398 and [Re(CO)5]+ m/z = 327. The 13C NMR spectra on the products of both reactions were recorded in H S 0 3 F. The reaction between [Re2 (CO)10] and H S 0 3F produced a 13C NMR spectrum with two resonances at 5182.1 and 5179.3 ppm, relative intensities 4:1. On the basis of the previous evidence these resonances are assigned to the species [R e(C 0) 5 (S 0 3 F)]. The reaction between [Re(CO)5 Cl] and H S 0 3F showed four resonances in the 13C NMR spectrum, 5182.0 and 5179.3 ppm with relative intensities 4:1, and 5181.4 and 5176.5 ppm with relative intensities 4:1. The two sets of signals have an integration ratio of 3:1, suggesting that the major product constitutes 75% of the species produced. Considering the vibrational spectra, which did not show absorptions due to the 171 presence of [Re(CO)5 Cl], and the similarity of the chemical shifts to that observed for the [Re2 (CO) 1 0 ]-HSO3F reaction, the major species formed was [R e(C 0) 5 (S 0 3 F )]. The second species present was undoubtedly [Re(CO)5 Cl] and the failure to observe the expected v(CO) absorptions at 2155, 2046 and 1983 cm '1, was a consequence of the presence of the strong absorptions of [R e(C 0) 5 (S 0 3 F)] which dominated that region of the spectrum. Separation of the products was attempted using a range of solvents, however, the solubilities of the two materials appeared very similar. 4.12. The Reaction Between HSO3 F and [CP2 MX 2 ] (M = Ti, Zr or Hf, X = Me or Cl). In order to extend the range of fluorosulphate derivatives an attempt was made to synthesise [Cp2 M (S 0 3 F)2], where M = Ti, Zr or Hf. Recent work carried out at Leicester revealed that [Cp2 MX2], where X = Me or Cl, underwent clean displacement reactions with teflic acid, HOTeF5, to form [Cp 2 M (OTeF5)2] Analysis of the products revealed bis-teflate substitution at the metal centre and intact r|5-cyclopentadienyl ligands. Proton-1 NMR spectroscopic data for [Cp2 MX2] (M = Ti, Zr or Hf, X = Cl or OTeF5) and the teflate spectroscopic data indicated that the metal-teflate bond was extremely ionic in nature. Complexes of the general type [Cp2 MX2] (where X = Cl, alkyl or aryl) are well studied and exchange or reduction reactions are of significant synthetic potential for example, in alkene polymerisation. The reactions between fluorosulphonic acid and [Cp2 MX2] (M = Ti, Zr or Hf, X = Me or Cl) were attempted. Samples of [Cp2 TiCl2], [Cp2 ZrMe2], [Cp 2 ZrCl2] and [Cp2 HfCl2] were placed separately into passivated FEP tubes. Using a metal vacuum line, excess of fluorosulphonic acid was condensed on to the samples at -196°C. On warming to room temperature, in each case an immediate reaction occurred, as evidenced by the evolution of a gas. The 172 reactions were extremely vigorous and required quenching several times with an acetone-cardice slush. The gases generated were identified using infrared spectroscopy as methane for the [Cp2 ZrMe2] reaction and hydrogen chloride from the other reactions. On completion, all four reactions gave black, viscous, almost solid materials. Excess of fluorosulphonic acid appeared to be incorporated in the product mixtures, and it proved impossible to remove under reduced pressures. The materials formed in each case were undoubtedly polymeric and analysis by a series of spectroscopic techniques failed to give any indication about their composition. Excess of fluorosulphonic acid was condensed on to a sample of orange [Cp 2 TiM e2] at -196°C. On warming to -78°C an immediate reaction occurred as evidenced by the evolution of a gas. The gas was identified as methane using gas phase infrared spectroscopy. The reaction continued for approximately twenty minutes after which time it appeared to be complete. The resulting solution was black as in the above examples, however, the mixture was significantly less viscous. Removal of the H S 0 3F proved difficult, and elevated temperatures (ca. ~100°C) and prolonged pumping under dynamic vacuum were required. The isolated solid was dark purple in colour. Analysis of the solid by FAB mass spectrometry identified the fragment [Cp 2 T i]+ m/z = 1 7 8 . The infrared vibrational data are presented in Table 4.12, also listed, for comparison, are the infrared vibrational data for K [B r(S0 3 F)4], for which a covalent monodentate interaction exists between the fluorosulphate ligand and the bromine centred13^ The infrared spectroscopic data provide conclusive evidence for the formation of a covalent monodentate interaction between the titanium centre and the fluorosulphate group. The distinguishing features of the spectrum are contained in the region 1500-800 cm '1, i.e. the sulphur-oxygen and sulphur-fluorine stretching region (Section 4.5.1). A total of 7 absorptions were observed in the above region and this indicated that the symmetry of the fluorosulphate anion had been lowered to Cs (N.B. 3 173 absorptions were expected for a fluorosulphate anion with C3v symmetry). This lowering of symmetry indicated the presence of a mono- or bi-dentate interaction. Furthermore, the sulphur-fluorine stretching frequency was significantly shifted to higher wave numbers, which indicates a covalent interaction. The covalent nature was also emphasised by the magnitude of the splitting of the v 4 mode, v 4 —> v 7 (1440 cm '1) + v 4 (1047 cm"1), which at 393 cm " 1 is similar to that observed for K [B r(S0 3 F)4], 454 cm '1. The data clearly indicates, that in the solid phase, a covalent monodentate interaction exists between the titanium centre and the fluorosulphate anion. Table 4.12. Infrared spectroscopic data for [Cp2 T i(S 0 3 F)2] and K [B r(S0 3 F)4]. K [B r(S 0 3 F)4] / cm - 1 [Cp2 T i(S 0 3 F)2] / cm ' 1 Assignment - 2968 s sh v(CH) - 2929 s v(CH) - 2865 s sh v(CH) 1424 m 1440 m v7 1407 w 1401m 970 m 1047 m - 1023 s sh 1237 s 1 2 2 0 w sh 834 s 1236 s br V4 Vl 1205 m sh 8 8 6 m v2 615 vs 603 s - 578 ms 587 s sh - 553 ms - - 406 w - - 239 vs - - 174 Multinuclear NMR spectra were recorded for the reaction mixture, [Cp 2 TiMe 2 ]-H S 0 3 F. Proton-1 NMR experiments showed a single resonance at 87.2 ppm which originated from the equivalent protons of the rj5- cyclopentadienyl groups. Fluorine-19 NMR experiments showed a doublet at 5 57.8 ppm, J = 5.8 Hz, which became a singlet when the experiments were run in proton decoupling mode. Carbon-13 ^ H ) NMR experiments showed a singlet at 8129.3 ppm, which originated from the r\5-cyclopentadieny 1 groups. A comparison of the *H NMR chemical shifts for the compounds [Cp2 TiX2] (X = -S 0 3 F, -OTeF5, -F and -Cl) is presented in Table 4.13. As can be seen, the resonance arising from the cyclopentadienyl protons is shifted to higher frequency as one ascends the table. A similar pattern was observed for the Zr and H f derivatives, but are not due to steric effects Excluding the fluorosulphate group, the trend observed is interpreted in terms of the ability of the ligands to undergo p n —» dn bonding, fluorine being the strongest n donor but the weakest n acceptor. As the ability to p n —> dn bond decreases so the Cp-metal interaction increases. The ]H NMR data for [Cp2 T i(S 0 3 F)2] is not strictly comparable as solvent effects have not been accounted for. However, the NMR data suggests, that the interaction of the fluorosulphate groups in solution is highly ionic in nature. This is not surprising in view of the high ionising ability of fluorosulphonic acid which presumably dominates in favour of covalent bonding. The resulting interaction requires the Cp ligands to donate more electron density to the metal centre, therefore, resulting in deshielding of the cyclopentadienyl protons. In the solid state, however, the fluorosulphate groups are covalently bound to the titanium centre. This is in contrast to [Cp2 Ti(OTeF5)2], where the spectroscopic data clearly indicated that the titanium-teflate bonds possess a large degree of ionic character. This obviously reflects the fundamental differences between the teflate and fluorosulphate groups as ligands. Although the central atoms, S and Te, are both in their maximum oxidation states their 175 geometries are completely different. Due to the fact that the fluorosulphate group is tetrahedral and the teflate group is octahedral, different sized molecular orbitals are formed and, presumably, in the case of the fluorosulphate anion these orbitals are suitable to undergo covalent bonding with the titanium metal centre. Table 4.13. Proton-1 NMR chemical shifts for [Cp2 TiX2] (X = -SO 3 F, -OTeF5, -F and -Cl). Compound 8 [Cp2Ti(S0 3 F)2]a 7.2 [Cp2 Ti(OTeF5)2]b 6.9 [Cp2 TiCl2]b 6.56 [Cp 2 TiF2]b 6.44 lH / ppm a Recorded in HSO 3 F. b Recorded in CD 2 C12. Ref. 8 6 . The 19F NMR spectra showed a resonance at 857.8 ppm, and Table 4.14 lists this and other recently reported values for other covalent monodentate ligands. For the molecules Cs 2 [Pt(S0 3 F)6], Cs[Sb(S0 3 F)6] and Cs 2 [A u(S 0 3 F)4] which are listed in Table 4.14, a strong covalent interaction is present between the metal centre and the fluorosulphate anion (Section 4.6). It is concluded, that the 19F NMR resonance for [Cp2 T i(S 0 3 F) 2 ]-H S 0 3 F, observed at 557.8 ppm, does not originate from a covalently bound fluorosulphate group, and three important observations support this:- i) The high frequency shift (5fluorosulphate-8compiex) was much larger than has previously been observed for fluorosulphate complexes, ii) The presence of proton coupling in such a molecule cannot be accounted for, i.e. coupling to the solvent is unlikely in view of the very low basicity of peripheral fluorine atoms expected in this type of molecule, and iii) The *H NMR spectral data (Table 176 4.13) suggests a highly ionic environment, when dissolved in H S 0 3 F. Therefore, it seems apparent that, in solution, the titanium complex present is best represented as [Cp2 Ti]2+, consistent with the high ionising ability of fluorosulphonic acid. Table 4.14. Fluorine-19 NMR chemical shifts for various covalent monodentate fluorosulphate complexes. Complex 8 19F / ppm HSO 3 F solvent peak 8 19F / ppm [Cp2 T i(S 0 3 F)2] 57.8 40.8 Cs 2 [P t(S 0 3 F)6]a 47.7 40.7 C s[S b(S0 3 F)6]a 46.4 40.9 Cs 2 [A u(S0 3 F)4] a 45.5 40.8 a Ref. 29. 4.13. The Reaction Between HSO3 F and [W(CO)6 J. The reaction between [W(CO)6] and HSO 3 F was carried out in an analogous manner to those reactions described earlier. Tungsten hexacarbonyl was loaded into a passivated FEP tube and attached to the metal line. All connections were leak tested and passivated. An excess of fluorosulphonic acid was condensed into the FEP tube at -196°C. On warming to room temperature a solvation commenced which continued at a very slow rate. The mixture was left for four days, after which time all the solid had gone into solution. The resulting solution was a red-brown colour. Carbon-13 NMR spectra of this solution showed a single resonance at 8 207.5 ppm accompanied by tungsten satellites / ( 1 8 3 W - 1 3 C) = 1 1 6 Hz. Proton-1 NM R experiments provided no evidence of any hydride containing species. 177 Attempts to remove the fluorosulphonic acid and obtain a solid sample proved difficult, elevated temperatures and prolonged pumping resulting in a black viscous oil. Mass spectrometry did not show any identifiable patterns, and infrared spectroscopy only revealed a single, strong absorption at 2004 cm '1. The spectroscopic evidence obtained for this reaction compares well to that reported in the literature for [W(CO)6].[87] The reported 13C NMR resonance for [W(CO)6] in CH 2 C12 is 5192.1 ppm, / ( 1 8 3 W - 1 3 C) = 126 Hz and the infrared carbonyl absorption is at 1998 cm '1. This has led us to conclude that no reaction occurs between [W(CO)6] and H S 0 3 F. The variation in the NMR data is almost certainly due to solvent effects and this has previously been highlighted (Table 2.5). The infrared spectroscopic data obtained for the reaction product is also consistent with that reported in the literature for [W(CO)6]. Our inability to remove all the fluorosulphonic acid and obtain solid [W(CO)6] is indicative of some degree of association. Several broad, rather small intensity absorptions observed in the region 1500-400 cm '1, are indicative of the sulphur-oxygen and sulphur-fluorine stretching region, and are almost certainly associated with HSO 3 F wrapped in the [W(CO)6] lattice. 4.14. The Reaction Between [Mo(CO)6 ] and HSO3 F. Molybdenum hexacarbonyl was loaded into a passivated FEP tube, attached to a metal vacuum line and all connections leak tested and then passivated. Fluorosulphonic acid was condensed on to the white solid at -196°C. Upon warming to room temperature, a slow reaction occurred as evidenced by the production of a gas which, was identified using infrared spectroscopy as CO. The reaction was left over a period of four days after which time no gas evolution was observed and it was judged to be complete. Removal of the fluorosulphonic acid was very difficult and required pumping 178 under dynamic vacuum for one week. The solid isolated was a dark blue, almost black, colour. Infrared spectroscopy of the residual solid showed the following absorptions; 2158 s, 2059 s, 1377 vs, 1106 vs vbr, 985 wsh, 841 m, 722 m, 619 w and 556 m cm '1. A comparison of the CO absorptions, 2158 and 2059 cm '1, to that for [Mo(CO)6], c f 2004 c m 'V 88^ shows a shift to higher wavenumbers indicating oxidation of the Mo centre. Proton-1 NMR experiments did not show any hydride containing species and carbon-13 NMR experiments showed a single resonance at 5203.3 ppm. The 13C NM R spectral resonance at 5203.3 ppm is due to the presence of unreacted [Mo(CO)6] {cf 5202.0 ppm for [Mo(CO)6] in CH 2 C12).[88] A comparison of the vibrational data with the data discussed in Section 4.5.1 shows the only certain feature in the fluorosulphate region is the sulphurfluorine stretch which, at 841 cm '1, implies a significant covalent interaction. The sulphur-oxygen region was dominated by a broad absorption centred at 1107 cm " 1 {N.B. 1320-950 cm '1), making a firmer assignment impossible. 4.15. The Reactions Between and [Co2 (CO)s] or [Cr(CO)6 ] and HSO 3 F. The reaction between fluorosulphonic acid and [Co2 (CO)8] or [Cr(CO)6] was carried out in an analogous manner to those described previously. Both reactions produced carbon monoxide gas and over the course of the reaction a black insoluble solid formed. Attempts to dissolve these solids in a variety solvents failed, and analysis by mass spectrometry and infrared spectroscopy produced no characterisable spectra. The infrared spectra did not show any absorptions in the carbonyl stretching region. 179 4.16. Summary. Although some success was obtained with these reactions, the combination of metal vacuum lines and FEP apparatus with fluorosulphonic acid is far from ideal. Work-up may take as long as a week and usually involves prolonged pumping under dynamic vacuum. Further development of these synthetic reactions is likely to be more suited to the use of glass reaction vessels and Schlenk vacuum lines. The reactions of fluorosulphonic acid with the organometallic complexes outlined in this Chapter is not a general route to metal fluorosulphate derivatives. However, in certain cases, identifiable products have been obtained. The compounds [Fe(S0 3 F)2] and [Re(CO)5 (S 0 3 F)], although previously available, have now been prepared by more direct routes. These reactions also demonstrate that fluorosulphonic acid possesses mild oxidising abilities. W oolf et al. previously reported that iron and rhenium metal are both inert to boiling fluorosulphonic acidj49^ so that the presence of the carbonyl ligands has clearly played a part in facilitating the oxidation of the metal (0 ) centre. The reaction between [Cp2 TiMe2] and H S 0 3F has established a route to the previously unknown [Cp2 T i(S 0 3 F)2] . This represents a new synthetic approach to obtaining fluorosulphate derivatives and also demonstrates, for the first time, that cyclopentadienyl and fluorosulphate groups are mutually compatible. It is likely, therefore, that further developments will occur in this area. 180 References Chapter Four [1] Thorpe, T. F. and Kirman W. J., J. Chem. Soc., 1892, 6 1 , 921. [2] A. Engelbrecht, Angew. Chem., Int. Ed. 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Chem. Soc., 1968, 2839. [85] N. N. Greenwood and A. Eamshaw, Chemistry o f the Elements, Pergamon, Oxford, 1st ed., 1984, ch. 25, pg. 1250. [8 6 ] M. C. Crossman, Ph.D. Thesis, University of Leicester, 1995 and references cited therein. [87] Comprehensive Oranometallic Chemistry, eds. E. W. Abel, F. G. A. Stone and G. Wilkinson, Pergamon Press, Oxford, 1982, vol 3, ch 28 and references cited therein. [8 8 ] Comprehensive Oranometallic Chemistry, eds. E. W. Abel, F. G. A. Stone and G. Wilkinson, Pergamon Press, Oxford, 1982, vol 3, ch 27 and refences cited therein. 185 CHAPTER FIVE Experimental 5.1. Handling of Materials. Most of the inorganic materials prepared and studied in this thesis are air- and moisture-sensitive. To prevent decomposition they were either handled on a metal vacuum line with facilities to connect glass or fluoroplastic reaction vessels via Teflon™ couplings or in an inert atmosphere dry box. 5.1.1. M etal vacuum line. This consisted of 316 stainless steel or Monel Autoclave Engineers' valves [AE-30 series, Autoclave Engineers Inc.] connected via Autoclave Engineers connectors. Argon-arc welded nickel 'U' traps were incorporated to permit separation and condensation of gases in the metal manifold. Inlets for argon [BOC Special Gases] and fluorine [Distillers MG] were positioned as shown in Figure 5.1. Rough pump vacuum outlets were connected to a soda lime chemical scrubber to neutralise any volatile fluorides, thereby protecting the rough pump [Model PSR/2, NGN Ltd.] which provided a vacuum of 10‘ 2 mmHg. High vacuum was obtained via outlets to a mercury diffusion pump coupled to a second NGN pump. This gave a vacuum in the region of 10‘ 5 mmHg. The mercury diffusion pump was protected by a glass trap in liquid nitrogen between the metal line and the diffusion pump to condense volatile products or fluorides from the metal line that remained after evacuation using the rough pump. The second rotary pump was protected by a second glass trap cooled with solid carbon dioxide to condense any mercury vapour before it could enter the rotary pump. 5.1.2. Inert atmosphere dry box. Involatile fluorides were manipulated in an auto-recirculating positivepressure dry box [Vacuum Atmospheres Co., VAC NE-42 Dri Lab.] 186 Figure 5.1. Metal vacuum line. Bourdon Gauge To high vacuum F„ hi let Ar Inlet To high vacuum A.E. Tee piece A E . Cross piece 9 = fi (r 8 = 1 m 13= ^ (r^ Metal connector A.E. tap To rough vacuum To rough vacuum To attached reactors To attached reactors attached attached reactors which provided a nitrogen atmosphere with an oxygen content of less than 5 ppm. The quality of the atmosphere was maintained by circulation through columns of molecular sieves and manganese dioxide which removed water and oxygen respectively. The dry box was equipped with a Sartorius electronic balance [Model 1601 MPS, Sartorius, Surrey, UK]. If static charge proved to be a problem during the weighing or transfer of materials, a Zerostat 3 anti-static gun was used to minimise the problem. 5.2. Reaction Vessels. 5.2.1. M etal reactors. Metal reactors (Figure 5.2) were always prepared, prior to use, by the following procedure. After evacuation to high vacuum, reactors were pre seasoned with either 500 mmHg of fluorine or the reaction pressure of fluorine, which ever was the greater (maximum reaction pressure of 1 0 atmospheres), for ca. one hour at either room temperature or the planned reaction temperature followed by re-evacuation to high vacuum. 5.2.2. Glass apparatus. Pyrex-glass apparatus were blown as required and equipped with Young's Greaseless taps [J. Young (Scientific glassware) Ltd.], Acton, London UK]. The most commonly used design is shown in Figure 5.3. Before use, each apparatus was pumped to high vacuum and seasoned with 500 mmHg of fluorine for ca. thirty minutes. After this time, the fluorine was pumped away and the apparatus re-evacuated to high vacuum. 188 Figure 5.2. Metal reactor. T o A u toclave E ngineers va lv e t W ater-cooled lid W ater out W ater in PTFE ‘O ’ R ing S tainless-steel body Figure 5.3. Glass apparatus. Y o u n g ’s greaseless tap 6 m m glass connector G lass v e sse l M olecular sieves Solvent 189 5.2.3. Fluoroplastic apparatus. In general, for the reactions carried out in fluoroplastic, a straightened 6 nun o.d. FEP reactor was first prepared by sealing at one end by heat moulding into a 7 mm i.d. glass tube. This was then connected to Chemcom coarsecontrol needle valves [Type STD/VC-4, Production Techniques] by a PTFE 'O' compression union. Before the introduction of the reagents, the system was evacuated to approximately 10' 4 mmHg to ensure that a vacuum-tight system had been obtained, passivated with 500 mmHg of fluorine and re-evacuated to high vacuum. Non-volatile products were loaded into the evacuated FEP tubes in a dry box and then placed back on the vacuum line. The connectors were re evacuated and passivated as above. Volatile reagents and solvents could be transferred into these tubes under static vacuum (Figure 5.4). For large scale reactions, 12 mm o.d. FEP tubes were prepared (sealed at one end by heat moulding into a 13 mm i.d. glass tube) and used in a similar manner. For small scale reactions, 4 mm o.d. FEP tubes were prepared (sealed at one end by heat moulding into a 5 mm NMR tube) and connected to the vacuum line in a similar manner to that described above. After reaction, the solvent was either removed to permit the analysis of the resulting product, or the tubes were sealed under vacuum, by heating with a small ring oven whilst the solution remained frozen at -196°C. The resulting sealed FEP tube and its contents could then be examined by NMR spectroscopy. Highly reactive and corrosive liquids were stored in passivated Kel-F vessels fitted with a Chemcom tap. These vessels were connected to a metal vacuum line in an identical manner to that above, thus facilitating transfer of these materials under static or dynamic vacuum. 190 Figure 5.4. Apparatus for the transfer of volatile reagents under static vacuum. A.E. valve o f vacuum line Chemcom needle valve A.E. connector with 6mm adaptor 9868 UTT) Youna s greaseless tap r J 1j 6mm Glass connector Chemcom tee coupling FEP Tubing (6mm o.d., 2mm i.d.) Glass vessel Molecular sieves Solvent FEP Tubing ( 4 mm o.d., 2 mm i.d.) 191 5.3. Analytical Techniques. 5.3.1. Nuclear magnetic resonance spectroscopy. 1 H, 19 F, 31P and 13C NMR spectra were recorded on a Bruker DRX 400 spectrometer at 400.13, 376.50, 161.97 and 100.61 MHz respectively and also 19F and 81Br NMR spectra were recorded on a Bruker AM 300 spectrometer at 300.13 and 81.09 MHz respectively. Spectra were recorded on air-sensitive samples in 4 mm o.d. FEP tubes held coaxially in 5 mm precision glass NMR tubes containing a small amount of D20 as the external lock substance (Figure 5.5). *H and 13C NMR spectra were referenced to external TMS, 19F NMR spectra to external CFC13, 31P NMR spectra to 85% H 3 PO 4 and 81Br NMR spectra to 1 M KBr in water, using the high frequency positive convention. 5.3.2. Infrared spectroscopy. Infrared spectra were recorded for solid samples either as dry powders or dispersed in Nujol mulls compressed between KBr plates, on a Digilab FTS40 FTIR spectrometer. For air-sensitive materials, sample preparation was performed in the dry box. Gas-phase spectra were recorded in a copper cell of length 10 cm fitted with AgCl windows. A seal was achieved between the windows and the cell body by two PTFE gaskets. 5.3.3. Mass spectrometry. Electron impact (El) and fast atom bombardment (FAB) mass spectra were recorded on a Kratos concept 1H double focusing, forward geometry, mass spectrometer. 3-nitrobenzyl alcohol was used as the matrix when 192 Figure 5.5. NM R samples fitted inside a 5 mm o.d. precision NM R tube. FEP Tube (4mm o.d., 3mm i.d.) Cap r' Precision Glass NMR Tube (5mm o.d.) Deuterated Solvent Sample 193 operating in positive FAB mode and the samples were introduced directly into the ionising chamber. 5.3.4. EXAFS spectroscopy. Bromine K-edge EXAFS data were collected at the Daresbury synchrotron radiation source operating at 2 GeV (ca. 3.2 x 10"1 0 J) with an average ring current of 205 mA on station 9.2 using a double-crystal Si (220) monochromator offset to 50% of the rocking curve for harmonic rejection. Selenium K-edge EXAFS data were collected under the same conditions with an average ring current of 227 mA on station 9.3 using a double-crystal Si (220) monochromator, offset to 50% of the rocking curve for harmonic rejection. The EXAFS data were collected in transmission mode for the solid samples diluted with fully fluorinated Teflon and closed in thin-wall FEP cells (Figure 5.6). The EXAFS data treatment utilised the programs EX^ and E X C U R V 92.^ Several data sets were collected for each sample in k space (k = photoelectron wave vector / A '1), and averaged to improve the signal to noise ratio. The pre-edge background was removed by fitting the spectrum to a quadratic polynomial, and subtracting from the whole spectrum. The atomic contribution to the oscillatory part of the absorption spectrum was approximated using polynomials and the optimum function judged by minimising the intensity of the chemically insignificant shells at low r (r = radial distance from primary absorbing atom) in the Fourier transform. To compensate for the decreased intensity at higher k, the data was multiplied by k3. Modelling and analysis was performed using EXCURV92, utilising curvedwave theory with phase shifts and back-scattering factors calculated using the normal ab initio methods. 194 Figure 5.6. FEP cell used for the collection of EXAFS data. Assem bled cell Gap 0.12 mm Wall 0.6 mm Exploded section Section of wall and gap FEP Stainless steel 5.4. Solvents. 5.4.1. Anhydrous hydrogen fluoride. Hydrogen fluoride (ICI pic) was distilled direct from the cylinder into a passivated Kel-F vessel fitted with a Chemcom tap. It was then dried for twelve hours with one atmosphere of fluorine. The fluorine was removed and the HF stored over BiF5. 5.4.2. Dichloromethane. Dichloromethane was purified and dried by first shaking it with portions of concentrated H2S 0 4. This was repeated until the acid layer remained colourless. It was then washed with water containing 5% Na2C 0 3 and then water again. The solvent was then pre-dried with CaCl2, distilled from P2Os and finally distilled from CaH2 under dry nitrogen. The dichloromethane was 195 stored in a glass Schlenk flask over dried 4 A molecular sieves. The solvent was degassed prior to use. 5.4.3. Acetonitrile. The acetonitrile was initially dried by shaking with 4 A molecular sieves. It was then stirred under nitrogen with CaH 2 for approximately four hours. The resulting liquid was then distilled on to fresh P 2 0 5 under nitrogen, retaining only the middle fraction. Finally the middle fraction again was distilled into a glass Schlenk vessel and stored over dry 4 A molecular sieves. It was degassed prior to use. 5.4.4. Fluorosulphonic acid. Fluorosulphonic acid, H S 0 3 F, was transferred, using a dry box, into a round bottom flask. The round bottom flask, which had previously been dried with fluorine, had a connector making it possible to attach it to the metal vacuum line. The fluorosulphonic acid was then degassed and transferred by distillation into a Kel-F vessel fitted with a Chemcom tap for storage until use. 5.5. Preparation of Fluorides, Oxide Fluorides, Seflate and Fluorosulphate species. 5.5.1. Preparation ofX eF 2- Xenon difluoride was prepared as described by H ollow ay,^ 1966. Xenon was mixed with a 10% excess of fluorine in a preseasoned glass bulb (1 litre volume). The reaction mixture was UV photolysed with mercury discharge 196 lamps (350 nm) for a week, after which time the reaction was considered to be complete. Unreacted xenon and fluorine were removed in vacuo. The xenon difluoride was purified by sublimation under dynamic vacuum through a trap at -78°C. The crystalline solid (yield 100%) was stored in a preseasoned FEP vessel in the dry box. 5.5.2. Preparation ofXe(O SeF5)2. Xenon bis(seflate) was prepared as described by S eppelt,^ 1986. Selenium dioxide, S e 0 2 (ca. 40 mmol) was loaded into a passivated autoclave reaction vessel containing a magnetic stirrer bar. The reaction vessel was then cooled to -196°C and SF4 (ca. 36 mmol) was condensed on to the S e0 2. The reaction vessel was sealed and then, under constant stirring, heated to 120°C for twelve hours. The metal trap of the vacuum line was cooled to -78°C and the contents of the reaction vessel was pumped through it under dynamic vacuum. Selenium oxide difluoride, SeOF2, the least volatile product of this reaction, was the only compound isolated in the trap. Xenon difluoride (ca. 29 mmol) was loaded into a passivated FEP Utube containing a magnetic stirrer bar. The FEP U-tube was connected to the line via Chemcom taps and the connectors were evacuated and passivated. The FEP U-tube was then cooled to -78°C and pumped to high vacuum. Under dynamic vacuum the SeOF 2 was condensed on to the XeF 2 and the FEP U-tube was then allowed to warm to room temperature. A steady reaction occurred. The reaction mixture was left open to the line and stirred for twelve hours. After this time, the system was considered to be at equilibrium and the volatile products were pumped away using the rough pump. The solid was pumped for a total of three hours to remove HF and XeF2. The solid white Xe(OSeF5)2, yield 2.8g (56%), was stored in an preseasoned FEP vessel in the dry box. 197 5.5.3. Preparation o f K[BrO 4]. Potassium perbromate was prepared using the method described by Appleman, 1 9 7 2 .^ The initial oxidation involves the action of elemental fluorine on an aqueous alkaline solution of K [B r03]. Consequently, the experiment was undertaken using fluoroplastic reaction vessels. Ice was packed around the reaction vessel to dissipate heat produced during the oxidation stage. Sodium bromate, N a[B r03], (ca. 1.3 mol) was added to a 900 ml solution of 5M NaOH and stirred mechanically until all the solid dissolved. A FEP tube was placed into the solution, the other end of the tube was connected to a metal line which, in turn, was connected via copper piping to cylinders of fluorine and argon gas. Elemental fluorine was passed into the solution and the metal line was used to control the rate. Fluorine was introduced into the solution at a rapid rate, however, care must be taken to avoid undue splattering. W arn in g , the reaction must be carefully monitored as, if the temperature of the solution approaches its boiling point, small detonations may occur in the vapour above the mixture. Care must also be taken to avoid deposits of solid blocking the end of the FEP tube. To remove any solid formed, the FEP tube was removed from the solution and the end cut off, the tube was then flushed with fluorine before introduction back into the solution. The oxidation stage was complete within one and a half hours and was monitored by measuring the pH of the solution. Oxidation only occurs in alkaline conditions and the pH was measured by extracting a drop of the solution using FEP tubing and placing it on universal indicator paper. Once the fluorination was complete (ca. solution turned acidic), the solution was flushed with a vigorous stream of argon gas (ca. 5 minutes) to remove unreacted fluorine and oxygen fluoride from the solution and the space above it. The resulting solution was colourless, free of any deposits and left to cool to room temperature (ca. 2 0 minutes). 198 The following stages involve the purification of K [B r04] where the use of glass vessels was avoided except when using the rotary evaporator. Washing the precipitate involves the use of distilled water to remove any K [B r04] from the precipitate, the filtrate was added to the original filtrate. Once the solution had cooled to room temperature, anhydrous Ba(OH ) 2 (ca. 1.75 mol) was slowly added to the solution. The temperature of the solution rose and it was stirred mechanically until it cooled to room temperature (ca. 1 hour). The solution was filtered and the precipitate washed several times. The precipitate, which consisted largely of BaF 2 and B a[B r03]2, was discarded. The solution was then acidified using Dowex 50X8 cation exchange resin, 20-50 mesh, in the hydrogen form. The pH of the solution was raised to 1.3, and the Dowex cation exchange resin served to remove sodium from the solution. The solution was then filtered and the precipitate washed and discarded. The volume of the solution had now risen to 2 litres and it was reduced to 400 ml using a rotary evaporator. The bromate concentration^ was assayed and this showed the presence of 4.96 g of unreacted bromate. Enough AgF (ca. 6 8 mmol) was added to the solution to provide a 0.15 M excess over the amount needed to precipitate the bromate present. The solution was filtered and washed with aqueous AgF (ca. 30 ml of 0.1 M AgF). Calcium hydroxide (ca. 41 mmol) was slowly added to the solution and this was sufficient to provide a 1 0 % excess above the amount needed to precipitate the fluorine added in the form of AgF. The solution was left for one hour to cool to room temperature and then the precipitate was removed by filtration and washed. The solution was acidified to pH ~ 1.3 using the Dowex cation exchange resin, filtered and washed. The solution was neutralised using Ca(OH ) 2 and then filtered using diatomaceous earth filter aid. The precipitate was washed with a saturated solution of Ca(OH)2. 199 Dowex cation exchange resin was then added to the solution, which was acidified up to pH 0.8. The solution was filtered, washed and then reduced in volume using a rotary evaporator to 2 0 0 ml. Initially, the solution was neutralised using 4 M KOH but, as the end point approached, 0.1 M KOH was used. The solution was then reduced in volume using a rotary evaporator. On cooling with ice, a brown solid was obtained which was then recrystallised from the minimum volume of water. The yield of the white K [B r04] was 6 g (2.5 %), it was dried in an autoclave vessel at 150°C and stored in a dry box. 5.5.4. Reactions involving Xe(OSeF5)2. The reactions of xenon bis(seflate) were carried out in identical ways. The reactions were performed using modified apparatus to prevent scorching of the materials used. A Chemcom tap was connected via Chemcom T-piece. Two separate 6 6 mm FEP tubing to a mm FEP tubes, sealed at one end, were connected to the T-piece. The Chemcom tap was then connected to a metal line, pumped to high vacuum and passivated with 600 torr of fluorine. Using an inert atmosphere dry box, xenon bis(seflate) (ca. 0.5 mmol) was placed in one of the FEP tubes. A stoicheiometric amount of the reactant (ca. [Re 2 (CO)10] 0.5 mmol) was then added to the remaining FEP tube. The vessel was attached to the line and the connectors were leak tested and passivated. The vessel was evacuated and dry dichloromethane was condensed on to the xenon bis(seflate). Once all the xenon bis(seflate) had dissolved it was decanted on to the reactant. The reaction was quenched, if required, using an acetone / cardice slush. On completion of the reaction all the volatiles were removed and the products were stored in the dry box. 200 5.5.5. Preparation ofB rF 3. Bromine trifluorine, BrF3, was distilled from the cylinder into a FEP Utube. The BrF 3 was brown due to the presence of bromine and this was removed by direct reaction with fluorine. Fluorine was slowly allowed into the tube were an immediate reaction occurred. Warning, the addition of fluorine must be very slow to avoid ignition. The brown colour slowly disappeared to leave a straw coloured liquid. Two atmospheres of fluorine were placed above the liquid and was left for two hours with constant stirring. The fluorine was removed at -78°C and the bromine trifluoride was transferred by distillation to a Kel-F vessel fitted with a Chemcom tap for storage. 5.5.6. Preparation ofB rF 5. This was distilled directly from the cylinder into an FEP U-tube. The BrF 5 was brown in colour due to the presence of bromine, also present were bromine trifluoride and HF. The bromine was removed by direct reaction with fluorine to produce bromine trifluoride. Two atmospheres of fluorine were placed above the mixture and left, with constant stirring, for two hours. The fluorine was then removed at -78°C. The U-tube was warmed to -13°C at which temperature bromine trifluoride has a vapour pressure of 0.3 torr and bromine pentafluoride a vapour pressure of 62 torr. The bromine pentafluoride was then distilled into a passivated Kel-F vessel which contained dried NaF: the NaF removed trace amounts of HF and BrF3. 5.5.7. Preparation ofK [BrF4], K[BrF6] and Cs[BrF6]. Dried KF (ca. 4 mmol) was loaded into a passivated 6 mm FEP tube fitted with a Chemcom tap, using an inert atmosphere dry box. The tube was connected to a metal vacuum line and all connections were leak tested and 201 passivated. Bromine trifluoride was condensed under static vacuum on to the KF. The reaction vessel was shaken for one week at room temperature. After this time the volatiles were removed and the solid complex stored in the dry box. Dried CsF (ca. 4 mmol) was loaded into a passivated 6 mm FEP tube fitted with a Chemcom tap, using an inert atmosphere dry box. The tube was connected to a metal vacuum line and all connections were leak tested and passivated. Bromine pentafluoride was condensed under static vacuum on to the CsF. The reaction vessel was shaken for one week at room temperature. After this time the volatiles were removed and the solid complex stored in the dry box. The preparation of K[BrF6] was carried out on a larger scale using a 12 mm FEP vessel. Potassium fluoride (ca. 16 mmol) was allowed to react with BrF 5 (ca. 62 mmol) as described above. The reaction vessel was shaken for one week after which time all the volatiles were removed and the solid stored in the dry box. 5.5.8 . Preparation o f [BrF 2 ][AsF6] and [BrF 4 ][Sb2Fj j]. A 6 mm FEP tube was connected via a satellite line to the metal manifold. Also connected to the satellite were BrF 3 and AsF 5 or BrF 5 and SbF5. The connectors and reaction vessels were leak tested and passivated. Bromine trifluoride (ca. 0.6 mmol) or bromine pentafluoride (ca. 0.4 mmol) were then condensed into the tube. Arsenic pentafluoride was allowed into the metal line and the reaction vessel. The uptake of AsF 5 was carefully monitored using the metal-line gauge and once the pressure remained constant all the BrF 3 was judged to have reacted. The solid was pumped under dynamic vacuum to remove excess of AsF 5 and then stored in the dry box. 202 Antimony pentafluoride (ca. 0.5mmol) was condensed on to the BrF5, which was in excess. An immediate reaction occurred at room temperature and the solid adduct was obtained by removal of the excess of BrF 5 using the rough pump. The solid was then stored in the dry box. 5.5.9. The Preparation o f Cs[BrOF4]. This was prepared using the method described by Chirste et al., 1987.^ Using an inert atmosphere dry box, C s[N 03] (ca. 2 mmol) was loaded into a passivated nickel reaction vessel. The reaction vessel was then attached to the metal line and the connectors were leak tested and passivated. A five molar excess of BrF 5 was condensed at -196°C into the nickel reaction vessel . The reaction vessel was warmed to -31°C and shaken occasionally for two to three hours. The reaction vessel was then re-attached to the metal line and the connectors were leak tested and passivated. The volatile products were removed using the rough pump and the solid product stored in the dry box. 5.5.10. The Preparation o f BrO^F. Potassium perbromate, K [B r04] (ca. 0.7 mmol), was loaded into a passivated 6 mm FEP tube in a dry box. The FEP tube, along with BrF 5 and HF, were attached to a metal line via a satellite connection. The connections were then evacuated and passivated. Enough AHF was condensed into the tube to completely dissolve the K [B r04] . Finally BrF 5 (ca. 2 mmol) was condensed into the FEP tube, and on warming to room temperature, an immediate reaction occurred. The reaction vessel was cooled to -78°C and at this temperature B r0 3F is the most volatile component of the reaction. The volatiles were condensed into an FEP tube containing dried NaF at -196°C, the NaF formed an adduct with the HF leaving a pure source of B r0 3 F. 203 Over time, B r0 3F decomposes with the formation of bromine. This can be kept to a minimum by storage of the B r0 3F at liquid nitrogen temperatures. The bromine may be removed by the addition of a small amount of fluorine into the FEP tube at liquid nitrogen temperatures. On warming the fluorine and bromine react to form BrF3, which itself reacts with the NaF to form a solid, involatile, adduct. 5.5.11. Reactions involving HSO 3 F. In a typical reaction, the solid metal complex (ca. O.lg) was weighed out in the dry box and loaded into a preseasoned FEP tube. The FEP tube was connected to the metal vacuum line via a Chemcom tap, and all the connections were leak tested, passivated and re-evacuated. The FEP tube containing the metal complex was then cooled to liquid nitrogen temperatures and leak tested. Fluorosulphonic acid (ca. 3.5g) was condensed into the reaction vessel at -196°C. The FEP tubing was warmed to -78°C using an acetone-cardice slush, and then slowly warmed to room temperature. The reaction was quenched, if necessary, using the acetone-cardice slush. On completion of the reaction all the volatile materials were removed. This involved distillation of the excess of fluorosulphonic acid into a second, empty, FEP tube. This procedure is very time consuming and care must be taken to avoid bumping of the H S 0 3F at reduced pressures. Once the excess of H S 0 3F was removed, the remainder of the volatile materials were pumped away under dynamic vacuum using the rough pump. This was also very timeconsuming since most of the reaction products appeared to adsorb the fluorosulphonic acid. This often required the use of elevated temperatures (ca. 100°C) and prolonged pumping (ca. one week). The solid products, if obtained, were stored in the dry box before analysis. 204 5.5.12. Attempted synthesis ofBrO F3. The reaction between L i[N 03] and BrF 5 was carried out using the method described by Christe et alP^ 1987. Using a dry box, L i[N 03] (ca. 2 mmol) was loaded into a gold-seal nickel reaction vessel. The reaction vessel was attached to a metal vacuum line and the connection was leak tested, passivated and then re-evacuated. The reaction vessel was then opened to the metal line and leak tested. Bromine pentafluoride (ca. 30 mmol) was distilled from a Kel-F storage vessel into the reaction vessel. The vessel was sealed and placed in a Dewar containing acetone. The acetone was cooled using a refrigeration unit to a temperature of 0°C. The vessel was left at 0°C for twenty days with occasional agitation. The vessel was then reconnected to the metal line and all connections were leak tested and passivated. The reaction cylinder was then cooled to -196°C and opened to the vacuum line, a small pressure rise was noted. The reaction was allowed to warm slowly to room temperature under dynamic vacuum. The volatile materials present were separated by fractional condensation through a series of traps at -64°C and -196°C. No volatiles were isolated in the trap at -64°C. Analysis using 19F NMR spectroscopy, at 0°C, only produced resonances due to BrF5. No resonances at or around 8 164 ppm were observed. Analysis of the solid product identified only K[BrF4]. In view of the presence of K[BrF4] the reaction was attempted at lower temperatures: in order to minimise the apparent decomposition of BrOF3. Subsequent reactions were carried out at -10°C and -20°C, in an analogous manner to that outlined above. Two further reactions were performed using larger and smaller excesses of BrF5. These reactions only produced BrF 5 and K[BrF4] as the identifiable products. 205 5.5.13. Attempted synthesis o f B r0 2F. This reaction was carried out using the method described by Gillespie et al., 1976.^ Potassium bromate (c.a. 8 mmol) was loaded into a passivated Kel- F vessel fitted with a Chemcom tap. The Kel-F vessel and vessels containing AHF and BrF 5 were attached to a metal vacuum line, via a satellite. All connections were leak tested, passivated and re-evacuated. The Kel-F reaction vessel was opened to the metal line, leak tested, cooled to -78°C and then BrF 5 (ca. 21 mmol) and AHF (ca. 0.3 mmol) were condensed into it. The tube was then slowly warmed to room temperature. Warning, the necessary safety protocols need to be followed as an explosion occurred the first time this reaction was performed (i.e. safety shields and screens). At room temperature a violent reaction occurred, accompanied by the production of a large volume of gas. After this initial phase, the violence of reaction subsided and a much smoother reaction occurred as evidenced by effervescence. This continued for approximately two hours and the resulting solution was a dark brown colour. The reaction mixture was cooled to -78°C and degassed. The volatile products were pumped under dynamic vacuum through a FEP trap cooled to -48°C (n-hexyl alcohol-cardice). At -48°C B r0 2F should have been the only volatile material collected in the trap, but no such material was obtained. 5.5.14. The attempted synthesis o f K [B r02F2] and K[BrOF4]. The following reactions were carried out using the method described by Gillespie et al., 1 9 7 6 Potassium bromate (ca. 1.3 mmol) and K[BrF6] (ca. 1.45 mmol) were loaded into a passivated Kel-F vessel in a dry box. The reaction vessel was attached to a metal line and all the connections were leak tested, passivated and re-evacuated. The Kel-F tube was cooled to -78°C using an acetone-cardice slush and leak tested. Acetonitrile (ca. 21 mmol) was 206 condensed into the vessel, which was then sealed and shaken for one day. The solvent was removed and the white solid stored in the dry box. The separation of K [B r0 2 F2] and K[BrOF4] involved the use of a glass vessel. Two glass containers, each fitted with a Young’s greaseless tap, were connected via a piece of glass tubing. The glass tubing contained a glass frit which enabled the filtration of the acetonitrile mixture; and one of the glass vessels possessed a 6 mm o.d. glass arm suitable for connection to a metal line. Using a dry box, the K [B r0 2 F 2 ]-K [Br0F4] mixture (ca. 0.3 g) was loaded into one of the arms of the passivated vessel. The vessel was then attached to a metal vacuum line and all the connections were leak tested and passivated. Acetonitrile (ca. 40 mmol) was condensed into the vessel, which was then sealed and shaken for two hours. The vessel was reattached to the metal line via FEP tubing and the liquid filtered into the second arm of the vessel. Removal of the acetonitrile under reduced pressure did not afford separation. 207 5.6. Sources of Chemicals and Methods of Purification. Antimony pentafluoride, SbF5 : Fluorochem. Used as supplied, stored and degassed in a glass Schlenk vessel. Arsenic pentafluoride, AsF5 : Fluorochem. Used as supplied. Bromine pentafluoride, BrF5 : Ozark Mahoning, now known as Atochem North America. Purified as described in Section 5.5.6. Bromine trifluoride, BrF3 : Fluorochem. Purified as described in Section 5.5.5. Fluorine, F2 : Distillers MG. This was used as supplied after being transferred into 1 dm 3 nickel cans for convenience. Fluorosulphonic acid, H S03F : Aldrich Chemical Company Ltd. Purified as described in Section 5.4.4. Hydrofluoric acid, HF : ICI pic. Purified as described in Section 5.4.1. Sulphur tetrafluoride, SF4 : ICI pic. Used as supplied. Xenon, Xe : BOC gases. Used as supplied. Acetonitrile, CH3CN : Aldrich Chemical Company Ltd. Dried and stored as described Section 5.4.3. Dichloromethane, CH2C12 : Aldrich Chemical Company Ltd. Dried and stored as described Section 5.4.2. Bis-cyclopentadienyl titanium dichloride, [Cp2TiCl2] : Aldrich Chemical Company Ltd. Used as supplied. Bis-cyclopentadienyl titanium dimethyl, [Cp2TiMe2] : Prepared according to the literature^ method. Dried in an autoclave vessel under dynamic vacuum at 100°C and stored in the dry box. Bis-cyclopentadienyl hafnium dichloride, [Cp2HfCl2] : Aldrich Chemical Company Ltd. Used as supplied. 208 Bis-cyclopentadienyl zirconium dimethyl, [Cp2ZrMe2] : Prepared according to the literature^ method. Dried in an autoclave vessel under dynamic vacuum at 100°C and stored in the dry box. Bis-cyclopentadienyl zirconium dichloride, [Cp2ZrCl2] : Aldrich Chemical Company Ltd. Used as supplied. Caesium fluoride, CsF : Aldrich Chemical Company Ltd. Dried in an autoclave vessel under dynamic vacuum at 150°C and stored in the dry box. Caesium nitrate, Cs[N03] : Aldrich Chemical Company Ltd. Dried in an autoclave vessel under dynamic vacuum at 150°C and stored in the dry box. Chromium hexacarbonyl, [Cr(CO)6] : Aldrich Chemical Company Ltd. Used as supplied. Dicobalt octacarbonyl, [Co2(CO)8] : Aldrich Chemical Company Ltd. Used as supplied. Diiron nonacarbonyl, [Fe2(CO)9] : Donated by Dr G. Capper, used as supplied. Dimanganese decacarbonyl, [Mn2(CO)10] : Aldrich Chemical Company Ltd. Used as supplied and stored in the fridge. Dirhenium decacarbonyl, [Re2(CO)10] : Aldrich Chemical Company Ltd. Used as supplied. Dowex 50 X8 20-50 mesh cation ion exchangers : Fluka. Used as supplied. Lithium Fluoride, LiF : Aldrich Chemical Company Ltd. Dried in an autoclave vessel under dynamic vacuum at 150°C and stored in the dry box. Lithium nitrate, Li[N03] : Aldrich Chemical Company Ltd. Dried in an autoclave vessel under dynamic vacuum at 150°C and stored in the dry box. Iodine, I2 : Aldrich Chemical Company Ltd. Used as supplied. Methyl manganese pentacarbonyl, [MeMn(CO)5] : Prepared according to the literature^10^ method. Dried in an autoclave vessel under dynamic vacuum at 100°C and stored in the dry box. Molybdenum hexacarbonyl, [Mo(CO)6] : Aldrich Chemical Company Ltd. Used as supplied. 209 Potassium bromate, K[Br03] : Aldrich Chemical Company Ltd. Used as supplied or dried in an autoclave vessel under dynamic vacuum at 150°C and stored in the dry box. Potassium fluoride, KF : Aldrich Chemical Company Ltd. Dried in an autoclave vessel under dynamic vacuum at 150°C and stored in the dry box. Rhenium pentacarbonyl chloride, [Re(CO)5CI] : Aldrich Chemical Company Ltd. Used as supplied. Ruthenium tris-carbonyl bis-triphenylphosphine, [Ru(CO)3(PPh3)2] : Prepared according to the literature ^ 1 ^ method. Dried in an autoclave vessel under dynamic vacuum at 100°C and stored in the dry box. Selenium dioxide, S e02 : Aldrich Chemical Company Ltd. Dried in an autoclave vessel under dynamic vacuum at 150°C and stored in the dry box. Silver fluoride, AgF : Aldrich Chemical Company Ltd. Used as supplied. Tetrairidium dodecacarbonyl, [Ir4(CO)12] : Aldrich Chemical Company Ltd. Used as supplied. Trisruthenium dodecacarbonyl, [Ru3(CO)12] : Aldrich Chemical Company Ltd. Used as supplied. Trisosmium dodecacarbonyl, [Os3(CO)12] : Aldrich Chemical Company Ltd. Used as supplied. Tungsten hexacarbonyl, [W(CO)6] : Aldrich Chemical Company Ltd. Used as supplied. 210 References Chapter Five [1] EX, A. K. Brisdon, University of Leicester, 1992. [2] EXCURV92, SERC Daresbury Laboratory Program, N. Binstead, J. W. Campbell and S. J. Gurman, 1992. [3] J. H. Holloway, J. Chem. Soc., Chem. Commun., 1966, 22. [4] K. Seppelt, D. Lentz and G. Kloter, Inorg. Synth., 1986, 24, 27. [5] E. H. Appleman, Inorg. Synth., 1972,12, 1. [6 ] E. H. Appleman, Inorg. Chem., 1969, 8, 223. [7] W. W. Wilson and K. O. Christe, Inorg. Chem., 1987, 26, 916. [8 ] R. J. Gillespie and P. Spekkens, J. Chem. Soc. Dalton Trans., 1976, 2391. [9] E. Samuel and M. D. Rausch, J. Am. Chem. Soc., 1973, 95, 6263. [10] R. J. Mckinney and S. S. Crawford, Inorg. Synth., 1989, 26, 155. [11] N. Ahmad, J. J. Levison, S. D. Robinson and M. F. Uttley, Inorg. Synth., 1974,15, 55. 211