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The Oxide Fluoride Chemistry of Bromine, Selenium and Sulphur

Thesis submitted for the degree of

Doctor o f Philosophy

at the

University of Leicester

by

LEE JO H N W O O TTO N

Department o

Faculty o f

University o

1997

UMI Number: U095961

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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 I7R H

i

Abstract

The Oxide Fluoride Chemistry of Bromine, Selenium and Sulphur

Lee J. Wootton

The transition metal carbonyls [Re2 (CO)10], [Mn2 (CO)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 [BrF2 ][AsF6], [BrF4 ][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 03F and a range of 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 [Ru3 (CO)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 4

1.3 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(OSeF5 ) 2 39

2 . 1 1 The Reaction Between [Mn2 (CO)10] and Xe(OSeF5 ) 2 46

2 . 1 2 The Reaction Between [Ru(CO)3 (PPh3)2] and Xe(OSeF5 ) 2 53

2.13 The Reaction Between I2 and Xe(OSeF5 ) 2 57

2.14 Discussion 6 6

References

Chapter Three

Bromine Oxide Fluoride Chemistry

70

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 The Halogen Oxide Trifluorides, XOF3 80

3.4.1 Chlorine oxide trifluoride 80

3.4.2 Bromine oxide trifluoride 81

3.4.3 Iodine oxide trifluoride 82

3.5 The Perhalogenyl Fluorides, X 0 3F 83

3.5.1 Perchloryl fluoride 83

3.5.2 Perbromyl fluoride 84

3.5.3 Periodyl fluoride 84

3.6 The Halogen Dioxide Trifluorides, X 0 2 F3 85

3.6.1 Chlorine dioxide trifluoride 85

3.6.2 Bromine dioxide trifluoride 8 6

3.6.3 Iodine dioxide trifluoride 8 6

3.7 The Halogen Oxide Pentafluorides, XOF5 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

3.8 The Halogen Oxide Fluorides 89

3.8.1 Chlorosyl fluoride 89

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[Br04] 109

3.15 The Synthesis of B r0 3F 110

3.16 The Attempted Synthesis of BrOF3 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 Syntheses involving S2 0 6 F2 or S2 0 6 F2 -H S 03F 129

4.3.1.1 Limitations of the S2 0 6 F 2 -H S 03F system 131

4.3.2 Displacement reactions 134

4.3.3 Syntheses involving B rS03F 136

4.3.4 Insertion reactions 136

4.3.5 Oxidising reactions involving H S 03F 137

4.4 Decomposition of Fluorosulphates 137

4.5 Spectroscopic Characterisation of Fluorosulphate 139

Compounds

4.5.1 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

4.6 Single Crystal X-ray Analysis of Fluorosulphate 147

Compounds

4.7 Recent Developments in Fluorosulphate Chemistry 153

4.7.1 Cationic carbonyl metal species 153

4.7.2 Superacids 157

4.8 Area of Study 160

4.9 The Reaction of [Ir4 (CO)12], [Ru3 (CO)12] and 161

[Os3 (CO)12] with H S 03F

vi

4.9.1 Summary 164

4.10 The Reaction Between [Fe2 (CO)10] and H S 03F 166

4.11 The Reaction Between Re or Mn Carbonyl 168

Derivatives and H S 03F

4.12 The Reaction Between [Cp2 MX2] (M = Ti, Zr or 172

Hf and X = Me or Cl) and H S 03F

4.13 The Reaction Between [W(CO)6] and H S 03F 177

4.14 The Reaction Between [Mo(CO)6] and H S 03F 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 Handling of Materials 186

5.1.1 Metal vacuum line 186

5.1.2 Inert atmosphere dry box 186

5.2 Reaction Vessels 188

5.2.1 Metal reactors 188

5.2.2 Glass apparatus 188

5.2.3 Fluoroplastic apparatus 190

5.3 Analytical T echniques 192

5.3.1 Nuclear magnetic resonance spectroscopy 192

5.3.2 Infrared spectroscopy 192

5.3.3 Mass spectrometry 192

5.3.4 EXAFS spectroscopy 194

5.4 Solvents 195

vii

5.4.1 Anhydrous hydrogen fluoride 195

5.4.2 Dichloromethane 195

5.4.3 Acetonitrile 196

5.4.4 Fluorosulphonic acid 196

5.5 Preparation of Fluorides, Oxide Fluorides, Seflate and 196

Fluorosulphate Species

5.5.1 Preparation of XeF2 196

5.5.2 Preparation of Xe(OSeF5 ) 2 197

5.5.3 Preparation of K [Br04] 198

5.5.4 Reactions involving Xe(OSeF5 ) 2 200

5.5.5 Preparation of BrF3 201

5.5.6 Preparation of BrF5 201

5.5.7 Preparation of K[BrF4], KtBrFg] andCs[BrF6] 201

5.5.8 Preparation of [BrF2 ][AsF6] and [BrF4 ][Sb2 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 03F 204

5.5.12 Attempted synthesis of BrOF3 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

5.6 Sources of Chemicals and Methods of Purification 208

References 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 X2, 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(OSeF5 ) 2 33

2.9 EXAFS and crystal data for Xe(OSeF5 ) 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 I2 and five molar equivalents

of Xe(OSeF5 ) 2

2.13 Fluorine-19 NMR spectral data for the products of 65

the reaction between I2 and three molar equivalents

of Xe(OSeF5 ) 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 76

compounds of bromine (V) and bromine (VII)

ix

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], [BrF2 ][AsF6] and [BrF4 ][Sb2 F n ]

3.4 EXAFS data for Cs[BrOF4]

4.1 Physical properties of H2 S 0 4, H S 0 3 F, S 0 2 F2,

CF3 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

[Co(S0 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 [Br(S0 3 F)4] and K [I(S0 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], [Sn(S0 3 F)2 Cl2],

K [Br(S0 3 F)4] and K [S0 3 F]

4.7 Comparison and assignment of the infrared vibrational

data for K [S 0 3 F] and [Re(S0 3 F)(C 0)5]

4.8 Bond lengths and angles for C s[S0 3 F]

4.9 Bond lengths and angles for Cs[Au(S0 3 F)4] and

Cs[Sb(S0 3 F)6]

4.10 Infrared spectroscopic data for K [S 0 3 F] and

[Fe(S0 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 [Br(S0 3 F)4]

4.13

4.14

Proton-1 NMR chemical shifts for [Cp2 TiX2] 176

(X = -S 0 3 F, -OTeF5, -F and -Cl)

Fluorine-19 NMR chemical shifts for various covalent 177

monodentate fluorosulphate complexes

xi

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 F5 SeOSeF5 28

2.4 The X-ray crystal structure of Xe(OSeF5 ) 2 32

2.5 Fuorine-19 NMR spectrum of Xe(OSeF5 ) 2 35

2.6 Background-subtracted EXAFS and the Fourier 37

transform spectra for Xe(OSeF5 ) 2

2.7 Fluorine-19 NMR spectrum for the product of the 41

reaction between [Re2 (CO)10] and Xe(OSeF5 ) 2

2.8 Carbon-13 NMR spectrum for the product of the 42

reaction between [Re2 (CO)10] and Xe(OSeF5 ) 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 NMR spectrum for the product of the 49

reaction between [Mn2 (CO)10] and Xe(OSeF5 ) 2

2.12 Carbon-13 NMR spectrum for the product of the 50

reaction between [Mn2 (CO)10] and Xe(OSeF5 ) 2

2.13 B ackground-subtracted EXAFS and the Fourier 5 2

transform spectra for [Mn(CO)5 (OSeF5)]

2.14 Fluorine-19 NMR spectrum for the products of the 55

reaction between [Ru(CO)3 (PPh3)2] and Xe(OSeF5 ) 2

2.15 Fluorine-19 NMR spectrum for the products of the 56


xii

2.16

2.17

2.18

2.19

2.20

3.1

3.2

3.3

3.4

3.5

3.6

3.7

3.8

Phosphorus-31 NMR spectrum for the products of the 56


Fluorine-19 NMR spectrum for the products of the 60

reaction between I2 and five molar equivalents of

Xe(OSeF5 ) 2


reaction between I2 and five molar equivalents of

Xe(OSeF5 ) 2


reaction between I2 and five molar eqivalents of

Xe(OSeF5 ) 2


reaction between I2 and three molar eqivalents of

Xe(OSeF5 ) 2

The isomeric forms of I 0 2 F 3 87

Background-subtracted EXAFS and the Fourier 98

transform spectra for K[BrF4]


transform spectra for Cs[BrF6]


transform spectra for [BrF2 ][AsF6]


transform spectra for [BrF4 ][Sb2 F n ]

Proposed reaction scheme for the reaction between 105

bromine pentafluoride and the alkali metal nitrates


transform spectra for Cs[BrOF4]

Fluorine-19 NMR spectrum of B r0 3F in BrF5 112

xiii

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 Molecular structures of [Au(S0 3 F)4]‘ and [Sb(S0 3 F)6] ' 149

4.5 Molecular structure of [A u(S0 3 F)3] 151

4.6 Crystal structure of mer-[Ir(S0 3 F)3 (CO)3] 156

4.7 Proposed structure of [Ir4 (CO)1 2 H2]2+ 163

4.8 Carbon-13 and 13C {1 H } NMR spectra of [Ir4 (CO) j2] 164

in H S 0 3F

5.1 Metal vacuum line 187

5.2 Metal reactor 189

5.3 Glass appartatus 189

5.4 Apparatus for the transfer of volatile reagents under 191

static vacuum

5.5 NMR samples fitted inside a 5 mm o.d. precision NMR 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 NMR 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

memorable 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 -C5 H5)

8 : NMR 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

NMR : 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 of doublets

( t ) : triplet

( q ) : quintet

cm '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 XeF8 is unknown,

whereas, X e 0 3 F2 and X e0 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 (CrOF3 and M n 0 3 F, CrF2 and

MnF3).

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, SeF6 and TeF6 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

state

Oxide Fluoride Oxide

fluoride

II - XeF2 -

IV - XeF4 XeOF2

VI X e0 3 XeF6 XeOF4

X e0 2 F2

VIII X e0 4 - X e 0 3 F2

As will be looked at in Section 2.4, polyatomic ligands such as -OSeF5

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. Br0 3 F[6] (pseudo tetrahedral: d(Br-C>) = 1.582(1) A, d(Br-F) =

1.708(3) A, ZOBrO = 114.9(3)° and ZOBrF = 103.3(3)°) using electron

diffraction, and SeOF2[7] (trigonal: </(Se-0 ) = 1.576(3) A, d(Se-F) = 1.729(1) A, ZFSeF = 92.22(10)° and ZOSeF = 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

[NMe4 ]+[IOF6] '[8] (pseudo pentagonal-bipyramidal: <7(1-0) = 1.775(6) A and

d(I-F) mean = 1.854(9) 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 intemuclear

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. Mn0 3F (rf(Mn-O) = 1.59(2) A and d(Mn-F) = 1.72(2) A) and

C r0 2 F 2 W(Cr-O) = 1.55(2) A and </(Cr-F) = 1.71(2) A).

3

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. Within 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 one-

electron 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 largeo

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 F2 and SOF4, and SeOF2, S e0 2 F2

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 S2 0 5 F2 - S7 O2 0 F2, and those whose structural group is

-SF5 e.g. SF5 OF, (SF5)20 and (SF5 O) 2. For selenium, a series of complex oxide

fluorides is known e.g. F5 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. (F5Te)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 -SO3 F and -OSeF5. The two

ligands have been described as “pseudo fluorides”, and indeed, the high valent

complexes [Sb(S0 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 -OSeF5 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,

27,157 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

Methods 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

Methods 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 eppelt^ in 1972.

This followed the similarly unexpected synthesis of pentafluoroorthotelluric

(VI) acid, HOTeF5, (teflic acid) by Engelbrecht^ and Sladky in 1964.

Engelbrecht et al. intended to synthesise T e0 2 F2 by combining Ba[Te04] and

H S 0 3F at 160°C, in a reaction analogous to that used in the preparation of

S e0 2 F2. Tellurium dioxide difluoride, T e0 2 F2, still unknown today, was not

isolated, but instead, HOTeF5 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

of non-metals; whereas an increase in coordination number is observed on

going to the fourth row, e.g. [SO4 ]2', [Se04]2" cf. [H4 T e0 6]2".

Seppelt was attempting to synthesis SeOF4 when he discovered HOSeF5.

The methods he employed were the same as those used to synthesis the

analogous SOF4. The fluorination of SeOF2 using a range of halogen fluorides

or elemental fluorine produced SeF4 and SeF6 respectively. However, the

fluorination of SeOF2 in the presence of HF afforded HOSeF5 (Eqn. 2.1).

Selenium oxide tetrafluoride has since been prepared by the vacuum pyrolysis

of Na[OSeF5 ] 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.

SeOF2 + F 2 + HF HOSeF5 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 HOTeF5 and HOSeF5 are strong Bronsted Lowry acids, and studies

have shown the acid dissociation constant of HOTeF5 to be between that of

concentrated hydrochloric and nitric acids with a pK& = 9 .2 .^ Further work by

Engelbrecht et al. demonstrated that HOTeF5 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, [TeF5 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(OTeF5 ) 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 CH4

respectively ii) boron tris(teflate), B(OTeF5)3, which undergoes metathesis to

generate BF3 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 HgF2 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]

Homoleptic

B(OTeF5)3[13’14]

C(OTeF5)4[15>16>i7]

Teflate

Organometallic

[Cp2Ti(OTeF5)2][8]

[CpFe(OTeF5)2]t8]

High Valent

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[OSeF5 ] [ 1 0 >2 1 ’2 2 1 [N 0 2 ][0SeF5] [23) F(OSeF5)[26]

Na[OSeF5] [ia21'22] [POF2 ][OSeF5] [24) Cl(OSeF5 ) [ 2 6 -2 7 1

K[OSeF5 ] [ 1 0 ’2 1 ’2 2 1 As(OSeF5)3[25] Br(OSeF5)[26]

Rb[OSeF5 ] [ 2 1 ' 2 3 1 Br(OSeF5)3[26]

Cs[OSeF5] [21'23] Group 16 Rb[Br(OSeF5 ) 4 ] [ 2 3 -2 6 1

(OSeF5)2[26,27] F,I(OSeF5 )5 J 17'28]

( x = 1-5)

Group 14 F2 Se(OSeF5)2[34l

C(OSeF5)4[29] F S 0 2 -0 -S 0 2 -0SeF5[31] Group 18

CH^(OSeF5)4. ; 29! SF5 -OSeF5[35l36] Xe(OSeF5)2[27,37]

( x = 0 -4 )

o C 5 F5 -OSeF5[30] 0=Se(0SeF 5 ) 2 [ 2 3 1 FXe-OSeF5[37’38]

FCO-OSeF5[30] F5 Se-0-SeF5[35] F5 SeO-Xe-OTeF5I27'38]

CF3 CO-OSeF5[31]

(CH3 )3 Si-OSeF5t32>33]

Table 2.2. Seflate derivatives of the transition metals.

Compound Reference

[F(4-i)Ti( ° s eF5) J 39,40

ii i

[0=V (0SeF5)3] 39,40

[0 2 Cr(OSeF5)2] 39,40

[Hg(OSeF5)2] 10,21,27

structural -OSeF5 and -OTeF5 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]).

XeF2 + 2 HOSeF5 --------► Xe(OSeF5 ) 2 + 2 HF Eqn. 2 .2 .

The compounds Br(OSeF<;)^2̂ and [Hg(OSeF5)2]F1,21’27J have also been

prepared from their respective precursors, BrF3 and HgF2, by this method.

14

ii) The difficulty associated with using HOSeF5 can be avoided when using

[Hg(OSeF5)2], this reacts gently with chlorides (Eqn. 2.3).

C H ^C l^ + Vi Hg(OSeF5 ) 2 --------► CHx(OSeF5)4.x + V4HgCl2 Eqn. 2.3.

(x = 0 - 3 )

By this route [CrO2 (OSeF5)2],[39'40] [VO(OSeF5)3],[39'40] C(OSeF5)4[29] and

As(OSeF5)3 2̂9' have been prepared from C r0 2 Cl2, VOCl3, CC14 and AsC13

respectively.

iii) The use of Xe(OSeF5 ) 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).

U VF 5 SeOXeOSeF5 — — ► 2 OSeF5 -------- ► F5 SeOOSeF5 Eqn. 2.4.

” X G

F5 SeOOSeF5 A » F5 SeOSeF5 + SeF4 + SeF6 + 0 2 Eqn. 2.5.

F,SeOXeOSeF, ---- ► Xe + F5 SeOSeF5 Eqn. 2.6.130 C

Other reagents have found isolated uses in the preparation of seflate

species and these include F2 OPOSeF5 1̂7,28̂ and C10SeF5 2̂8 ̂ (Eqn.’s 2.7 and

2 . 8 respectively).

IF5 + F 2 OPOSeF5 --------► Fxl(OSeF5)5.x + POF3

(x = 0 - 5 )

IC13 + C!OSeF5 ► I(OSeF5 ) 3 + Cl2

Eqn. 2.7.

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(OSeF5 ) 2 Xe + (F5 SeO ) 2 Eqn. 2.9.

The elimination of selenium oxide tetrafluoridej39,40 ̂ 0=SeF4, from a

coordinated seflate group is encountered when the -OSeF5 group is attached to

very electropositive, or coordinatively unsaturated central atoms (Eqn. 2.10).

Ti(OSeF5 ) 4 -> FJCTi(OSeF5)4_;c + 0=SeF 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 F5 SeOOSeF5 -> F5 SeOSeF5 + 0 2 Eqn. 2.11.

2 F5 S 0 -0 -0 S F 5 -> F5 SOOSF5 + 0 2 Eqn. 2.12.

16

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(OSeF5 ) 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 NMR chemical shifts of the methyl protons of CH 3 X, X = I, Br, Cl, F

and OSeF5 2̂9J are presented in Table 2.3. By extrapolation of 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, CH2 (OSeF5)2.

Table 2.3. Proton-1 NMR chemical shiftsa for CH3X and CH 2 X2, X = halogen

or seflate.

CH 3 (OSeF5) 4.50 CH2 (OSeF5 ) 2 6.30

c h 3f 4.26 c h 2 f 2 5.54

c h 3c i 3.05 c h 2 c i 2 5.33

CH3Br 2 . 6 8 CH2 Br2 4.94

c h 3i 2.19 c h 2 i 2 3.90

a ppm relative to TMS.

17

ii) The reaction of IF5 with F2 OPOSeF5 leads to substitution of the fluorine

atoms on the iodine by seflate^18,28̂ (Eqn. 2.13).

IF5 + F2 OPOSeF5 -> FxI(OSeF5)5. ̂ + POF3 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).

IF3 + B(OTeF5 ) 3 -> I(OTeF5 ) 3 + BF3 Eqn. 2.14.

I(OTeF5 ) 3 + Xe(OTeF5 ) 2 -> I(OTeF5 ) 5 + Xe Eqn. 2.15.

iii) The ligand properties of -OSeF5 and -OTeF5 groups in the pseudo-trigonal-

bipyramidal molecules F 2 Se(OSeF5)2, F2 Se(OTeF5 ) 2 and F2 Te(OTeF5 ) 2 have

also been investigated^34 ̂ by 77Se and 125Te NMR spectroscopy. All three

compounds possess axial -OSeF5 or -OTeF5 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 POF2 -X, X = Cl, F or OSeF5 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 J41 ̂

The NMR 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 p p m ^ 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 Mossbauer 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.

Of the studies carried out, those performed by Schrobilgen appear to be

the most definitive, and indicate that fluorine possesses a greater

19

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 su 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(OSeF5 ) 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,

[IO2 F4 ]', 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.

Whilst 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 f in 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 1 9 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 Fax due to deshielding of the axial fluorine.

2.5.1. Fluorine-19 NMR spectroscopy.

Of 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 AX4 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 AB4 spectra. This is the case for seflic a c i d ^

where 5Fa 75.9 and 8 Fe 66.1 ppm. For F5 SeO-OSeF5,[27] the A and B4 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

CF3 CO-OSeF5 ,̂ 31J 5Fa 61.2 and 5Fe 73.5 ppm.

The appearance of the spectrum depends on the ratio, R ,[49,50] Qf the 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 215-

240 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.

Within 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 SFa 5 F e J ( FaFe) A ab Solvent Ref.

ppm ppm Hz ppm

[N 0 2 ][OSeF5] 108.9 72.1 224 -36.8 c h 3c n 23

IOSeF5 92.1 74.0 229 -18 Neat 26

Xe(OSeF5 ) 2 80.5 69.4 234 - 1 1 . 1 c f c i 3 27,37

FOSeF5 54.2 52.1 230 -2 . 1 Neat 26

F5 SeOOSeF5 55.2 54.1 230 -0 . 8 Neat 27,26

F5 SeOCOCF3 61.2 73.5 211 12.3 Neat 31

F 5 SeOSeF5 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 2 1 Neat 31

23

Table 2.5. Solvent effects on the value of R for [Ti(OTeF5)4] .

Solvent §Fa SFe J (FaFe) R

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

c h 3 n o 2 -40.3 -48.9 187 0.38

Genetron 113 -47.3 -45.2 187 -1.58

CC14 -49.5 -43.3 183 -0.52

2.5.2. Vibrational spectroscopy.

The seflate group, [OSeF5]", possesses C4v symmetry for which the

following vibrational representation is obtained;

rV 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.

Se,

o -

F \ / F .S e ,

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 ][0SeF5] and F5 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 Ref.

[N 0 2 ][0SeF5] 918 23~~

Xe(OSeF5 ) 2 787 27,37

F5 SeOOSeF5 765 27,36

F5 SeOSeF5 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]^11,53̂ 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

point group

Description of vibration

V i( A j) v (Se - 0 )

v 2 ( A i ) ^ sy m ( ^ 6 - Feq)

v 3 ( A i ) v (Se - Fax)

v4 ( A j) 8 sym (out-of-plane SeF4)

v 5 ( B ! ) vsym (out-of-phase SeF4)

V6 ( B ! ) 8a sy m (out-of-plane SeF4)

v 7 ( B 2 ) 8 sym (in-plane-SeF4)

v 8 ( E ) ^ a sy m (ScF4)

v 9 ( E ) 8 (F - Se - F4)

v i o ( E ) 8 (O - Se - F4)

V n ( E ) S asym (in-plane SeF4)

2.5.3. Mass 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=SeF4, leaving one fluorine behind, is also

common. This was what was found for Br(OSeF5 ) 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=SeF 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,

(<103 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.

While 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(OSeF5 ) 5 and

Xe(OSeF5 ) 2 have few analogues outside of fluorine chemistry.

Using electron diffraction a structural investigation was carried out on

bis(pentafluoroselenium) oxide, F5 SeOSeF5 J35,36 ̂ The structure consists of

octahedra linked via an oxide bridge (Figure 2.3).

The gas phase structure of F5 SeOSeF5 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.

Figure 2.3. The gas phase structure of F5 SeOSeF5.

28

The bridge angle is large and constant (about 143°) for the three

chalcogen species F5 SOSF5, F5 SeOSeF5 and F5 TeOTeF5. This is at variance

with the fact that steric interactions between the equatorial fluorines diminishes

considerably in the sequence F5 SOSF5 > F5 SeOSeF5 > F 5 TeOTeF5.

Characterisation of F5 SOSF5 3̂5,36 ̂ shows the equatorial fluorines are bent 2.1°

away from the octahedral orientation, but this effect diminishes in F 5 SeOSeF5

(1.1°) and disappears for F5 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) A for F5 SeOSeF5, is between that

of a double and a single bond value (Se0 2 (Se=0 ) = 1.61(1) A and ethylene

selenite, (Se-O) = 1.80(2) A)J35,361 Therefore, 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 F5)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, F5 SOSF5 and F5 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 anionJ5 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 oxygen-

fluorine 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 HOTeF5 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 OTeF5 group with respect to the

Mn(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 ][0SeF 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(OTeF5 ) 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).

XeF2 + 2H O SeF 5 -» Xe(OSeF5 ) 2 + 2H F Eqn. 2.18.

This involves the use of seflic acid, HOSeF5 which is both difficult to

prepare and handle. The synthesis of seflic acid^1,27,48̂ 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 e0 2 F 2 and HOSeF5 from the involatile H2 S e04.

3 S e0 2 F2 + 4 HF H2 Se0 4 + 2 HOSeF5 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(OSeF5 ) 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 NMR 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).

O-Xe-O’ 180a F(2)-Se-F(3) 92(3)

F (l)-S e-F (l’) 8 8 (8 ) F(2)-Se-0 170(2)

F(l)-Se-F(2) 88(3) Xe-O-Se 125(2)

F(l)-Se-F(3)

X)0

00

r—H 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 5129Xe 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 e0 2 + SF4 —> SeOF2 + SOF2 Eqn. 2.21.

S e0 2 + 2 SF4 —) SeF4 + 2 SOF2 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 e0 2 is used, SeOF2 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 SF4 was condensed on to S e0 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 SeOF2 was then condensed on to the XeF2 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(OSeF5 ) 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 AB4 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) cm '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 k3 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, 6 fluorine atoms at 1.69(1) A and a xenon

atom at 3.07(1) A. 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 X-ray EXAFSe

Mean rf(Se-X) / Af 1.67(5)g

1.73(5)h

1.69(1)

2o2/ Ab - 0.008(2)

d(Se-Xe) / A 3.24(2) 3.07(1)

2o2/Ab - 0.014(2)

Fit index0 2.5

Rd 0.064 19.1

a Standard deviations in parentheses. b Debye-Waller factor. c Fit index = Xi[(%

T - X E ) k i 3 ] 2 . d R = [s(%T-%E)k3 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(OSeF5 ) 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

of 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

Xco-5

-15

* -1k /A ‘lb]

3.0

C 2.0 D>.> -(0k .13< 1.0

0 - 1 r /A

aEXAFS ( experimental x k3, curved-wave theory x k3)

bFourier transforms ( experimental, — theoretical)

37

In comparison, a microwave spectroscopic study of SeOF2 6̂1 ̂

determined the Se= 0 distance to be 1.576(3) A, and in 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 ][0SeF5] Aab = -36.8 and F5 SeOSeF5 Aab = 13.3. Covalent interactions

in the case of F5 XOXF5 (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) A, whereas, the value obtained using

EXAFS spectroscopy was 3.07(1) A. 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 1.6-

1.69 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(OSeF5 )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, XeF2 in solution is a

mild fluorinating agent and reaction of xenon difluoride, XeF2, with

[Re2 (CO)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.

[Re2 (CO)10] + 3 XeF2 -> [Re(CO)5F ReF5] + 5 CO + 3 Xe Eqn. 2.23.

The reaction between [Re2 (CO)10] and Xe(OTeF5 ) 2 also leads to a low-

valent 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.

[Re2 (CO)10] + Xe(OTeF5 ) 2 2 [Re(CO)5 (OTeF5)] + Xe Eqn. 2.24.

In an attempt to prepare a Re (I) seflate derivative the reaction between

[Re2 (CO)10] and Xe(OSeF5 ) 2 was investigated. Colourless [Re2 (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 NMR 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 AX4 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 NMR 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, 6 8 6

s, 6 6 6 sh, 592 s, 555 s, 505 w and 492 s c n r1.

The high v(Se-O) of 856 cm " 1 is exceeded only by those of

[N 0 2 ][0S eF 5] and the alkali metal salts. This compliments the 19F NMR 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 [Re^CO)!0] and Xe(OSeF5 ) 2

182.5 181.5182.0 181.0 180.5 180.0 179.5 179.0 178.5 178.0

(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 8 2 Se) (Figure 2.9). Accurate mass

spectrometry was used to 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(OSeF5 ) 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

l(X) —

0 —

518

455

460 480 500 520 540 580 6(X)

12 13 16 18 17 19 187 185 80 78 82Mass Abundance C C 0 0 0 F Re Re Se Se Se517.83400 29.3811 5 0 6 0 0 5 1 0 1 0 0515.83100 17.6287 5 0 6 0 0 5 0 1 1 0 0515.83480 13.8708 5 0 6 0 0 5 1 0 0 1 0513.83180 8.3225 5 0 6 0 0 5 0 1 0 1 0519.83420 5.4198 5 0 6 0 0 5 1 0 0 0 1513.83670 5.3195 5 0 6 0 0 5 1 0 0 0 0514.83740 4.4703 5 0 6 0 0 5 1 0 0 0 0517.83120 3.2519 5 0 6 0 0 5 0 1 0 0 1

Centroid10 0 -1

I nt 50 H

Llnowldth = 100.0 ppn lOOX = 330225

460 470 480 490 500Mass

510 520 530 540

517.83403lOOn

9 0-

80-

70-

s 60ium 50 -

4 0-

30^2 0 -

1 0 -

5 1 7 . 9 7 5 3 4

60 100 240120Channel

220140 160 180 200

43

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 6 fluorine

atoms at 1.71(1) A and 1 rhenium atom at 3.55(1) A (Table 2 . 1 0 and Figure

2.10). Each shell was tested for statistical significance J60]

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

[Re2 (CO)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

[Re2 (CO)10], [Re(CO)5 Cl] and [Re(CO)5 (OSeF5)].

Table 2.10. EXAFS data for [Re(CO)5 (OSeF5)].

Parameter EXAFSe

d(Se-X) / Af 1.71(1)

l a 2/ Ab 0.007(2)

rf(Se-Re) / A 3.55(1)

l a 2/ Ab 0 .0 1 2 (2 )

Fit indexc 4.1

Rd 25.7

a Standard deviations in parentheses. b Debye-Waller factor. c Fit index = Xi[(%

T-XE)ki3]2. d R = [s(%T-%E)k3 dk/s%Ek3 dk] x 100 e E0 5.0 (4), AFAC 0.86 and

VPI -4 .71 .f Mean bond length (X = F and O).

44


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)


45

The spectroscopic data presented in this section conclusively shows that

the reaction of [Re2 (CO)10] with Xe(OSeF5 ) 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(OSeF5 )2 -

A survey of the literature indicates that F2 and XeF2 do not react with

[Mn2 (CO)10] to produce any stable carbonyl fluorides. The reaction between

AgF and [Mn(CO)5 Br] was originally reported to yield the dimer

[{Mn(CO)4 F}2] or, with excess of AgF, [Mn(CO)3 F3].t65 ̂ However,

reinvestigation by Horn et. alS66̂ 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 Leicester^ has shown that [Mn2 (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 d5 Mn (II) species, believed to be cis-

[Mn(CO)4 (OTeF5)2]. Any further excess of Xe(OTeF5 ) 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] + HOTeF5 -> [Mn(CO)5 (OTeF5)] + CH 4 Eqn. 2.25.

The 19F NMR 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(OTeF5 ) 2 and [Mn2 (CO)10]. In view of these

results the synthesis of [Mn(CO)5 (OSeF5)] was attempted via the reaction

between [Mn2 (CO)10] and Xe(OSeF5)2.

Manganese carbonyl dimer, [Mn2 (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 J67 ̂ 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

[Mn2 (CO)10] and Xe(OSeF5)2.

208 198206 202 200

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 8 0 Se).

Attempts were made to record manganese K edge EXAFS data for

[Mn2 (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 A (k = photoelectron wave vector). This was

later truncated to 13.5 A due 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, 6 fluorine atoms at 1.70(1) A and 1 manganese atom at 3.38(1) A (Table 2.11 and Figure 2.11). Each shell was tested for statistical

significance

Table 2.11. EXAFS data for [Mn(CO)5 (OSeF5)].

Parameter EXAFSe

d(Sc-X) / Af 1.70(1)

2 c 2/ Ab 0.008(2)

</(Se-Mn) / A 3.38(1)

2 a 2/ Ab 0.008(2)

Fit index0 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)k3 dk/s%Ek3 dk] x 100 %. e E0 12.7 (4), AFAC 0.86

and VPI -4 .71 .f Mean bond length (X = F and O).

From this evidence it has been shown that xenon bis(seflate) does indeed

react with [Mn2 (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


spectra for [Mn(CO)5 (OSeF5)].

(a)

*CO

-5

-15

(b)

3.0

CO

cD>.i _C 6i —

6 82 104

r / A 1



52

2.12. The Reaction Between [Ru(CO)3(PPh3)2] and Xe(OSeF5)2.

As a result of the successes with similar systems, numerous reactions

have been undertaken between XeF2 and a range of ruthenium (0 )

complexes.[68,69̂ Of these reactions, that between [Ru(CO)3 (PPh3)2] and XeF2

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).

XeF2 + [Ru(CO)3 (PPh3)2] -> [RuF2 (CO)2 (PPh3)2] + Xe + CO Eqn. 2.26.

The products and intermediates formed in this reaction were

characterised using multinuclear NMR and infrared spectroscopy, and X-ray

crystallography. Although the reaction between [Ru(CO)3 (PPh3)2] and

Xe(OTeF5 ) 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 (PPh3 )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(OSeF5 ) 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

cm '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 AX4 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

[RuF2 (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 J69 ̂ The reported chemical shift of the metal-bound

fluorine atom in the 19F NMR data for [RuF2 (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.

100105

(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 22 21

(ppm)

56

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 distribution was detected for

[Ru(CO)(PPh3 )2 F]+ m/z 673 (for 1 0 2 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 I2 -> 5 Agl + IF5 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, IF5 is also accessible from I2, HI or I2 0 5 using a variety of

fluorinating agents.

57

The chemistry of IF5 has been quite extensively studied but, relevant to

the present work, in 1978, Seppelt et al. reported1 1 7 ,2 8 1 that IF5 and F2 PO-

OSeF5 undergo a metathetical reaction to produce I(OSeF5 ) 5 (Eqn. 2.28).

IF5 + F2 PO-OSeF5 -> FjCI(OSeF5)5.JC + POF3 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 IF5 using 19F

NMR spectroscopy.

Iodine tris(seflate), I(OSeF5)3, can be synthesised by the following

reaction (Eqn. 2.29).

3 Cl-OSeF5 + IC13 -> I(OSeF5 ) 3 + 3 Cl2 Eqn. 2.29.

Both I(OSeF5) and I(OSeF5 ) 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 unstable1 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).

The first reaction was intended to synthesis I(OSeF5 ) 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

2 IF3 -> IF + IF5 Eqn. 2.30.

Eqn. 2.31.5 IF -> 2 I2 + IF5

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 AX4 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).

8

ppm

2J (FaFe)

Hz

‘/(S e p ,)

Hz

9.5 (d)b 90 -

42.8 (s) - 862

56.3 (q)b 91 -

59.7 (d) 224 1384

6 6 . 8 (d) 231 1367

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.

7580 70 65 60

(ppm )


between I2 and five molar equivalents of Xe(OSeF5)2.

(ppm)


between I2 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(OSeF5 ) 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 IF5 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 AX4 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 AX4 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 AX4 patterns showed the

same basic features as those for the I2 and five molar equivalents of

Xe(OSeF5)2, namely I(OSeF5 ) 5 (Figure 2.17). However, the third AX4 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(OSeF5 )J23,26 ̂ which at 5Fe

92.0 ppm would be clearly visible.

The 19F NMR data reported^23,26] for I(OSeF5 ) 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(OTeF5 ) 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(OSeF5 ) 2 this signal disappeared and only resonances

assignable to I(OSeF5 ) 5 were observed. The presence of iodine in the reaction

mixture, coupled with the fact that no Xe(OSeF5 ) 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 I2 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)

Hz

lJ( SeFe)

Hz

9.5 (d) 90 -

42.2 (s) - 8 6 8

59.6 (d) 226 1383

66.7 (d)b 2 1 0 -

70.0 (d) 229 1343

73.5 (q)b 2 2 0 -

77.7 (q)a - -

83.3 (q) 227 -

1 L 1 ~~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,[3̂ Se2 0 2 F8,[77] S e0 2 F2[78̂

and SeF6 [ 7 9 1 all of which possess Se-F coupling constants in the region of

1500-1300 Hz. The 19F NMR spectrum of neat SeOF2 p ro 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 SeOF2 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 v(X -0)a /c m _1 Reference

[Re2 (CO)10] 2070, 2013, 1975 - b

[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

[Mn2 (CO)10] 2046, 2013, 1983 - b

[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

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 (PPh3 )2 ] may produce [Ru(CO)2 (PPh3 )2 (OTeF5)2], the spectroscopic

evidence is not conclusive. The reaction between xenon bis(seflate) and

[Ru(CO)3 (PPh3 )2] produced a seflate containing species, however, several

species were generated and identification of the products was difficult. The

only product identified was [RuF2 (CO)2 (PPh3 )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

[Mn2 (CO)io] XeF2 No fluoro products -

Xe(OTeF5 ) 2 [Mn(CO)5 (OTeF5)]

4

[Mn(CO)4 (OTeF5)2]

4,

Decomposition

8

Xe(OSeF5 ) 2 [Mn(CO)5 (OSeF5)]c

[Re2 (CO)10] XeF2 [Re(CO)5 .JuF.ReF5]

[Re(CO)6 ][ReF6]

64

Xe(OTeF5 ) 2 [Re(CO)5 (OTeF5)] 8

Xe(OSeF5 ) 2 [Re(CO)5 (OSeF5)] c

[Ru(CO)3 (PPh3)] XeF2 [OC-6-13] [RuF2 (CO)2 (PPh3)2] 69

Xe(OTeF5 ) 2 [Ru(CO)2 (PPh3 )2 (OTeF5)2]a 8

Xe(OSeF5 ) 2 [RuF2 (CO)2 (PPh3)2]bc

h XeF2 i f 5 -

Xe(OTeF5 ) 2 I(OTeF5 ) 5 17,28

Xe(OSeF5 ) 2 I(OSeF5 ) 5c

a Postulated.b The only one of several products which could be identified.

c This work.

68

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 manganese-

carbonyl 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.



[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., 1994,116, 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. Engl, 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.


[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.


[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


[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., 1978,17, 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.


[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 ] Parameter 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.



[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., 1976,113, 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.


[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.


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 theoryf1,2̂ (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 (m+n) Arrangement Approximate Symmetry

bonds lone pairs of bonds and shape

(m) (n) lone pairs_____ _ _ _ Pseudo “

trigonal

Br

o^llN

Br

c/'INF F

Br03F 4 0 4 Pseudo F

tetrahedral

O

3ha[Br03]+ 3 0 3 Trigonal jj D

/ Xo o

Br02F 3 1 4 Pseudo Cs

tetrahedral I

O

[BrOF2]+ 3 1 4 Pseudo .. Cs

tetrahedral

'3v

76

a[Br02F2]+ 4 0 4 Pseudo

tetrahedral

F

Br

o ^ l l \0 F

c2v

a[Br03F2]- 5 0 5 Pseudo

trigonal

bipyramidal

F

V lBr = 0

0^1F

D 3h

aBr02F3 5 0 5 Pseudo

trigonal

bipyramidal

F

v j. B r ----- F

o ^ |F

C2v

a[BrOF4]+ 5 0 5 Pseudo

trigonal

bipyramidal

F

' s L op / i

F

c2v

BrOF3 4 1 5 Pseudo

trigonal

bipyramidal

F

v jB r ----- 1

f/ |F

Cs

[Br02F2]- 4 1 5 Pseudo

trigonal

bipyramidal

F

J. B r ----- :

IF

C2v

[BrOF4]- 5 1 6 Pseudo

octahedral

0F\ II_F

Rr--C4v

l[Br02F4]

aBrOF<

Pseudo

octahedral

Pseudo

octahedral

OII

Br:

OO

D4h

Br:

F

-F'F

-F

‘F

'Aw

Species are unknown to date.

77

3.3. The Halogenyl Fluorides, XO2F.

3.3.1. Chloryl fluoride.

Chloryl fluoride, C102 F, has been prepared by several methods. Early

syntheses used shock-sensitive chlorine oxides such as C120 and C102 J 3,4̂ 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 C102 F.[5,6]

Chloryl fluoride, a colourless gas at room temperature, is a powerful

oxidising and fluorinating reagent which is stab le^ up to 300°C. Studies using

microwave spectroscopy have provided information on the molecule's internal

param eters'8,9 ̂ and shown that it has the predicted Cs symmetry.

The reaction of C102F with dried CsF at -80°C affords

Cs[C102 F2],^10,11̂ and analysis by infrared spectroscopy^ produces data

consistent with the anion being a pseudo-trigonal bipyramid of C2v symmetry.

Chloryl fluorides react with the Lewis acids BF3, PF5, AsF5, SbF5 and VF5^ to

form salts comprising of the cation [C102]+ and the anions [BF4]', [PF6]‘,

[AsF6]', [SbF6]' and [VF6]' respectively. The reaction with PtF6 gives a

mixture of [C102 F2]+[ PtF6]‘ and [C102 ]+[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 BrF5 and K [Br03] 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], BrF3 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 17̂ 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, Br2 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 AsF5 and SbF5 J 21,22 ̂ which are, however, of low stability

and decompose near room temperature. The reaction of PtF6 and B r0 2F at

-120°C yields a mixture of [BrOF2 ]+[PtF6]' and [B r0 2 ]+[PtF6] '.[23] The

attempted oxidative fluorination of B r0 2F using KrF2 does not yield bromine

(VII) oxide fluorides, but instead, proceeds via BrOF3, to yield BrF5 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 CH3 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 [Br0 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, IOF3 is formed by refluxing I 0 2F with IF5. Iodyl fluoride is also

produced by the fluorination of I2 0 5 in AHF, at 20°CJ3̂ Iodyl fluoride is a

stable colourless solid at room temperature, which slowly evolves HF in moist

79

air.t16̂ 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 salts such as

[I0 2 ]+[AsF6]" 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 [C102 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 CIOF3 by the UV irradiation of a mixture of C1F3 and

OF2 J 28] A large scale preparation^29̂ involves the fluorination of chlorine

nitrate, C 10N 02, at -35°C, which affords a mixture of C10F3 and F N 02. 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,[30̂ 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

CIF5 , and reacts with glass, quartz and most metals causing both fluorination

and oxygenationP 1̂ 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 N2 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, PF5 and BF3. The

vibrational spectra observed for the [C10F2]+ salts[33] showed the presence of

six fundamental vibrations, which is consistent with them being pseudo-

tetrahedral 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 [C10F4 ] 'J 32̂ 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 CIOF3 with SbF5 -F2 or PtF6 J 14,34̂ These failed but the

latter reaction produced [C10F2 ]+[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.[17̂ The HF is removed at

low temperature to leave K[HF2] and BrOF3. The BrOF3 cannot simply be

distilled from the BrOF3 -K[HF2] mixture since the reaction is reversible.

Instead, BrF5, into which the BrOF3 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 BrOF3 which

involves the reaction of L i[N 03] and an excess of BrF5 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 CIOF3 . The reaction between BrOF3 and the Lewis acids BF3, AsF5 and

SbF5 affords1 3 9 1 the adducts [BrOF2 ]+[BF4]-, [BrOF2 ]+[AsF6]’ and

[BrOF2 ]+[SbF6] ' respectively. The stability of the complex formed increases

considerably with the increasing strength of the Lewis acid employed. The

adduct [BrOF2 ]+[SbF6]" can also be prepared from the reaction between

I 0 2 F3 *SbF5 and BrF5. The vibrational spectroscopic data from these

complexes^23’37,39,40] are consistent with their containing a cation of pseudo-

pyramidal geometry (Cs symmetry) and the assignments made are in good

agreement with those for the isostructural species SeOF2, SOF2 and [C10F2]+.

Salts of the anion [BrOF4]" are easily synthesised using the method

described by Wilson and ChristeJ36̂ The reaction between BrF5 and the alkali

metal nitrates M [N 03], M = Na, K, Rb or Cs, yields the corresponding anionic

salts. From Na+ —» Cs+, smaller excesses of BrF5 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 ed ^ when

IOF3 was prepared by refluxing a saturated solution of I2 0 5 in IF5. On cooling,

colourless needles of IOF3 are formed. Iodine oxide trifluoride is stable up to

110°C at which temperature it decom poses^ to IF5 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 CIOF3 .

The anion [IOF4]‘ is accessible from the indirect reaction of KF and

I2 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 [I03]. X-ray crystallography^42! has shown that

[IOF4 ] ' 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, XO3F.

3.5.1. Perchloryl fluoride.

Perchloryl fluoride, C103 F, was first synthesised in the early 1950's and

has been extensively investigated since. Perchloryl fluoride is readily prepared

by several different rou 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 NaC104 in AHF also yields CIO3 F.

Perchloryl fluoride is a stable colourless gas with a melting point of -47

°C. The physical properties are well documented/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 CIO3 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 symmetry/44!

Applications include its selective fluorinating properties in organic

chemistry/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 C6 H5Li and C103F produces C6 H5 C103 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

CIO3 F and CIF3 , CIF5 , OF2 or F2 produces C10F3 J47 ̂ Perchloryl fluoride

behaves as a mild fluorinating agent and converts UF4 to UF6 via an unknown

uranium oxide fluoride,f48̂ 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, [BrF4 ]+[AsF6] ' or BrF5 on [KBr04] in

anhydrous HFJ49,50 ̂Perbromyl fluoride is a colourless gas with a melting point

of -110 °C. Electron diffraction studies confirm the structure is pseudo-

tetrahedral and vibrational studies are in agreement with the molecule

having C3v symmetry.

The chemical behaviour of B r0 3F is similar to that of CIO3 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 [I04] 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 AsF5 or BF3 to

give compounds containing the cation [I0 3 ]+.

84

3.6. Halogen Dioxide Trifluoride, XO2F3 .

3.6.1. Chlorine dioxide trifluoride.

Chlorine dioxide trifluoride is prepared by a multi-step synthesis.1[54̂ The

oxidation of C102F using PtF6 produces [C102 F2 ]+[PtF6]" and [C102 ]+[PtF6] \

The reaction of these cations with F N 0 2 or FNO, at -78°C, produces C102 F3

and C102F respectively. The C102 F3 is then separated from the C102F by

fractional condensation. Any remaining C102F can be removed by the addition

of BF3 to the mixture, and this produces the adducts [C102 F2 ]+[BF4]' and

[C102 ]+[BF4]". The species [C102 ]+[BF4]‘ is unstable above 20°C, and can be

removed as it is the only volatile product. The final step involves the reaction of

[C102 F2 +][BF4]" with F N 0 2 which liberates the volatile C102 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 C102 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 i o 2 f .

The synthesis of [C102 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 [C102 F2]+ to the spectrum of the isostructural SQ2 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 BrF5 in HF at low temperatures.[56]

Mass spectrometry of the hydrolysis products of BrF5 and BrF3 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 03F was allowed to react with [Ba3 H4 (I0 6)2] . The

mixture was then distilled under reduced pressure, and the fraction obtained

contained HOIOF4 and H S 03F in a 2:1 molar ratio. The addition of oleum

converted the HOIOF4 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 id ^ with a vapour

pressure of 5 torr at 25°C. It attacks glass and quartz slowly at room

temperature and is photosensitive, producing IOF3 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 F3 and AHF

which produces the acid HOIOF4. Alternatively, C s[I0 2 F4] can be prepared by

the reaction of C s[I04] with either AHF, BrF5, C1F3, C1F5 or F2 J62 ̂

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 OIOF4 group has been extended to xenon (II) derivatives,

compounds such as Xe(OIOF4 ) 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 AsF5 and

SbF5, to produce 1:1 adductsJ63 ̂ The spectroscopic data for these adducts are

not consistent with the ionic formulations [I0 2 F2 ]+[MF6]" (M = As or Sb). It

appears that I 0 2 F3 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, CIOF5 , has been reported as a product

from the photochemical reaction of C1F5 and OF2 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 BrF5 and excess of OF2 between -60 and -40°C j14̂

ii) the heating of a mixture of BrF5 and 0 2 at 4300 psi to 207°C,[37] iii) the

hydrolysis of [BrF6 ]+[AsF6]+ in HF^37̂ and iv) the reaction of BrOF3 and

KrF2 J 19̂ However, BrOF5 still remains unknown.

3.7.3. Iodine oxide pentafluoride.

Iodine oxide pentafluoride, IOF5, is formed by the reaction of IF7 and

water, silica or I2 0 5 J65"68 ̂ Another route, proposed by Christe and SchackJ69 ̂

employs the use of the ligand transfer reagent POF3, which reacts with IF7 to

produce IOF5 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 AsF5 and SbF5 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 IOF5 via the oxygen atom. Iodine oxide

pentafluoride readily reacts with the Lewis base [NMe4]+F to form the adduct

[IOF6 ]'[NM e4 ]+J72,73 ̂ 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, Na3 H2 [I06], 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 Se0 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[Br03] . The yield of this

reaction was rather poor at 1%. A more successful method involves the use of

aqueous XeF2 which produces [Br04]" in a 10% yield (Eqn. 3.2).

[BrO3] ' + H20 + XeF2 -» Xe + [Br04]’ + 2 HF Eqn. 3.2.

90

The most practical synthesis involves the oxidation of [B r03] ' using

elemental fluorine in aqueous sodium hydroxide (Eqn. 3.3). This preparation^80̂

involves a hazardous fluorination stage and a rather lengthy purification. The

yields obtained are around 2 0 %.

[B r03]" + F2 + 2 OH" -> [Br04]" + 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,^80̂ 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[82̂

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 [Br04] ' 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-SbF5 J 83̂ 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 3 x o 3 h 3 x 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 orbitalJ87̂ 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 BrF5 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 K[BrF4] J 9 0 1 Cs[BrF6] J 9 1 1

[BrF2 ][AsF6 ] [ 9 2 , 9 3 1 and [BrF4 ][Sb2 F n ] [ 9 4 , 9 5 1 were prepared as described in the

experimental. The adduct [BrF2 ][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 ̂

[BrF2 ][AsF6 ] [ 9 5 1 an(j [BrF4 ][Sb2 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], [BrF2 ][AsF6] and [BrF4 ][Sb2 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

transmission mode, out to k = 15 A'1 (k = photoelectron wave vector). This was

later truncated to 13.5 A - 1 for [BrF4 ][Sb2 F n ] and 13 A"1 for [BrF2 ][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

95

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 J97^

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],

[BrF2 ][AsF6] and [BrF4 ][Sb2 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],

[BrF2 ][AsF6] and [BrF4 ][Sb2 Fn ].

Compound K[BrF4] Cs[BrF6]

Crystal EXAFSe Crystal EXAFSf

</(Br-F) \ A 1.89(2) 1.88(1) 1.847(1) 1.85(1)

2a2 \ Ab 0.006(2) 0.007(2)

Fit Index0 2 . 8 2 . 1

Rd 23.4 20.3

Compound [BrF2 ][AsF6] [BrF4 ][Sb2 F n ]

Crystal EXAFSg Crystal EXAFSh

(/(Br-F) \ A 1.69(2) 1.70(1) 1.81(11) 1.69(1)

2 o 2 \ Ab 0.003(2) 0.007(1)

(/(Br-F) \ A 2.29(2) 2.35(1) 2.36(10) 2.40(1)

2 <J2 \ Ab 0.024(2) 0.028(3)

Fit Index0 2 . 0 2 . 0

Rd 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/sXEk3 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

-12° - i

k / A 1(b)

1.8

CO

k.< 0.6

4 62 8 10o _

r /A



98


spectra for Cs[BrF6].

(a)

-10

(b) 3.0

(0

C 2.0 D>.w(0

3k -

r / A



99


spectra for [BrF2][AsF6].

(a)

6

0

6

0k / A 1

(b)

1.0c3

(01 -4-»3 0.5<

v - ■

r /A

aEXAFS (----- experimental x k3, — curved-wave theory x k3)


100


spectra for[BrF4] [Sb2F x x ].

C O

-12

(b)

2.00)

1.0

<

4 6 8 102



101

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 sign ifican 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 8 1 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^102̂

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).

[BrF6 ]+[AsF6]' + NOF -> BrF5 + F2 + [NO]+[AsF6]’ Eqn. 3.4.

[IF6 ]+[AsF6]" + NOF -> IF7 + [NO]+[AsF6r 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 Br-

F 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 BrF7 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

GillespieJ17̂ Bougon obtained K[BrOF4] by the reaction of K [B r03] with a

large excess of BrF5 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 [Br0 2 F2] Eqn. 3.6.

In 1978, Christe reported an improved synthesis of [BrOF4]‘ salts

The reaction of K [Br04] or Cs[Br04] with BrF5 and F2 led to K[BrOF4] and

the previously unknown Cs[BrOF4]. Although this method results in essentially

pure products in high yield, the required [Br04]" salts are difficult to prepare.

The reaction of the caesium salt, Cs[Br04], 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 BrF5 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 02. 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 ]-BrF5 system at 0°C some BrOF3 was obtained along

with Na[BrOF4]. At 25°C any BrOF3 formed undergoes either fast

decomposition to BrF3 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[Br0 2 F2]. It was noted

that the formation of BrOF3 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, FN 0 2 or

M[BrOF4] and F N 02. 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

[BrF4 0 N 0 2] is presumably unstable and would eliminate FN 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 BrF5 and Cs[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 k3 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 charactero

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 [Br04](S),

N a[B r03](s), N a[Br02](s) and Na[BrO](aq),̂ 103̂ 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)

2c2 / Ab 0.007(1)

rf(Br-F) / A 1.87(1)

2a2 / Ab 0.006(1)

Fit Index0 2.29

Rd 18.0

a Standard deviations in parentheses. b Debye-Waller factor. c Fit index = Xj[(%

T-XE)ki3]2. dR = [s(XT-%E)k3 dk/s%Ek 3 dk] x 100 %. e E0 7.25 (0.40), AFAC

0.71 and V P I-1.86.

107


spectra for Cs[BrOF4].

M

-6

o -k / A 1

1.8

Z> 12

< 0.6

62 84 10

r /V


bFourier transforms ( experimental, —- theoretical)

108

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 + F2 + 2 OH* [Br04]' + 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 r04] and referenced to aqueous KBr (1 mol dm-3), the values being

corrected to infinite dilution using the published dataJ103̂ The 81Br NMR

resonance of [B r04]" was observed at 52470, relative to infinitely dilute

aqueous Br", in good agreement with the reported^103̂ 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 [Br04] and SbF5,

SbF5 -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 BrF5 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 BrF5 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 r03F decomposes slowly at room temperature to

produce Br2. The Br2 can be readily removed by the condensation of F2 into the

FEP tube at 196°C. On warming to room temperature the Br2 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[Br04] and BrF5, using AHF as the solvent has

been previously investigated^21̂ 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 AX4 pattern generated by the BrF5 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 r03F (cf. neat B r0 3 F, 8274 ppm, -80

°C)J3TI

2 K [B r04] + BrF5 + 2 HF -> 2 B r0 3F + B r02F + 2 K[HF2] Eqn. 3.8.

The presence of B r02F was used by Spekkens^2^ to explain the poor

resolution observed for the 19F NMR spectrum of the reaction media. Fluorine-

19 NMR 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 r02F to be 8210 ppm when

recorded as a solution in BrF5 J37̂ The presence of B r02F 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 [Br04] with XF5 -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 r04] + AsF5 —» B r0 3F + K[AsOF4] Eqn. 3.9.

I l l

Figure 3.8. Fluorine-19 NMR spectrum of B r0 3F in BrF5.

t o

276 274 272

(ppm )270 268 140 138 136 134

(ppm )132 130

Raman spectroscopy has shown that neither AsOF3 nor K[AsOF4] are

present^21! An alternative mechanism is that proposed for the formation of

CIO3 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 XF5 -> 2 [H2 F]+ + 2 [XF6]‘

2 [H2 F]+ + [Br04r -> [H2 O Br03]+ + 2 HF

[H2 O B r03]+ + HF -» B r03F + [H3 0 ] +

overall: [Br04] ' + 3 HF + 2X F 5 -> B r03F + [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-OH2 bond in [H2 0 B r0 3]+ to give H20 and [Br03]+ followed by reaction

of [B r03]+ 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 3̂6̂ did not produce

Na[BrOF4] as the only product. They observed that BrOF3 was also formed,

along with some Na[BrF4], which was presumably formed via the

decomposition of the BrOF3 as shown in Equations 3.10 and 3.11 (Figure 3.6).

BrOF3 —» BrF3 + V i0 2 Eqn. 3.10.

NaF + BrF3 -> 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 BrOF3 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 BrF5 at 0°C over twenty days, BrOF3 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, BrF5 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 BrOF3 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

acidJ17̂ 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 BrOF3 and K[HF2] are formed (Eqn. 3.12).

K[BrOF4] + HF —» BrOF3 + K[HF2] Eqn. 3.12.

The HF can be removed by pumping the mixture at -60°C to leave a

mixture of BrOF3 and K[HF2]. The BrOF3 cannot simply be removed under

dynamic vacuum as the reaction is reversible. This problem can be overcome

by condensation of BrF5 on to the mixture which solubilises the BrOF3 but not

the K[HF2] . Decantation of the solution affords separation. The BrF5 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 r03] and K[BrF6] in CH3 C N J17̂ 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 [Br0 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 BrF5 and a

catalytic amount of HF reportedly produces K [Br0 2 F2], K[BrOF4] and

B r0 2 F J 15,16,17̂ 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 r02F should have

condensed. No material was collected in the trap indicating that B r02F 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 [Br0 2 F2] and HF yields B r0 2F

and K[HF2].[17] Unlike the K[HF2 ]-BrOF3 mixture, B r02F can be removed

from the K[HF2 ]-B r02F mixture without the reverse reaction occurring;

indicating that BrOF3 is a stronger fluoride ion acceptor than B r0 2 F. As a

consequence of not being able to separate the mixture of K [Br0 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], [BrF2 ][AsF6] and

[BrF4 ][Sb2 F 11]. The expected trends in Br-F bond lengths have been clearly

illustrated and the improvement in the data available for [BrF4 ][Sb2 Fn ]

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 [Br04] 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. J. Gillespie, Molecular Geometry, Van Nostrand Reinhold, London,

1972.

[2] R. J. Gillespie, Chem. Soc. Rev., 1992, 59.

[3] A. J. Downs and C. J. Adams, Comprehensive Inorganic Chemistry,

Pergamon Press, Oxford, 1973, 2, 1386-96 and references cited therein.

[4] Y. Macheteau and J. Gillardeau, B ull Soc. Chem. Fr., 1967,11, 4075.

[5] K. O. Christe, R. D. Wilson and C. J. Schack, Inorg. Nucl. Chem. Lett.,

1975,11, 161.

[6 ] K. O. Christe and C. J. Schack, Adv. Inorg. Chem. Radiochem., 1976,

18, 319 and references cited therein.

[7] W. H. Basualdo and H. J. Schumacher, Angew. Chem., 1955, 67, 231.

[8 ] C. R. Parent and M. C. L. Gerry, J. Chem. Soc., Chem. Commun., 1972,

285.

[9] C. R. Parent and M. C. L. Gerry, J. Mol. Spectrosc., 1974, 49, 343.

[10] D. K. Huggins and W. B. Fox, US. Pat., 3 423 168, 1969.

[11] D. K. Huggins and W. B. Fox, Inorg. Nucl. Chem. Lett., 1970, 6 , 337.

[12] K. O. Christe and E. C. Curtis, Inorg. Chem., 1972,11, 35.

[13] K. O. Christe, US. Pat., 3 879 526, 1975.

[14] K. O. Christe, Inorg. Chem., 1973,12, 1580.

[15] M. Schmeisser and K. Brandle, Adv. Inorg. Chem. Radiochem., 1963, 5,

411 and references cited therein.

[16] R. Bougon and G. Tantot, C. R. Seances Acad. Sci. Paris, 1975, C 281,

271.

[17] R. J. Gillespie and P. Spekkens, J. Chem. Soc. Dalton Trans., 1976,

2391.

[18] R. J. Gillespie and P. Spekkens, J. Chem. Soc., Chem. Commun., 1975,

314.

119

[19] R. J. Gillespie and P. Spekkens, J. Chem. Soc. Dalton Trans., 1977,

1539.

[20] K. O. Christe, E. C. Curtis and E. Jacob, Inorg. Chem., 1978,17, 2744.

[21] P. Spekkens, Ph.D. Thesis, McMaster University, Canada, 1977.


[23] M. Aldehelm and E. Jacob, Angew. Chem. Int. Ed. Engl, 1977,16, 461.

[24] R. Bougon, P. Joubert and G. Tantot, J. Chem. Phys., 1977, 6 6 ,1562.

[25] R. E. Aynsley, R. Nichols and P. L. Robinson, J. Chem. Soc., 1953, 623


[26] H. A. Carter and F. Aubke, Inorg. Chem., 1971,10, 2296.

[27] D. Pilipovic, C. B. Lindahl, C. J. Schack, R. D. Wilson and K. O.

Christe, Inorg. Chem., 1972,11, 2189.

[28] R. Bougon, J. Isabey and P. Plurien, Compt. Rend., 1970, 271C, 1366.

[29] D. Pilipovich and C. J. Schack, US. Pat., 3 692 476, 1972.

[30] H. Oberhammer and K. O. Christe, Inorg. Chem., 1982, 21, 273.


[32] K. O. Christe, C. J. Schack and D. Pilipovich, Inorg. Chem., 1972,11,

2205.


[34] K. O. Christe, C. J. Schack and D. Pilipovich, Inorg. Chem., 1972,11,

2205.

[35] R. Bougon and T. Bui Huy, C. R. Seances Acad. Sci., Ser. C, 1976, 283,

461.

[36] W. W. Wilson and K. O. Christe, Inorg. Chem., 1987, 26, 916.

[37] R. J. Gillespie and P. H. Spekkens, Isr. J. Chem., 1978,17, 11.

[38] K. O. Christe, E. C. Curtis and R. Bougon, Inorg. Chem., 1978,17,

1533.

[39] R. Bougon, T. Bui Huy, P. Charpin, R. J. Gillespie and P. H. Spekkens,

J. Chem. Soc., Dalton Trans., 1979, 6 .

[40] N. Keller and G. J. Schrobilgen, Inorg. Chem., 1981, 20, 2118.

120

[41] J. W. Viers and H. W. Braid, J. Chem. Soc., Chem. Commun., 1967,

1093.

[42] K. O. Christe, R. D. Wilson, E. C. Curtis, W. Kuhlmann and W.

Sawodny, Inorg. Chem., 1978,17, 533.

[43] Gall J. F., Kirk Othmer Encyclopedia o f Chemical Technology, Wiley,

London, 2nd edn., 1966, vol. 9, 598.

[44] A. H. Clark, B. Beagley and D. W. J. Cruickshank, J. Chem. Soc.,

Chem. Commun., 1968,14.

[45] W. A. Sheppard, Tetrahedron Lett., 1969, 83.

[46] F. A. Cotton and G. Wilkinson, Advanced Inorganic Chemistry, Wiley,

New York, 4th edn., 1980, 570.

[47] D. Pilipovich, H. H. Rogers and R. D. Wilson, Inorg. Chem., 1972,11,

2192.

[48] H. Rude, R. Benoit and O. Hartmanshenn, Nucl. Sci. Abstr., 1972, 26,

6528.

[49] E. H. Appleman, J. Amer. Chem. Soc., 1968, 90, 1900.

[50] E. H. Appleman, Inorg. Synth., 1973,14, 29.

[51] E. H. Appleman, B. Beagley, D. W. J. Cruickshank, A. Foord, S. Rustad

and V. Ulbrecht, J. Mol. Struct., 1976, 35, 139.

[52] M. Schmeisser and K. Lang, Angew. Chem., 1955, 6 7 ,156.

[53] R. C. Paul, K. K. Paul and K. C. Malhotra, Indian J. Chem., 1970, 8 ,

1030.

[54] K. O. Christe and R. D. Wilson, Inorg. Chem., 1973,12, 1356.

[55] K. O. Christe, R. D. Wilson and E. C. Curtis, Inorg. Chem., 1973,12,

1358.

[56] E. N. Sloth, L. Stein and C. W. Williams, J. Phys. Chem., 1969, 73, 278.

[57] A. Engelbrecht and P. Peterfly, Angew. Chem. Int. Ed. Engl., 1969, 8 ,

768.

[58] D. Lay cock, Ph.D. Thesis, University of Leicester, 1981 and references

cited therein.

121

[59] R. G. Syvret, Ph.D. Thesis, McMaster University, Canada, 1987.

[60] A. Engelbrecht, O. Mayr, G. Ziller and E. Schandara, Montash. Chem.,

1974,105, 796.

[61] A. Engelbrecht, P. Peterfly and E. Schandara, Z. Anorg. Allg. Chem.,

1971,384, 202.

[62] K. O. Christe. R. D. Wilson and C. J. Schack, Inorg. Chem., 1981, 20,

2104.

[63] H. A. Carter, J. N. Ruddick, J. R. Sams and F. Aubke, Inorg. Nucl.

Chem. Let., 1975,11, 29.

[64] K. Ztichner and O. Glemser, Angew. Chem. Int. Ed. Engl., 1972,11,

1094.

[65] H. Selig and U. El-Gad, Inorg. Nucl. Chem., Herbert H. Hyman

Memorial Volume 1976, Pergamon Press, ed. J. J. Katz and I. Sheft,

pg. 91.

[6 6 ] R. J. Gillespie and J. W. Quail, Proc. Chem. Soc., 1963, 278.

[67] N. Barlett and L. E. Levchuk, Proc. Chem. Soc., 1963, 342.

[6 8 ] J. H. Holloway, H. Selig and H. H. Classen, J. Chem. Phys., 1971, 54,

4305.

[69] C. J. Schack and K. O. Christe, J. Fluorine Chem., 1990, 49, 167.

[70] L. S. Bartlett, F. B. Clippard and J. E. Jacob., Inorg. Chem., 1976,15,

3009.

[71] M. Brownstein, R. J. Gillespie and J. P. Krasznai, Can. J. Chem., 1978,

2253.

[72] Ali-Reza Mahjoub and K. Seppelt, J. Chem. Soc., Chem. Commun.,

1991, 840.

[73] K. O. Christe, J. C. P. Saunders, G. J. Schrobilgen and W. W. Wilson, J.

Chem. Soc., Chem. Commun., 1991, 837.

[74] O. Ruff and H. King, Z. Anorg. Allg. Chem., 1930,190, 270.

[75] T. D. Cooper, F. N. Dost and C. H. Wang, J. Inorg. Nucl. Chem., 1972,

34, 3564.

122

[76] L. Andrews, F. K. Chi and A. Arkell, J. Am. Chem. Soc., 1974, 96,

1997.

[77] F. von Stadion, Gilberts Ann., 1816, 52, 197.

[78] F. Ammurmiiller and G. Magnus, Pogg. Ann., 1833, 28, 514.

[79] E. H. Appleman, J. Am. Chem. Soc., 1968, 9 0 ,1900.


[81] G. K. Johnson, P. N. Smith, E. H. Appleman and W. N. Hubbard, Inorg.

Chem., 1970, 9, 119.

[82] E. H. Appleman, Acc. Chem. Res., 1973, 6 , 113 and references cited

therein.

[83] E. H. Appleman, J. Am. Chem. Soc., 1969, 91, 4561.

[84] R. J. Gillespie and G. J. Schrobilgen, Inorg. Chem., 1974,13 ,1230.

[85] R. C. Ferreira, Bull. Soc. Chim. Fr., 1950, 135.

[8 6 ] R. S. Nyholm, Proc. Chem. Soc., London, 1961, 273.

[87] D. S. Urch, J. Inorg. Nucl. Chem., 1963, 25, 771.

[8 8 ] H. Moissan, Le Fluor, et ses Composes, Steinheil, Paris, 1900, pg. 123.

[89] O. Ruff and W. Menzel, Z. Anorg. Allg. Chem., 1931, 202, 49.

[90] A. J. Edwards and G. R. Jones, J. Chem. Soc., Chem. Commun., 1969,

1936.

[91] A. R. Mahjoub, A. Hoser, J. Fuchs and K. Seppelt, Angew. Chem. Int.

Ed. Engl., 1989, 28, 1526.

[92] I. Sheft, A. F. Martin and J. J. Katz, J. Am. Chem. Soc., 1956, 78, 1557.

[93] K. Christe and C. Schack, Inorg. Chem., 1970, 9, 2296.

[94] M. D. Lind and K. O. Christe, Inorg. Chem., 1972, 11, 608.

[95] A. J. Edwards and G. R. Jones, J. Chem. Soc., Chem. Commun., 1969,

1467.

[96] N. Binstead, S. J. Gurman and J. W. Campbell, EXCURV92, SERC

Daresbury Laboratory Program, 1992.

123

[97] 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. Am.

Chem. Soc., 1987,109, 3669.

[98] R. J. Gillespie and G. J. Schrobilgen, J. Chem. Soc., Chem. Commun.,

1974, 90.

[99] R. J. Gillespie and G. J. Schrobilgen, Inorg. Chem., 1974,13, 1230.

[100] K. O. Christe and R. D. Wilson, Inorg. Chem., 1975,14, 694.

[101] A. R. 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


[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, HSO3 F, was first prepared by Thorpe and Kirman

in 1892^ by the reaction of S 0 3 with HF (Eqn. 4.1).

S 0 3 + HF H S 03F 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, P4 O10, B(OH ) 3 and KM n0 4

react with it to produce POF3, BF3 and M n03F 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 organo-

fluorine compounds from hydrocarbons, the refining of lubricating oils and the

removal of metals from crude petroleum.

4.2. Properties.

Fluorosulphonic acid is isoelectronic with H2 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 H2 S 0 4, HSO3 F, S 0 2 F2, CF3 SO3 H and HF.

Property^ h 2 s o 4 HSO3 F s o 2 f 2 CF3 SO3 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

-H0 1 2 . 1 15.1 - 15.1 15.1

Figure 4.1. The protonation of acetamide.

CH3C.n h 2

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, NH2 and CH3 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^

HSOgF^CH 3C

H

NH

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 in1977 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 H2 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, [H3 0 ]+[S0 3 F]", which then undergoes slow hydrolysis. Woolf 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 H2 S 0 4 but instead, a stable equilibrium is

rapidly achieved.

127

Figure 4.2. Variety of fluorosulphate derivatives.

M ono & B identateA ' ' v ’i'V -V / '

1

t V 3A /3J

[Br(S03F)3][>3]

Bidentate

[c-Pd2(n-C0)2][(S03F)2][>7]

[Sn(S03F)2(CH3)2][181

M onodentate

[Br(S03F)4]li3]

[Au(S03F)4]-[|61

Binary

[Zn(S03F)2][14!

[Nb(S03F)5]ll5)

17 l l l A V * A C 111 Vi n -

* 1 U O I..............

O S U I pna t c........

High O xidation State

[Re02(S03F)3][19l

[XeF5(S03F)]P°l

T ern a ry

Cs[Sb(S03F)6][|6)

Cs2[Pt(S03F)6]l16l

H etero-bim etallic

[AgSn(S03F)6]P>l

[CuSn(S03F)6]P»

T riden ta te

[Fe(S03F)2]P21

[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 F2 -HS0 3F mixtures,

although other routes are known.

4.3.1. Syntheses involving S20 6F2 or S20 6F2-HS03F.

Fluorosulphate derivatives are most commonly synthesised using bis-

fluorosulphuryl peroxide, S2 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(S0 3 F ) 2 is formed by the

reaction, at 67°C, between a ten fold excess of S2 0 6 F2 and silver powder over

seven days P 4^

A mixture of S2 0 6 F2 in H S 03F is of particular synthetic use as it

combines the oxidising power of S2 0 6 F2 and the ionising ability of H S 0 3 F.

Several distinct advantages arise from the use of this sy 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 F2 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 F2 in HSO3 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 F2 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 + S2 0 6 F 2 HS° 3F ► [Au(S03 F)3][12] Eqn. 4.3.

M + n/2 S2 0 6 F2 HSQ3F ► [M (S0 3 F)4] Eqn. 4.4.

M = Ti, Zr, Hf, Sn, Pt, Ir, Pd and Ru.[24]

HSO3F2 M + 5 S2 0 6 F2 ► 2 [M (S03 F)5] Eqn. 4.5.

M = Nb, Ta and Sb . [ 2 6 1

Metal oxidation using S2 0 6 F2 and H S03F in the presence of

stoichiometric amounts of M [S0 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 10J+ (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[S0 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 AgF2 at 100 to 170°C. Since then, a

number of alternative routes to S2 0 6 F2 have been published and these include

the low temperature electrolysis of dilute solutions of K [S0 3 F] in

fluorosulphonic acid,^31 ̂ the reaction of CrF5 and S 0 3 in a 5:1 molar ratioJ32̂

the photolysis of C10S03F at ambient temperatures for 2-4 hoursj33̂ the

131

reaction of Cs[AgF4] with S 0 3 J 34̂ the low temperature combination of HSO3 F

and 0 2 +[AsF6 ]V 35] and finally, the reaction of F 0 S 0 2F and S 0 3 J23̂ However,

none of these methods have provided a more synthetically viable route to bis-

fluorosulphuryl 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^36̂

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 F2 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! (AgF2 on copper turnings as a support) for the fluorination

of SO3 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 S2 0 6 F2 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 F2 -H S03F 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

/°C

Reaction time

/D ays

Reference

[Cd(S0 3 F)2] 90 28 24

[Zn(S0 3 F)2] 90 2 1 24

[Pd(S0 3 F)3] 25-120 3 28

[Ru(S0 3 F)3] 60 1 28

[Ti(S0 3 F)4] 25-60 1 0 28

[Zr(S0 3 F)4] 25-120 2 1 28

[Hf(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

[Nb(S0 3 F)5] 40 3 15

[Ta(S0 3 F)5] 40 3 15

Cs2 [Ti(S0 3 F)6] 25 5 28

Cs2 [Sn(S0 3 F)6] 25 0.5 28

Cs2 [Pt(S0 3 F)6] 80 3 28

Cs2 [Ir(S0 3 F)6] 150 1 0 28

Cs2 [Ge(S0 3 F)6] 50 2 28

Cs[Sb(S0 3 F)6] 70 0.5 29

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^37̂ 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 S2 0 6 F2

133

produces O2 and S2 0 5 F2 ,t38] hence reaction vessels must be vented regularly to

avoid explosions. At certain temperatures, the catalytic decomposition of

S2 0 6 F2 to 0 2 and S2 0 5 F2 may occurJ37̂ This was observed during the

formation of Ir(S0 3 F)4. At 130°C, Ir(S0 3 F ) 4 decomposes to Ir(S0 3 F)3, S2 0 5 F2

and 0 2, the Ir(S0 3 F ) 3 is then re-oxidised by S2 0 6 F2 to form Ir(S0 3 F)4.

H S 03F HF + S 0 3 Eqn. 4.6.

The reaction of bis-fluorosulphuryl peroxide and H S 03F with a variety

of metal carbonyls and chlorides has been investigated J24 ̂ 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 Woolf described the preparation of bivalent fluorosulphate

derivatives of Mn, Fe, Co, Ni, Cu, Zn and CdJ39̂ 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 [CH3 C 0 2]‘ >

[S04]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 J40 ̂Zr(C 0 2 CF3)4 4̂1 ̂

and A g(C 0 2 CF3)̂ 42 ̂ in H S 03F has been shown to produce Fe(S0 3 F)3,

Z r(S 0 3 F ) 4 and A gS03F 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 F2 and S2 0 5 F2 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).

TaCl5 + S2 0 6 F2 -> [TaO(S03 F)3] Eqn. 4.7.

WC16 + S2 0 6 F2 [W 0(S0 3 F)4] Eqn. 4.8.

Ti(OCH 3 ) 4 + S2 OsF2 -> [Ti(OCH3 )2 (S 0 3 F)2] Eqn. 4.9.

An alternative approach is the reaction of metal carbonyls with S2 0 6 F2

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)J2 4 1 Also as explained earlier, the use of

S2 0 6 F2 and CO is not ideal.

[Mn2 (CO)10] + S2 0 6 F2 —̂ [M n(S03 F)4] Eqn. 4.10.

Co2 (CO ) 8 + S2 0 6 F2 —» [Co(S03 F)2] Eqn. 4.11.

Cr(CO ) 6 + S2 0 6 F2 —» [Cr(S03 F)3] 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 rS03F.

Bromine fluorosulphate, B rS0 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 F2 for its synthesis. Its limited use

includes the synthesis of the noble metal fluorosulphates [Pd(S0 3 F)2]

[Pt(S0 3 F)4] / 45 ̂ [Ag3 (S 0 3 F)4] / 46 ̂ and [Au(S03 F)3 ]J45 ̂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 AgF2 and S 0 3 produces [Ag(S0 3 F)2],[42̂ and this is expected when

one considers the catalytic nature of AgF2 during the formation of S2 0 6 F2 from

S 0 3 and F2 (Section 4.3.1.1. and Eqn. 4.14). The reaction between Ag, BrF3

and S 0 3 involves the fluorination of silver followed by the insertion of S 0 3 J47 ̂

whilst the reaction between S 0 3 and CrF5, at -22°C, offers a route to

AgF2 + S 0 3 —> [Ag(S0 3 F)2] Eqn. 4.13.

[A g(S0 3 F)2] + F2 —> AgF2 + S2 0 6 F2 Eqn. 4.14.

3 Ag + 2 BrF3 + 3 S 0 3 —> 3[ A g(S0 3 F)2] + Br2 Eqn. 4.15.

WF6 + S 0 3 [WF2 (S 0 3 F)4] [48] Eqn. 4.16.

CrF5 + S 0 3 —> [Cr(S0 3 F)3] + S2 O^F2 Eqn. 4.17.

136

4.3.5. Oxidising reactions involving H S 03F.

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 HSO3 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,

Sb(S0 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) fluoro-

fluorosulphate 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

S2 0 6 F2, in HSO 3 F results in the formation of the unstable pentafluorosulphate,

137

[Nb(S0 3 F)5], this loses S 0 3 and eventually forms the solid [NbF2 (S 0 3 F)3]^15̂

(Eqn. 4.18).

N b(S0 3 F ) 5 ^ N bF(S0 3 F ) 4 + S 0 3 NbF2 (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 oxo-

fluorosulphate complex. As noted earlier, substitution reactions between metal

chlorides or carbonyls and S2 0 6 F2 usually produces an oxo-fluorosulphate as

the p ro d u c t^ (Eqn.’s 4.19 and 4.20).

WC16 + S2 0 6 F2 -» [W 0(S 0 3 F)4] + S2 0 5 F2 Eqn. 4.19.

W(CO ) 6 + S2 0 6 F2 ^ [W 0(S 0 3 F)4] + S2 0 5 F2 Eqn. 4.20.

Further examples of this type of reactivity include the formation of

[V 0 (S 0 3 F)3], [N b0(S 0 3 F)3], [M o0 2 (S 0 3 F)2] and [Re0 2 (S 0 3 F)3].[24] The

elimination of S2 0 5 F2 does not always result exclusively in the formation of an

oxo-fluorosulphate complex. For example, the decomposition of Ir(S0 3 F ) 4

yields S2 0 5 F2, but Ir(S0 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 S2 0 6 F2 (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(S0 3 F ) 2 and B rS0 3 F. A similar

process was observed for A g(S0 3 F ) 2 which decomposes at 210°C to produce

A g(S0 3 F) and S2 0 6 F2.[53]

o

Pdu [PdIV(S 0 3 F)6] 2 Pd(S0 3 F ) 2 + S2 OeF2 Eqn. 4.21.

138

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 (S0 3 F)x —) MFjp + x SO3 Eqn. 4.22.

2 M (S 0 3 F)a; —̂ + x SO2 F2 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(S0 3 F ) 2 is pyrolysed to CaF2 whereas Ba(S0 3 F ) 2

produces B a(S04). 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 O-Monodentate 0,0'-Bidentate 0 ,0 ’,0"-Tridentate

_ M ' — . ___ __Q ° \ / > 0 ^ 0 / 0

o/S\ ^A ,C3vatS Cs atS Cs atS C3vatS

For potassium fluorosulphate, K [S0 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 (E) S03 asym defn

570 m ^3 (A) S03 Sym defn

480 m V6 (E) SF defn

For C s[S0 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(SQ3 F)2].

Infrared spectra / cm" 1

[Co(S0 3 F)2] [Fe(S0 3 F)2] [Ni(S0 3 F)2] Assignment

1265 vs 1261 vs 1262 vs v4 (E)

1109 s 1118s 1 1 2 0 s Vi (A)

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)

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 Fe(S0 3 F ) 2 and N i(S0 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^1̂ and assignments for K[Br(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 1250 s Vi (A1) S03

592 s v5 (E) 1 2 2 0 w sh 1 2 2 2 w sh -

570 m V3 (A) 970 m 1 0 0 2 m v4(A') S03as

409 m V6 (E) 834 m 837 m v2 (A') SF- - 578 ms 582 ms V5 (A')- - 553 ms 554 ms v8 (A")- - 615 vs 620 vs V3 (A’)- - 406 w 407 w V9 (A")

- - 239 vs 239 vs V6 (A')

The fluorosulphate group can act as a monodentate ligand and to

demonstrate, this the Raman spectra of K[Br(S0 3 F)4] and K [I(S0 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 [S0 3 F]. The S-F stretching mode is

observed at 834 cm '1, which is higher than that observed for K [S0 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(S0 3 F)4]- 454 cm" 1 [v4 (E) v4 (A') + v 7 (A")] and [I(S0 3 F)4]'

407 cm '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 ]J56 ̂ the tin compound having been

crystallographically characterised. As a result of the 0 ,0-bidentate 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[Br(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 [S0 3 F] and K[Br(S0 3 F)4], whereas the

position of the third S-O mode remains virtually the same as that for K [S0 3 F].

Table 4.6. Infrared vibrational data, cm '1, for the -S 03F group in [Fe(S0 3 F)3],

[Sn(S0 3 F)2 Me2], [Sn(S0 3 F)2 Cl2], K[Br(S0 3 F)4] and K [S0 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 is the case for [Br(S03 F)3],[13) [I(S0 3 F)3] [13] and

[Sn(S0 3 F)4] J ' The infrared spectra are very complex and in the case of

B r(S0 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(S03 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(S0 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)(C0)5] [19]

(M = Mn or Re) and [Sn(S0 3 F)2] For example, the vibrational data for

[Re(S0 3 F)(C0)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 (-60 cm " 1 for Re) and this splitting is too large to be solely due to

site affects (for example in [NO][S0 3 F] site effects result in a splitting of 10-

20 cm_1 )J58 ̂ iii) The position of v(SO—M), 1030 cm"1, is consistent with the

highly ionic character proposed for the M -OS02F 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[Br(S0 3 F)4], Table

4.5).

144

Table 4.7. Comparison and assignment of the infrared vibrational data, cm '1,

for K [S0 3 F] and [Re(S0 3 F)(CO)5].

k s o 3f Mode [Re(S03 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 s y m ( S 0 2 )

- - 1255 m V a s y m ( S 0 3 )

1080 s V , (A) 1170 w V Sy m ( S 0 3 )

- - 1 1 2 0 w V s y m ( S O s )

- - 1030 m V ( S O - - M )

786 s v2 (A) 760 m v (S-F)

4.5.2. X-ray crystallography.

The polymeric nature of many transition metal fluorosulphates has

prevented their crystallographic analysis, although recently compounds such as

Cs[A u(S0 3 F)4],[29) Cs2 [Pt(S0 3 F)6],|29) [Sn(S0 3 F)2 Me2][59] and

Cs[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 A, sulphur-oxygen (terminal)

-1.37-1.45 A and sulphur-fluorine -1.45-1.57 A.

145

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 1 1 9 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 J64,65 ̂ 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[S0 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 [S0 3 F] and [NH4 ][S 0 3 F].

The addition of fluorosulphonic acid to C s[S0 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(S0 3 F)2]- is 1 .2 1 0 (2 ) A, which is rather short when compared to

[H(OTeF5 )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[S0 3 F]. It

appears that bonding to the peripheral oxygen and fluorine atoms in the pair

[S 0 3 F]' and [H (S0 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[S0 3 F]. Overall, the bonds involving peripheral atoms

strengthen slightly from C s[S0 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

Cs[A u(S0 3 F)4] and Cs[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:-^29̂ 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 Cs[S03F]. The expected trends may be slightly

modified by inter atomic contacts to the cation Cs+.

Table 4.9. Bond lengths and angles for Cs[Au(S03F)4] and Cs[Sb(S03F)6].

Parameter Cs[Au(S03F)4] Cs[Sb(S03F)6]

</(M-Ob) / A 1.968(4) 1.955(2)

d(S-Ob) / A 1.508(4) 1.516(2)

Z(M ObS) 125.2(3) 136.6(1)

d{ S-Ot) / A 1.393(5) 1.396(3)

1.402(6) 1.409(4)

d( S -F )/ A 1.523(6) 1.486(3)

Ob = bridging oxygen, Ot = terminal oxygen

Figure 4.4. Molecular structures1̂ of a) [Au(S03F)4]" and b) [Sb(S03F)6] ' .

149

The anion [Au(S0 3 F)4]" possesses Q symmetry and the Au-O bond

distances average 1.972(4) A. The coordination number of the caesium cation

in this molecule is 1 0 .

Within Cs[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

[Au(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 re 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[S0 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(S03 F)2]" <

125.2(3)° for [Au(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-C0)2][(S03F)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 Pd-

O 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[S03F].

Figure 4.5. Molecular structure of [Au(S03F)3].

The molecular structure of [Au(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) A for the mono and

Au-Oav = 2.018(7) A 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[S0 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,^66̂ 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 cm '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, [Au(S03 F)(C0)] J69̂ This was serendipitous

and stemmed from investigations to detect the formyl cation, [HCO]+. Attempts

to observe the cation by NMR spectroscopy using 1 3 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 back-

donation. 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(S0 3 F ) 3 should act as a fluorosulphate

ion acceptor J 7 0 1 Later, it was shown that A u(S0 3 F ) 3 does indeed behave as a

Lewis acid in fluorosulphonic acid to form the conjugate base [Au(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.

Work-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 [Au(S03 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(S0 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(S0 3 F)(C0)] 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 + [S0 3 F] + S2 O^F5 Eqn. 4.24.

Au+ + 2 CO -> [Au(CO)2]+ Eqn. 4.25.

154

[Au(CO)2]+ + [SO3F]- -> [Au(S03F)(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-[Pt(S0 3 F)4] ̂ was soon undertaken,^13,68̂ and led to the

complexes [cis-Pt(S 0 3 F)2 (C 0)2] and [cw-Pd(S03 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 7t-back donation is

highlighted by the high CO-stretching frequencies, vav(CO) of 2218 cm-1. The

CO bond lengths are also short, 1.106(6) and 1.114(6) A, when compared to

those of gaseous carbon monoxide (cf. 1.12822 A). 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 iridiumJ37̂ 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 F2 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(S03 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(S03 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(S0 3 F)3 (C0 )3 ] .

Again, as was observed for cw-[Pd(S03 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 [Ru(S0 3 F)3] and [Os(S03 F)3] in SbF5 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 [ds-P t(S0 3 F)2 (C 0)2]

and [c/.s-Pd(S03 F)2 (C 0)2], in SbF5 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. [Sb2 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 - 101 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 H2 0 , is

[H3 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 -H0 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 S03F results in the formation of the

following species, [H2 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

[H2 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- bridgesJ28̂

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(S0 3 F)n, where E (S0 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 HSO3F + E (S0 3F)n m [H 2S 0 3F]+ + [E(S03F)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-Au(S0 3 F)3:-

1) The proton donor strength should be very high and is limited by the acidity of

[H2 S 0 3 F]+.

2) The nucleophilicity or electron pair donor ability of the conjugate base ion,

[Au(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 (S0 3 F)4] and Cs2 [M (S0 3 F)6] species, where M = Ti, Zr or Hf However,

all three Group four binary fluorosulphates are insoluble in H S0 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 (S0 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. [Au(S0 3 F)3] which is dimeric, is it feasible to use

binary fluorosulphates in conjugate H S 03F 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 bis-

fluorosulphuryl 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 Woolf 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], [Ru3 (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 S03F.

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 H2 S0 4,̂ 8°] whereas more recent work by

Hope et al. used AHFJ81̂

Anhydrous HF is a convenient solvent for the fluorination of transition-

metal 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 [Ru3 (CO)12], [Os3 (CO)12] and [Ir4 (CO)12] react with AHF to produce

[Ru3 (CO)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 (CO)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 HSO3 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 /(CH) = 3.1 Hz) [c f 8159.0 ppm (d 2 /(CH) = 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% H2 S 0 4 J 82̂

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 /(CH) = 3.2 Hz), 8170.8, 8168.9,

8163.9 and 8161.8 ppm (d 2 /(CH) = 7.1 Hz) (cf 8176.4, 8169.8 (d 2 J(CH) = 2.9

Hz), 8165.1, 8167.0 (d 2 /(CH) = 2.0 Hz) and 8159.5 ppm (d 2 /(CH) = 7.0 Hz)

in HF), and these were assigned to [Os3 (CO)1 2 H]+J81,83̂ 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

[Ru3 (CO)12] in H S 03F 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 [Ru3 (CO)1 2 H]+ and also corresponds to the

peak at 8-19.4 ppm reported as the only resonance from a solution of

[Ru3 (CO)12] in 98% H2 S 0 4 J 80̂ 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 HSO3 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 Me2 S 0 4 J 80̂ 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 +

CO,OC: CO,

OC;

OC CO.CO,

CO,CO,

2+

aTrans hydride. bCis hydride.

Figure 4.8. Carbon-13 and 13C{ 1H} NMR spectra of [Ir4(CO)12] in HSO3F.

147.2 147.0 146.8 146.6 146.4 146.2 146.0 145.8 145.6 145.4 145.2 145.0 144.8 144.6 144.4 144.2(ppm)

4.9.1. Summary.

The reactivity of fluorosulphonic acid with the three metal carbonyls

[Ir4(CO)12], [Ru3 (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 (H0 = -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°CJ82̂ 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 HSO3 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 J81!

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 S0 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. Deeming et a l

demonstrated that the substitution of carbonyl ligands in [0 8 3 (0 0 )!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 NMR 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(S0 3 F ) 3 / H S 03F 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 [Au(S0 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

[M3 (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 HSO3F.

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)J67,85̂ 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 03, passivate the metal and prevent

complete reaction. Brazier and Woolf observed that iron did not react with

boiling fluorosulphonic acid[49̂ 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 S03F was attempted as a convenient

route to [Fe(S0 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(S03 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 03F affords [Fe(S0 3 F)2] . In view of the polymeric nature

of [Fe(S0 3 F)2] it is understandable why mass spectrometry failed to produce

any identifiable patterns.

Table 4.10. Infrared spectroscopic data for K[SQ3 F] and [Fe(SQ3 F)2].

K [S 0 3 F]a

cm"1

[Fe(S0 3 F)2]b

cm'1

[Fe(S0 3 F)2]c

cm"1

Assignment.

1280 s 1270 s 1261 vs v4 (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

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[S0 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 [Mn2 (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 HS0 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 S03F 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 [Mn2 (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 [Mn(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 solidJ19̂

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 [S0 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 3F)].

K [S0 3 F]b Assignment [Re(C0)5 (S 0 3 F)]b [Re2 (CO)10] +

h s o 3f

[Re(CO)5 Cl] +

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 v7 (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 v5 (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 - - v6 (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 [S0 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

[Re(C0)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:- [Re(C0)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 03F 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 [Re(C0)5 (S 0 3 F)]. The reaction between [Re(CO)5 Cl] and H S03F

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

[Re(C0)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

[Re(C0)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 HSO3F and [CP2MX2 ]

(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 (S0 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

[Cp2 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],

[Cp2 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

[Cp2 TiMe2] 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 03F 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

[Cp2 Ti]+ m/z =178. The infrared vibrational data are presented in Table 4.12,

also listed, for comparison, are the infrared vibrational data for K[Br(S03 F)4],

for which a covalent monodentate interaction exists between the fluorosulphate

ligand and the bromine centred13̂ 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) + v4 (1047 cm"1), which at 393

cm " 1 is similar to that observed for K[Br(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(S0 3 F)2]

and K[Br(S0 3 F)4].

K [Br(S0 3 F)4] / cm - 1 [Cp2 T i(S0 3 F)2] / cm ' 1 Assignment

- 2968 s sh v(CH)

- 2929 s v(CH)

- 2865 s sh v(CH)

1424 m 1440 m v 7

1407 w 1401m

970 m 1047 m V4

- 1023 s sh

1237 s 1236 s br Vl

1 2 2 0 w sh 1205 m sh

834 s 8 8 6 m v 2

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,

[Cp2 TiMe2 ]-H S0 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 Hf 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 pn —» dn bonding, fluorine being the strongest n donor

but the weakest n acceptor. As the ability to pn —> dn bond decreases so the

Cp-metal interaction increases. The ]H NMR data for [Cp2 Ti(S0 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 = -SO3 F, -OTeF5, -F and -Cl).

Compound 8 lH / ppm

[Cp2Ti(S03 F)2]a 7.2

[Cp2 Ti(OTeF5)2]b 6.9

[Cp2 TiCl2]b 6.56

[Cp2 TiF2]b 6.44

a Recorded in HSO3 F. b Recorded in CD2 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 Cs2 [Pt(S0 3 F)6], Cs[Sb(S0 3 F)6] and

Cs2 [A u(S0 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(S0 3 F)2 ]-H S0 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 HSO3 F solvent peak

8 19F / ppm

[Cp2 T i(S 0 3 F)2] 57.8 40.8

Cs2 [Pt(S0 3 F)6]a 47.7 40.7

Cs[Sb(S0 3 F)6]a 46.4 40.9

Cs2 [Au(S0 3 F)4] a 45.5 40.8

a Ref. 29.

4.13. The Reaction Between HSO3F and [W(CO)6 J.

The reaction between [W(CO)6] and HSO3 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) =116 Hz. Proton-1

NMR 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 CH2 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 HSO3F.

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 cm 'V 88̂ 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 NMR 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 CH2 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 sulphur-

fluorine 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

HSO3F.

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. Woolf 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 03F has established a route to

the previously unknown [Cp2 Ti(S0 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, 61,921.

[2] A. Engelbrecht, Angew. Chem., Int. Ed. Engl., 1965, 4 ,641.

[3] R. C. Thompson, Inorganic Sulphur Chemistry, Elsevier, ed. G.

Nickless, 1968, ch. 17, 587-606 and references cited therein.

[4] R. J. Gillespie, Acc. Chem. Res., 1968,1, 202.

[5] R. J. Gillespie and T. Birchall, Can. J. Chem., 1963, 4 1 ,148.

[6 ] CRC Handbook o f Chemistry and Physics, ed. R. C. Weast, 57th edn.

1979.

[7] G. A. Olah, G. K. S. Prakash and L. Sommer, Superacids, Wiley, New

York, 1985 and references cited therein.

[8 ] R. J. Gillespie and J. Liang, J. Am. Chem. Soc., 1988,110,6053.

[9] R. Jost and J. Sommer, Rev. Chem. Intermed., 1988, 9,1.

[10] R. D. Howells and J. D. McCown, Chem. Rev., 1977, 77,69.

[11] A. A. Woolf, J. Chem. Soc., 1954, 2840.

[12] H. Willner, S. J. Rettig, J. Trotter and F. Aubke, Can. J. Chem., 1991,

69,391.

[13] H. A. Carter, S. P. L. Jones and F. Aubke, Inorg. Chem., 1970, 9,2485.

[14] J. Gorbeau and J. B. Milne, Can. J. Chem., 1967, 45,2321.

[15] W. V. Cicha and F. Aubke., J. Fluorine Chem., 1990, 47, 317.

[16] D. Zhang, S. J. Rettig, J. Trotter and F. Aubke., Inorg. Chem., 1996,35,

6113.

[17] C. Wang, M. Bodenbinder, H. Willner, S. J. Rettig, J. Trotter and F.

Aubke., Inorg. Chem., 1994, 33,779.

[18] S. P. Mallela, S. Yap and F. Aubke, Inorg. Chem., 1986, 25,4074.

[19] S. P. Mallela and F. Aubke, Inorg. Chem., 1985, 24, 2969.

[20] M. Einsenberg and D. D. Des Marteau, Inorg. Chem., 1972,11, 2641.

[21] S. P. Mallela, K. Lee, P. F. Gehrs, J. I. Christensen, J. R. Sams and F.

Aubke, Can. J. Chem., 1987, 65, 2649.

181

[22] C. S. Alleyne, K. O. Mailer and R. C. Thompson, Can. J. Chem., 1974,

52, 336.

[23] A. W. Jache, Adv. Inorg. Chem. Radiochem., 1974,16, 177 and

references therein..

[24] F. Aubke, M. S. R. Cader and F. Minstry, Synthetic Fluorine Chemistry,

eds. G. A. Olah, R. D. Chambers and G. K. S. Prakash, Wiley, New

York, 1992, pg. 43 and references therein.

[25] J. M. Shreeve and G. H. Cady., Inorg. Synth., 1963, 7 ,124.

[26] D. Zhang, S. J. Rettig, J. Trotter and F. Aubke., Inorg. Chem., 1995, 34,

2269 and references therein.

[27] W. V. Cicha, F. G. Herring and F. Aubke, Can. J. Chem., 1990, 6 8 ,102.

[28] F. Mistry and F. Aubke, J. Fluorine Chem., 1994, 6 8 , 239.


6113.

[30] F. B. Dudley and G. H. Cady, J. Am. Chem. Soc., 1957, 79,513.

[31] S. Singh and R. D. Verma, Ahstr. Int. Symp. Fluorine Chem. 10th,

Vancouver, Aug. 1982 Abstr. 1-76.

[32] S. D. Brown and G. L. Gard, Inorg. Nucl. Chem. Lett., 1975,11 ,19.

[33] C. J. Schack and K. O. Christe, Inorg. Nucl. Chem. Lett., 1978,14,293.

[34] P. C. Leung and F. Aubke, J. Fluorine Chem., 1979, 14, 347.

[35] A. Smalc, Vestn. Slov. Kem. Drus., 1974, 21,5.

[36] G. H. Cady, Inorg. Synth., 1968,11, 155.

[37] C. Wang, A. R. Lewis, R. J. Batchelor, F. W. B. Einstein, H. Willner and

F. Aubke, Inorg. Chem., 1996, 3 5 ,1279.

[38] F. B. Dudley and G. H. Cady, J. Am. Chem. Soc., 1963, 85 ,3375.

[39] A. A. Woolf, J. Chem. Soc. (A), 1967, 355.

[40] A. Storr, P. A. Yeats and F. Aubke, Can. J. Chem., 1971, 50,452.

[41] S. Singh, M. Bedi and R. D. Verma, J. Fluorine Chem., 1982, 20, 107.

[42] P. C. Leung and F. Aubke, Inorg. Chem., 1978,17, 1765.

[43] J. E. Roberts and G. H. Cady, J. Am. Chem. Soc., 1960, 92, 352.

182

[44] K. C. Lee and F. Aubke, Can. J. Chem., 1977, 55,2437.

[45] W. M. Johnson, R. Dev and G. H. Cady, Inorg. Chem., 1972,11,2260.

[46] K. C. Lee and F. Aubke, J. Fluorine Chem., 1982,19, 501.

[47] A. A. Woolf, J. Fluorine Chem., 1971,1,127.

[48] H. C. Clark and H. J. Emeleus, J. Chem. Soc., 1957, 4778.

[49] J. N. Brazier and A. A. Woolf, J. Chem. Soc. (A), 1967, 99.


3153.

[51] R. J. Gillespie and O. C. Vaidya, J. Chem. Soc., Chem. Commun., 1972,

40.

[52] K. C. Lee and F. Aubke, Can. J. Chem., 1981, 59 ,2835.

[53] P. C. Leung and F. Aubke, Inorg. Nucl. Chem. Lett., 1977,13, 263.

[54] E. L. Muetterties and D. D. Coffman, J. Am. Chem. Soc., 1958, 80,

5914.

[55] E. L. Muetterties, J. Amer. Chem. Soc., 1958, 80, 5911.

[56] P. A. Yeats, B. L. Poh, B. F. E. Ford, J. R. Sams and F. Aubke, J. Chem.

Soc. (A), 1970, 2188.

[57] G. Hwang, C. Wang, M. Bodenbinder, H. Willner and F. Aubke, J.

Fluorine Chem., 1994, 6 6 , 159.

[58] A. M. Qureshi, H. A. Carter and F. Aubke, Can. J. Chem., 1971, 49, 35.

[59] F. H. Allen, J. A. Lerscher and J. Trotter, J. Chem. Soc. (A), 1971, 2507.

[60] F. A. Hohorst and J. M. Shreeve, Inorg. Chem., 1966, 5, 2069.

[61] T. Birchall, G. Denes, R. Faggiani, C. S. Frampton, R. J. Gillespie, R.

Kapoor and J. E. Vekris, Inorg. Chem., 1990, 2 9 ,1527.

[62] D. C. Adams, T. Birchall, R. Faggiani, R. J. Gillespie and J. E. Vekris,

Can. J. Chem., 1991, 69, 2122.

[63] S. P. Mallela, S. T. Tomic, K. Lee, J. R. Sams and F. Aubke, Inorg.

Chem., 1986, 25, 2939.

[64] J. R. Sams, R. C. Thompson and T. B. Tsin, Can. J. Chem., 1977, 55,

115.

183

[65] A. L. Arduini, M. Garnett, R. C. Thompson and T. C. T. Wong, Can. J.

Chem., 1975, 53, 3812.

[6 6 ] F. Aubke, Inorganic Fluorine Chemistry : Towards the 21st Century,

eds. J. S. Thraser and S. H. Strauss, American Chemical Society,

Washington, 1994, pg. 350.

[67] F. A. Cotton and G. Wilkinson, Advanced Inorganic Chemistry, 5th ed.,

1988.

[6 8 ] P. Schtitzenberger, C. R. Hebd. Seances Acad. Sci., 1870, 7 0 ,1134; P.

Schtitzenberger, C. R. Hebd. Seances Acad. Sci., 1870,7 0 ,1287; P.

Schtitzenberger, Bull. Chim. Fr., 1870,14,97.

[69] H. Willner and F. Aubke, Inorg. Chem., 1990, 29,2195.

[70] W. M. Johnson, R. Dev and G. H. Cady, Inorg. Chem., 1972,11,2260.

[71] K. C. Lee and F. Aubke, Inorg. Chem., 1979,18,389.

[72] K. C. Lee and F. Aubke, Inorg. Chem., 1984, 23, 2124.

[73] G. H. Hwang, C. Wang, F. Aubke, H. Willner and M. Bodenbinder,

Can. J. Chem., 1993, 7 1 ,1532.

[74] C. Wang, H. Willner, M. Bodenbinder, R. J. Batchelor, F. W. B.

Einstein and F. Aubke, Inorg. Chem., 1994, 33, 3521.

[75] C. Wang, B. Bley, G. Balzer-Jollenbeck, A. R. Lewis, S. C. Siu, H.

Willner and F. Aubke., J. Chem. Soc., Chem. Commun., 1995, 2071.

[76] B. Bley, H. Willner and F. Aubke, Inorg. Chem., 1997, 3 6 ,158.

[77] C. Bach, H. Willner, C. Wang, S. J. Rettig, J. Trotter and F. Aubke,

Angew. Chem., Int. Ed. Engl., 1996, 35,1974.

[78] M. Bodenbinder, G. Balzer-Jollenbeck, H. Willner, R. J. Batchelor, F.

W. B. Einstein, C. Wang and F. Aubke, Inorg. Chem., 1996, 35,92.

[79] R. J. Gillespie and J. Liang. J. Am. Chem. Soc., 1988,110, 6053.

[80] J. Knight and M. J. Mays, J. Chem. Soc. (A), 1970, 711.

[81] S. A. Brewer, J. H. Holloway and E. G. Hope, J. Fluorine Chem., 1995,

7 0 ,167.

184

[82] A. J. Deeming, B. F. G. Johnson and J. Lewis, J. Chem. Soc. (A), 1970,

2697.

[83] A. A. Koridze, O. A. Kizas, N. M. Astakhova, P. V. Petrovskii and Y.

K. Grishinn, J. Chem. Soc., Chem. Commun., 1981, 853.

[84] C. O. Quicksall and T. S. Spiro, Inorg. Chem., 1968, 7, 2365; B. F. G.

Johnson, J. Lewis and P. A. Kilty, J. 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


[87] Comprehensive Oranometallic Chemistry, eds. E. W. Abel, F. G. A.

Stone and G. Wilkinson, Pergamon Press, Oxford, 1982, vol 3, ch 28


[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. Metal 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 positive-

pressure dry box [Vacuum Atmospheres Co., VAC NE-42 Dri Lab.]

186

Figure 5.1. Metal vacuum line.

To high vacuum

9 = f i (r

To rough vacuum

F„ hi letAr Inlet

8 = 1 mMetal

connectorA.E. tap

attached

attachedreactors

Bourdon Gauge

To high vacuum

A.E. Tee piece

A E . Cross piece

13= ^ ( r ^

To rough vacuum

Toattachedreactors

Toattachedreactors

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. Metal 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.

To A utoclave Engineers valve

tW ater-cooled lid

Water out Water in

PTFE ‘O ’ R ing

Stainless-steel body

Figure 5.3. Glass apparatus.

Y oung’s greaseless tap

6mm glass connector

Glass vessel

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 coarse-

control 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

A.E. connector with 6mm adaptor

Chemcom needle valve

U T T )

Youna sgreaseless tap

6mm Glass connector

rJ

1jChemcom

tee coupling

Glass vessel

FEP Tubing (6mm o.d., 2mm i.d.)

Molecular sieves

Solvent

FEP Tubing (4mm o.d., 2mm i.d.)

191

9868

5.3. Analytical Techniques.

5.3.1. Nuclear magnetic resonance spectroscopy.

1 H, 1 9 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% H3 PO4 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. NMR samples fitted inside a 5 mm o.d. precision NMR 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 E X ^ and

EXCURV 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 curved-

wave 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.

Assembled 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 CaH2 for approximately four

hours. The resulting liquid was then distilled on to fresh P2 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 ofXeF2-

Xenon difluoride was prepared as described by Holloway,^ 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(OSeF5)2.

Xenon bis(seflate) was prepared as described by Seppelt,^ 1986.

Selenium dioxide, S e0 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 Se02. 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 U-

tube 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 SeOF2 was condensed on to the XeF2 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[BrO4].

Potassium perbromate was prepared using the method described by

Appleman, 1972.^ The initial oxidation involves the action of elemental

fluorine on an aqueous alkaline solution of K [Br03]. 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, Na[Br03], (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 arning, 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 [Br04] 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[Br04] 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 BaF2 and Ba[Br03]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[Br04] 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 6 mm FEP tubing to a

Chemcom T-piece. Two separate 6 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. [Re2 (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 ofBrF3.

Bromine trifluorine, BrF3, was distilled from the cylinder into a FEP U-

tube. The BrF3 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 ofBrF5.

This was distilled directly from the cylinder into an FEP U-tube. The

BrF5 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

BrF5 (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 [BrF2 ][AsF6] and [BrF4 ][Sb2Fj j].

A 6 mm FEP tube was connected via a satellite line to the metal

manifold. Also connected to the satellite were BrF3 and AsF5 or BrF5 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 AsF5 was carefully monitored using the metal-line gauge

and once the pressure remained constant all the BrF3 was judged to have

reacted. The solid was pumped under dynamic vacuum to remove excess of

AsF5 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 BrF5 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, Cs[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 BrF5 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[Br04] (ca. 0.7 mmol), was loaded into a

passivated 6 mm FEP tube in a dry box. The FEP tube, along with BrF5 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[Br04] . Finally BrF5 (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 r03F

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 r03F 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 HSO3 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 S03F at reduced pressures. Once the excess of

H S 03F was removed, the remainder of the volatile materials were pumped

away under dynamic vacuum using the rough pump. This was also very time-

consuming 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 ofBrOF3.

The reaction between Li[N 03] and BrF5 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 BrF5 and

K[BrF4] as the identifiable products.

205

5.5.13. Attempted synthesis o f B r02F.

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 BrF5 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 BrF5

(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 r02F

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[Br02F2] 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[Br0 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 [Br0 2 F2 ]-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 dm3 nickel cans for convenience.

Fluorosulphonic acid, HS03F : 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


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


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


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


Lithium nitrate, Li[N03] : Aldrich Chemical Company Ltd. Dried in an


Iodine, I2 : Aldrich Chemical Company Ltd. Used as supplied.

Methyl manganese pentacarbonyl, [MeMn(CO)5] : Prepared according to

the literature^10̂ 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


Rhenium pentacarbonyl chloride, [Re(CO)5CI] : Aldrich Chemical


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, Se02 : Aldrich Chemical Company Ltd. Dried in an


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.



[3] J. H. Holloway, J. Chem. Soc., Chem. Commun., 1966, 22.



[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

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